Download 1999 U. S. NATIONAL CHEMISTRY OLYMPIAD

1999 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
James S. Bock, Gateway High School, PA
Edward DeVillafranca (retired), Kent School, CT
Peter E. Demmin (retired), Amherst Central High School, NY
John Krikau, American Chemical Society, DC
Patricia A. Metz, University of Georgia, GA
Ronald O. Ragsdale, University of Utah, UT
Helen M. Stone (retired), Ben L. Smith High School, NC
Diane D. Wolff, Ferrum College, VA
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron® sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to
the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 26, 1999, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 questions
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 15 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write you name on the answer sheet; an ID number is already entered for you. Make a record of
this ID number as you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement which is followed by four possible choices. Select the single choice that best answers the question or completes
the statement. Then use a pencil to blacken the space on your answer sheet having the same letter as your choice. You may write on
the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 26, 1999.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K milli- prefix
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G mole
A frequency
ν Planck’s constant
N A gas constant
R pressure
°C gram
g rate constant
c hour
h second
C joule
J speed of light
E kelvin
K temperature, K
Ea kilo- prefix
k time
H liter
L volt
S measure of pressure mmHg volume
CONSTANTS
m
m
M
mol
h
P
k
s
c
T
t
V
V
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(289)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as a USNCO National Exam after April 26, 1999
DIRECTIONS
§ When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
§ Make no marks on the test booklet. Do all calculations on scratch paper provided by your instructor.
§ There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
§ Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which substance is most likely to be soluble in a
nonpolar solvent?
(A) glucose
(B) graphite
(C) lithium fluoride
(D) sulfur
2. A solution of which substance can best be used as both a
titrant and its own indicator in an oxidation–reduction
titration?
(A) I2
(B) NaOCl
(C) K2Cr2O7
(D) KMnO4
(A) 0.166 M
(B) 0.180 M
(C) 0.333 M
(D) 0.666 M
7. When ionic hydrides react with water, the products are
(A) acidic solutions and hydrogen gas.
(B) acidic solutions and oxygen gas.
(C) basic solutions and hydrogen gas.
(D) basic solutions and oxygen gas.
36
8. 0.250 g of an element, M, reacts with excess fluorine to
produce 0.547 g of the hexafluoride, MF6. What is the
element?
32
T, °C
3. What value of
∆T should be
used for the
calorimetry
experiment that
gives these
graphed results?
6. A 20.00 mL sample of a Ba(OH)2 solution is titrated with
0.245 M HCl. If 27.15 mL of HCl is required, what is the
molarity of the Ba(OH)2 solution?
28
(A) Cr
(B) Mo
(C) S
(D) Te
24
20
0 2 4 6 8 10 12 14 16 18 20
time, min
(A) 10 ˚C
(B) 12 ˚C
(C) 15 ˚C
(D) 19 ˚C
Fe3+(aq) + SCN–(aq) ¾ FeSCN2+(aq)
4.
The equilibrium constant for this reaction can best be
determined by means of
(A) chromatography.
(B) conductance.
(C) ion exchange.
(D) spectrophotometry.
9. How many moles of Na + ions are in 20 mL of 0.40 M
Na 3PO4?
(A) 0.0080
(B) 0.024
(C) 0.050
(D) 0.20
10. What is the mass
percent of oxygen in
Al2(SO4)3·18H2O?
(A) 9.60
Molar Mass, M
Al2(SO4)3·18H2O 666.43 g·mol
(B) 28.8
(C) 43.2
(D) 72.0
11. What is the coefficient for H +(aq) when the equation is
balanced with whole number coefficients?
5. Which solid reacts with dilute hydrochloric acid at 25 ˚C
to produce a gas that is more dense than air?
(A) Zn
(B) Pb(NO3)2
(C) NaBr
(D) NaHCO3
__Mn2+(aq) + __BiO3–(aq) + __H+(aq) →
__Bi3+(aq) + __MnO4–(aq) + __H2O(l)
(A) 3
Not valid for use as an USNCO National Examination after April 26, 1999
(B) 4
(C) 7
(D) 14
Page 3
12. What is the number of O2 molecules in the 2.5 g of O 2
inhaled by the average person in one minute?
(A) 1.9 × 1022
(B) 3.8 × 1022
(C) 4.7 × 1022
(D) 9.4 × 1022
(A) –103 J
B
A
C
P, atm
13. Which point in the
phase diagram best
represents
supercritical
conditions?
19. What is the change in internal energy, ∆E, for a reaction
that gives off 65 joules of heat and does 38 joules of
work?
(B) B
(A)
D
(C) C
(B)
(C)
(A) 1 only
(B) 2 only
(C) 1 and 3 only
(D) 1, 2 and 3
15. What is the maximum number of phases that can be in
equilibrium in a one component system?
(A) 1
(B) 2
(C) 3
(D)
(D) D
14. The vapor pressure of a liquid in a closed container
depends on
1. temperature of the liquid
2. quantity of liquid
3. surface area of the liquid
(D) 4
(B) 20 g·mol–1
(C) 150 g·mol–1
(D) 190 g·mol–1
(A) CH3OCH3
(B) C 2H5OH
(C) CH3CH(OH)CH3
(D) CH2(OH)CH2OH
18. 3N2O(g) + 2NH3(g) → 4N2(g) + 3H2O(g) ∆H = –879.6 kJ
What is ∆Hf˚ for N2O in
Heats of Formation
kJ·mol–1?
NH3
–45.9 kJ·mol–1
H2O
–241.8 kJ·mol–1
(A) +246
(B) +82
(C) –82
–
–
+
+
(D) –246
–
+
+
–
(A) 1.3 × 10–7
(B) 2.0 × 10–7
(C) 3.0 × 10–7
(D) 4.5 × 10–7
22. Which statements are true?
1. S˚ values for all elements in their standard states
are positive.
2. S˚ values for all aqueous ions are positive.
3. ∆S˚ values for all spontaneous reactions are
positive.
(A) 1 only
(B) 1 and 2 only
(C) 2 and 3 only
(D) 1, 2 and 3
Ag+(aq) + 2 NH3(aq) ¾ Ag(NH3)2+(aq)
23.
17. Which substance would be expected to exhibit the
greatest surface tension at 25 ˚C?
(D) +103 J
21. The rate of formation of O3(g) is 2.0 × 10–7 mol·L–1·s–1 for
the reaction
3O2(g) → 2O3(g)
What is the rate of disappearance of O2(g) in mol·L–1·s–1?
16. The molar mass of a gas with a density of 5.8 g·L–1 at
25 ˚C and 740 mm Hg is closest to
(A) 10 g·mol–1
(C) +27 J
20. What are the signs of ∆H and ∆S for this reaction?
2C (s) + O2(g) → 2CO(g)
∆H
∆S
T, °C
(A) A
(B) –27 J
For this reaction, K = 1.7 × 107 at 25 ˚C. What is the
value of ∆G˚ in kJ?
(A) –41.2
(B) –17.9
(C) +17.9
(D) +41.2
24. The value of ∆H for a reaction can be found by
appropriate combination of bond enthalpies (the energy
required to break a particular bond, represented BE).
Which expression will give ∆H for this reaction?
C 2H4(g) + H2(g) → C2H6(g)
(A) BEC=C + BEH–H – [BEC–C + 2BEC–H]
(B) BEC–C + 2BEC–H – [BEC=C + BEH–H]
(C) 1 /2BEC=C + BEH–H – 2BEC–H
(D) 2BEC–H – 1/2BEC=C + BEH–H
Page 4
Not valid for use as an USNCO National Examination after April 26, 1999
25. What is the sign of ∆G˚ and the value of K for an
o
electrochemical cell for which Ecell
= 0.80 V?
(A)
(B)
(C)
(D)
K
Phosphorus reacts with chlorine as shown. What is the
equilibrium constant expression, K p, for this reaction?
–
+
+
–
>1
(A)
4 PPCl
(B)
3
6 PPCl ⋅ PCl
3
>1
(C)
<1
PPCl3
4 PPCl3
6 PCl2
2
(D)
PP4 ⋅ PCl6 2
4
PPCl
3
PCl6 2
<1
(B) 4
(C) 6
(D) 8
27. The decomposition of ethane into two methyl radicals
has a first order rate constant of 5.5 × 10–4 sec–1 at
700 ˚C. What is the half-life for this decomposition in
minutes?
(A) 9.1
P 4(s) + 6Cl2(g) ¾ 4PCl 3(g)
∆ G˚
26. The reaction between NO(g) and O2(g) to give NO2(g) is
second order in NO (g) and first order in O2(g). By what
factor will the reaction rate change if the concentrations
of both reactants are doubled?
(A) 2
31.
(B) 15
(C) 21
(D) 30
32. The equilibrium constant for the reaction
N2O4(g) ¾ 2NO2(g)
is 6.10 × 10–3 at 25˚C. Calculate the value of K for this
reaction:
NO2(g) ¾ 1 /2N2O4(g)
(A) 327
(B) 164
(C) 12.8
(D) 3.05 × 10–3
33. The ion-product constant for water at 45 ˚C is 4.0 × 10–14.
What is the pH of pure water at this temperature?
(A) 6.7
28. The dependence of the rate constant of a reaction on
temperature is given by the equation k = e – E kT .
Under what conditions is k the smallest?
a
(A) high T and large Ea
(B) high T and small Ea
(C) low T and large Ea
(D) low T and small Ea
(B) 7.0
(C) 7.3
(D) 13.4
34. The position of equilibrium lies to the right in each of
these reactions.
N2H5+ + NH3 ¾ NH4+ + N2H4
NH3 + HBr ¾ NH4+ + Br –
N2H4 + HBr ¾ N2H5+ + Br–
29. The reaction
CHCl3(g) + Cl2(g) → CCl 4(g) + HCl(g)
is believed to proceed by this mechanism:
Cl2(g) → 2Cl(g)
fast
Cl(g) + CHCl3(g) → HCl(g) + CCl3(g)
slow
CCl3(g) + Cl(g) → CCl 4(g)
fast
What rate equation is consistent with this mechanism?
(A) Rate = k[Cl2]
Based on this information, what is the order of acid
strength?
(A) HBr > N2H5+ > NH4+
(B) N2H5+ > N2H4 > NH4+
(C) NH3 > N2H4 > Br–
(D) N2H5+ > HBr > NH4+
35. HCN is a weak acid (K a = 6.2 × 10–10). NH3 is a weak
base (K b = 1.8 × 10–5). A 1.0 M solution of NH4CN
would be
(B) Rate = k[Cl][CHCl3]
(A) strongly acidic
(B) weakly acidic
(C) Rate = k[Cl2][CHCl3]
(C) neutral
(D) weakly basic
(D) Rate = k[Cl2]1/2[CHCl 3]
30. The activation energy of a certain reaction is 87 kJ·mol–1.
What is the ratio of the rate constants for this reaction
when the temperature is decreased from 37 ˚C to 15 ˚C?
(A) 5/1
(B) 8.3/1
(C) 13/1
36. What is the percent ionization of a 0.010 M HCN
solution? (Ka = 6.2 × 10–10)
(A) 0.0025%
(B) 0.025%
(C) 0.25%
(D) 2.5%
(D) 24/1
Not valid for use as an USNCO National Examination after April 25, 1999
Page 5
37. How many moles of HCOONa must be added to 1.0 L of
0.10 M HCOOH to prepare a buffer solution with a pH of
3.4? (HCOOH K a = 2 × 10–4)
(A) 0.01
(B) 0.05
(C) 0.1
(D) 0.2
38. The acid–base indicator methyl red has a Ka of 1 × 10–5.
Its acidic form is red while its alkaline form is yellow. If
methyl red is added to a colorless solution with a pH = 7,
the color will be
(A) pink
(B) red
(C) orange
(Ksp = 5.0 × 10–13)
(B) AgCl
(Ksp = 1.8 × 10–10)
(C) Ag2CO3
(Ksp = 8.1 × 10–12)
(D) Ag3AsO4
(Ksp = 1.0 × 10 )
43. Which expression gives the value for ∆G˚ in kJ·mol–1 for
this reaction at 25 ˚C?
(A) –6 × 8.31 × 0.43 × 1000
(B)
−6 × 96500 × 0.43 × 1000
8.31
(C)
−6 × 96500 × 0.43
1000
(D)
−6 × 8.31 × 0.43
1000
(D) yellow
39. Silver ions are added to a solution with
[Br–] = [Cl–] = [CO32–] = [AsO43–] = 0.1 M.
Which compound will precipitate at the lowest [Ag+]?
(A) AgBr
Questions 43. and 44. should be answered with reference to the
reaction.
2Cr(s) + 3Cu2+(aq) → 2Cr3+(aq) + 3Cu(s)
E˚ = 0.43 V
44. What is the voltage for this cell when [Cu2+] = 1.0 M and
[Cr3+] = 0.010 M?
(A) 1.2
(B) 0.87
(C) 0.47
(D) 0.39
–22
40. Consider a voltaic cell based on these half–cells.
Ag+(aq) + e– → Ag(s)
E˚ = +0.80 V
2+
–
Cd (aq) + 2e → Cd (s)
E˚ = –0.40 V
Identify the anode and give the voltage of this cell under
standard conditions.
(A) Ag; Ecell = 0.40 V
(B) Ag; Ecell = 2.00 V
(C) Cd; Ecell = 1.20 V
(D) Cd; Ecell = 2.00 V
45. All of these sets of quantum numbers are permissible
except
n
l
ml
ms
(A)
1
0
0
+ 1 /2
(B)
2
2
0
–1 /2
(C)
3
1
1
–1 /2
(D)
3
2
–1
+ 1 /2
46. Which element can exhibit more than one oxidation state
in compounds?
1. Cr
2. Pb
3. Sr
41. Which two species react spontaneously?
(A) Cu (s) + Ag+(aq)
(B) Br2(l) + Cl–(aq)
(A) 1 only
(B) 1 and 2 only
(C) H2O(l) + Ca2+(aq)
(D) Au(s) + Mg2+(aq)
(C) 2 and 3 only
(D) 1, 2 and 3
42. When aluminum oxide is electrolyzed in the industrial
process for the production of aluminum metal, aluminum
is produced at one electrode and oxygen gas is produced
at the other. For a given quantity of electricity, what is
the ratio of moles of aluminum to moles of oxygen gas?
(A) 1:1
(B) 2:1
(C) 2:3
(D) 4:3
47. When the isoelectronic species, K+, Ca2+, and Cl–, are
arranged in order of increasing radius, what is the correct
order?
(A) K+, Ca2+, Cl–
(B) K+, Cl–, Ca2+
(C) Cl–, Ca2+,K+
(D) Ca2+, K+, Cl–
48. Which Group 2 element has chemical properties least
like the other members of the group?
(A) Be
Page 6
(B) Ca
(C) Sr
(D) Ba
Not valid for use as an USNCO National Examination after April 26, 1999
49. In the vapor state which atom has the largest ionization
energy?
(A) Na
(B) K
(C) Mg
(D) Ca
50. All of these species have the same number of valence
electrons as NO3– except
(A) CO32–
(B) HCO3–
(C) NF3
(D) SO3
51. Which set contains no ionic species?
(A) NH4Cl, OF2, H2S
(B) CO2, Cl2, CCl4
(C) BF3, AlF3, TlF3
(D) I2, CaO, CH3Cl
56. How many carbon–carbon bonds are in a molecule of
2-methyl-2-butanol?
(A) 2
(B) 3
(C) 4
(D) 5
57. Which molecule can exist as stereoisomers?
(A) CHF=CHF
(B) F 2C=CCl2
(C) CH2F–CHF2
(D) CF3–CH3
58. What are the most likely products in the reaction between
CH3CH2CH2OH and HI?
(A) CH3CH2CH2I and H2O
(B) CH3CH2CH3 and HOI
52. When these species are arranged in order of increasing
bond energy, what is the correct sequence?
(A) N2, O2, F2
(B) F 2, O2, N2
(C) O2, F2, N2
(D) O2, N2, F2
(C) CH3OH and CH3CH2I
(D) ICH2CH2CH2OH and H2
59. Addition polymers include
1. polyamide 2. polyethylene
53. The geometry of the atoms in the species PCl 4+ is best
described as
(A) tetrahedral
(B) see–saw
(C) square
(D) trigonal bipyramidal
(A) 1 only
(B) 2 only
(C) 2 and 3 only
(D) 1, 2 and 3
60. All of these are aromatic compounds except
(A) hexene, C6H12
54. Which are nonpolar molecules?
1. NCl3
2. SO3
3. PCl 5
(B) toluene, C6H5CH3
(A) 1 only
(B) 2 only
(C) p-dichlorobenzene, C6H4Cl2
(C) 1 and 3 only
(D) 2 and 3 only
(D) naphthalene, C10H8
55. What are the hybridizations of
carbon 1 and carbon 2 in the
hydrocarbon?
3. polyester
CH3CHCH2
1
(A) sp3, sp
(B) sp3, sp2
(C) sp2, sp2
(D) sp, sp2
2
END OF TEST
Not valid for use as an USNCO National Examination after April 25, 1999
Page 7
Page 8
Not valid for use as an USNCO National Examination after April 26, 1999
US National Chemistry Olympiad – 1999
National Examination—Part I
SCORING KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Answer
D
D
B
D
D
A
C
B
B
D
D
C
B
A
C
C
D
B
A
B
Number
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
Answer
C
A
A
A
A
D
C
C
D
C
D
C
A
A
D
B
B
D
A
C
Property of the ACS Society Committee on Education
Number
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
A
D
C
C
B
B
D
A
C
C
B
B
A
D
B
C
A
A
B
A
1999 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Lucy Pryde Eubanks, Clemson University, Clemson, SC
Chair
Robert Becker, Kirkwood High School, Kirkwood, MO
Craig W. Bowen, Clemson University, Clemson, SC
J. Emory Howell, University of Southern Mississippi, Hattiesburg, MS
Sheldon L. Knoespel, Michigan State University, East Lansing, MI
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Robert G. Silberman, SUNY-Cortland, NY
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 75 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 questions
laboratory practical
1 hour, 15 minutes
A periodic table is provided on page 8 for reference. Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
DETAILED DIRECTIONS CAREFULLY BEFORE YOU PROCEED. NOTE THE PERIODIC TABLE ON PAGE 8.
There are two laboratory-related tasks for you to complete during the next 75 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 75 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 26, 1999
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC
1999 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
You have been given a sample of 7-Up® that has been allowed to stand open. You also have been
provided with some table sugar (sucrose), distilled or deionized water, some measuring devices, and a variety of
containers. Graph paper has been provided on page 5 of this test booklet. Devise and carry out an experiment to
determine the percent by mass of sugar in a sample of 7-Up. You will be asked to describe the method you
developed to solve this problem.
Given:
The molar mass of sucrose, C12H22O11, is 342.30 g·mol–1 .
Lab Problem 2
You have been given a sample of Crystal Drano®. There are two components in the Drano – some small shiny
metallic pieces, and some pale green beads. (The green color is a dessicating substance.) The metallic pieces are
either zinc, magnesium, or aluminum. The beads are either NaOH, Ca(OH)2, or Al(OH)3. You also have 1.0 M
NaOH, 3.0 M HCl, and some phenolphthalein indicator. Devise and carry out an experiment to identify both
components of Crystal Drano. You will be asked to describe the method you developed to solve this problem.
Special Safety Consideration: Crystal Drano is quite caustic and must only be handled with the scoops or
spatulas provided. Also, Crystal Drano will readily absorb moisture from the air so only open the container
when you need a sample. Recap the container as quickly as possible.
Page 2
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the
steps you plan to take to solve this problem.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
2. Record your data and other observations.
Page 3
Examiner’s Initials:
3. Calculate the percent by mass of sugar in 7-Up. You may choose to use the graph paper on the next page.
Show your methods clearly.
Percent by mass sucrose in 7-Up®
4. Explain what assumptions were made in determining the mass percent of sugar in 7-Up. How does each
assumption influence the mass percent you calculated?
Page 4
Graph Paper for Possible Use with Laboratory Practical Problem 1
Page 5
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name : ________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the
steps you plan to take to solve this problem.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
2. Record your data and other observations.
Page 6
Examiner’s Initials:
3. Identify the two components of Crystal Drano. Support your choices with conclusions drawn from your
observations.
Identification of metallic pieces:
____________________________________
Identification of beads:
____________________________________
Page 7
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(289)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Page 8
1999 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
ANSWER KEYS
Lab Problem 1
5 pts
Plan is expected to include:
• method to measure density of 7-Up®
• method for determining the composition of sugar in water.
• method to compare 7-Up solution to the known sugar/water mixtures.
• replications of values of data expected.
8 pts
Analysis of sugar/water solutions is expected to include:
• data to support the method chosen for determining the composition of the sugar/water solutions.
• multiple samples.
• volume changes.
• well-organized data tables .
6 pts
Analysis of 7-Up solution is expected to include:
• data to support the method chosen for determining the composition of the 7-Up.
• replications of samples.
4 pts
Results are expected to include:
• clear explanations of calculations used to evaluate data.
• reasonable results.
Example of a reasonable approach (other approaches were evaluated as possibly acceptable):
• Determine mass and volume of samples of 7-Up from which an average density value can be determined.
• Determine mass and volume of a wide variety of sugar/water mixtures. Use these data to determine the
density and corresponding mass % solute values of these mixtures.
• Determine mass a volume of a wide variety of sugar/water mixtures. Use these data to determine the density and
corresponding mass % solute for sugar/water mixtures.
• Having the calculated average density value for 7_Up, use the graph to determine the mass % of 7-Up.
• Data points on graph should bracket the 7-Up value.
• Graph paper should be used efficiently.
2 pts
Discussion of assumptions made should include:
• recognition that other components may be present in the 7-Up other than sucrose.
• limitations inherent in the methodology and equipment.
Page 9
Lab Problem 2
5 pts
Plan is expected to include:
• method to separate the metal and green beads in Crystal Draino®.
• method for testing the metal with water, HCl, and NaOH.
• method to testing the green beads with water.
• replications of tests.
8 pts
Observations are expected to include that:
• the green beads dissolve in water.
• the metal turns black in water.
• the metal reacts with both HCl and NaoH, forming bubbles.
• well-organized observation tables .
12 pts Identification and support section is expected to include:
• Reasoning for selecting NaOH.
• The green beads dissolved easily in water, so the beads must be NaOH. Neither Ca(OH)2 nor Al(OH)3 are
soluble in water.
• Using 3 M HCl, a titration can be done to determine the amount of hydroxide in the beads.
• Reasoning for selecting Al
• All three possible metals would react with 3M HCl, so qualitative observations of reaction with HCl are
not definitive.
• Only Zn and Al react with 1 M NaOH. Because the metal is observed to react with NaOH, the Mg can be
eliminated.
• To distinguish between Zn and Al, a quantitative titration can be done with 1 M NaOH or with the
3M HCl.
Page 10
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
Edward DeVillafranca (retired), Kent School, CT
Alice Johnsen, Bellaire High School, TX
John A. Krikau (retired), Lyons Township High School, IL
Patricia A. Metz, University of Georgia, GA
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
Diane D. Wolff, Western Virginia Community College, VA
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 16, 2000, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number as you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement which is followed by four possible choices. Select the single choice that best answers the question or completes
the statement. Then use a pencil to blacken the space on your answer sheet having the same letter as your choice. You may write on
the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 16, 2000.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K milli- prefix
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G mole
A frequency
ν Planck’s constant
N A gas constant
R pressure
°C gram
g rate constant
c hour
h retardation factor
C joule
J second
E kelvin
K speed of light
Ea kilo- prefix
k temperature, K
H liter
L time
S measure of pressure mmHg volt
CONSTANTS
m
m
M
mol
h
P
k
Rf
s
c
T
t
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as a USNCO National Exam after April 16, 2000.
DIRECTIONS
When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
Make no marks on the test booklet. Do all calculations on scratch paper provided by your instructor.
There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
(A) I only
(B) III only
(C) I and II only
(D) I, II, and III
2. Which substance is stored in contact with water to
prevent it from reacting with air?
(A) bromine
(B) lithium
(C) mercury
(D) phosphorus
6. The molarity of a Cu2+
solution is to be
determined from its
absorbance, measured
under the same conditions
as those used to prepare
this calibration curve.
What will be the percent
uncertainty in the
concentration of a 0.050 M
solution if the uncertainty
in the absorbance reading
is ±0.01 absorbance units?
(A) 5%
3. A solution of concentrated aqueous ammonia is added
dropwise to 1 mL of a dilute aqueous solution of
copper(II) nitrate until a total of 1 mL of the ammonia
solution has been added. What observations can be made
during this process?
(A) The colorless copper(II) nitrate solution turns blue
and yields a dark blue precipitate.
(B) The colorless copper(II) nitrate solution yields a
white precipitate which turns dark blue upon
standing.
(C) The light blue copper(II) nitrate solution yields a
precipitate which redissolves to form a dark blue
solution.
(D) The light blue copper(II) nitrate solution turns dark
blue and yields a dark blue precipitate.
4. What gas is produced when dilute HNO3 is added to
silver metal?
(A) NO
(B) H2
(C) NH3
(D) N2
5. A substance is analyzed by paper chromatography,
giving the chromatogram shown.
start
solvent front
0.0
2.0
4.0
6.0
8.0
10.0
(B) 10%
0.40
Absorbance
1. Which of these ions is expected to be colored in aqueous
solution?
I Fe3+
II Ni2+
III Al3+
0.30
0.20
0.10
0
0
0.05 0.10 0.15
[Cu2+], M
(C) 15%
(D) 20%
7. A 1.50 g sample of an ore containing silver was
dissolved, and all of the Ag+ was converted to 0.124 g of
Ag2S. What was the percentage of silver in the ore?
(A) 6.41%
(B) 7.20%
(C) 8.27%
(D) 10.8%
8. Methyl-t-butyl ether, C5H12O, is added to gasoline to
promote cleaner burning. How many moles of oxygen
gas, O 2, are required to burn 1.0 mol of this compound
completely to form carbon dioxide and water?
(A) 4.5 mol
(B) 6.0 mol
(C) 7.5 mol
(D) 8.0 mol
9. A 0.200 g sample of
benzoic acid, C6H5COOH, Substance Molar Mass
is titrated with a 0.120 M
C 6H5COOH 122.1 g·mol–1
Ba(OH)2 solution. What
volume of the Ba(OH)2 solution is required to reach the
equivalence point?
(A) 6.82 mL
(B) 13.6 mL
(C) 17.6 mL
(D) 35.2 mL
12.0 cm
What is the Rf value of the substance represented by the
spot at 8.0 cm?
(A) 0.80
(B) 0.75
(C) 0.67
(D) 0.60
Not valid for use as an USNCO National Examination after April 16, 2000.
Page 3
(A) 5.15 g
(B) 14.3 g
(C) 19.4 g
(D) 26.4 g
11. What is the Na+ ion concentration in the solution formed
by mixing 20. mL of 0.10 M Na2SO4 solution with
50. mL of 0.30 M Na3PO4 solution?
(A) 0.15 M
(B) 0.24 M
(C) 0.48 M
(D) 0.70 M
12. A solution prepared by
Compound
Kb
dissolving a 2.50 g sample
C 6H6
2.53 °C·m–1
of an unknown compound
dissolved in 34.0 g of benzene, C6H6, boils 1.38 °C
higher than pure benzene. Which expression gives the
molar mass of the unknown compound?
2.50
(A) 2.53 ×
1.38
(B) 1.38 ×
2.53
1
×
34.0 1.38
3
(D) 2.50 × 10 ×
1.38
× 2.53
34.0
740 mmHg + 20 mmHg
740 mmHg
(C) 300 mL ×
740 mmHg
740 mmHg – 20 mmHg
(D) 300 mL ×
740 mmHg
740 mmHg + 20 mmHg
17. What is the normal melting
point of the substance
represented by the phase
diagram?
1.0
A B C
T, °C
(A) A
13. What is the total pressure in a 2.00 L container that holds
1.00 g He, 14.0 g CO, and 10.0 g of NO at 27.0 °C?
(A) 21.6 atm
(B) 13.2 atm
(C) 1.24 atm
(D) 0.310 atm
14. What type of solid is generally characterized by having
low melting point and low electrical conductivity?
(A) ionic
(B) metallic
(C) molecular
(D) network covalent
15. How many nearest neighbors surround each particle in a
face-centered cubic lattice?
Page 4
(B) 300 mL ×
34.0
× 2.50
2.53
3
(C) 2.50 × 10 ×
(A) 4
16. Hydrogen is
Compound
Vapor Pressure
collected over water
at 22 °C
at 22 °C and a
H2O
20. mmHg
barometer reading
of 740 mmHg. If 300. mL of hydrogen is collected,
which expression will give the volume of dry hydrogen
at the same temperature and pressure?
740 mmHg – 20 mmHg
(A) 300 mL ×
740 mmHg
P,atm
10. Chlorine can be prepared by reacting HCl with MnO2.
The reaction is represented by this equation.
MnO2(s) + 4HCl(aq) → Cl2(g) + MnCl2(aq) + 2H2O(l)
Assuming the reaction goes to completion what mass of
concentrated HCl solution (36.0% HCl by mass) is
needed to produce 2.50 g of Cl 2?
(B) 6
(C) 8
(D) 12
(B) B
(C) C
D
(D) D
18. A bomb calorimeter has a heat capacity of 783 J·°C–1
and contains 254 g of water, which has a specific heat of
4.184 J·g–1·°C–1. How much heat is evolved or absorbed
by a reaction when the temperature goes from 23.73 °C
to 26.01 °C?
(A) 1.78 kJ absorbed
(B) 2.42 kJ absorbed
(C) 1.78 kJ evolved
(D) 4.21 kJ evolved
19. Consider this equation and the associated value for ∆Ho.
2H2(g) + 2Cl2(g) → 4HCl(g)
∆Ho = –92.3 kJ
Which statement about this information is incorrect?
(A) If the equation is reversed, the ∆Ho value equals
+92.3 kJ.
(B) The four HCl bonds are stronger than the four bonds
in H2 and Cl2.
(C) The ∆Ho value will be –92.3 kJ if the HCl is
produced as a liquid.
(D) 23.1 kJ of heat will be evolved when 1 mol of
HCl (g) is produced.
Not valid for use as an USNCO National Examination after April 16, 2000.
(A) –1074.0 kJ
(B) –22.2 kJ
(C) +249.8 kJ
(D) +2214.6 kJ
25. A reaction follows this
concentration-time
diagram. The instantaneous
rate for this reaction at 20
seconds will be closest to
which value?
0.40
Molarity
20. Determine the heat of reaction for this process.
FeO(s) + Fe2O3(s) → Fe3O4(s)
Given information:
2Fe(s) + O2(g) → 2FeO(s)
∆Ho = –544.0 kJ
4Fe(s) + 3O2(g) → 2Fe2O3(s)
∆Ho = –1648.4 kJ
Fe3O4(s) → 3Fe(s) + 2O2(g)
∆Ho = +1118.4 kJ
0.30
0.20
0.10
0
0
21. For which process will ∆Ho and ∆Go be expected to be
most similar?
(A) 2Al(s) + Fe2O3(s) → 2Fe(s) + Al2O3(s)
(A) 4 × 10–3 M·sec–1
(B) 8 × 10–3 M·sec–1
(C) 2 × 10–2 M·sec–1
(D) 1 × 10–1 M·sec–1
(C) 2NO2(g) → N2O4(g)
(A) zero.
(B) first.
(D) 2H2(g) + O2(g) → 2H2O(g)
(C) second.
(D) third.
Bond
Bond Energy
H–H
O–O
O=O
H–O
436
142
499
460
(A) –127 kJ
(B) –209 kJ
(C) –484 kJ
(D) –841 kJ
60
26. If the half-life of a reaction increases as the initial
concentration of substance increases, the order of the
reaction is
(B) 2Na (s) + 2H2O(l) → 2NaOH (aq) + H2(g)
22. Use bond energies to
estimate ∆H for this
reaction.
H2(g) + O2(g) → H2O2(g)
20 40
Time, sec
kJ·mol–1
kJ·mol–1
kJ·mol–1
kJ·mol–1
23. For a particular reaction, ∆Ho = –38.3 kJ and
∆So = –113 J·K–1. This reaction is
(A) spontaneous at all temperatures.
(B) nonspontaneous at all temperatures.
27. The radioisotope N-13, which has a half-life of
10 minutes, is used to image organs in the body. If
an injected sample has an activity of 40 microcuries
(40 µCi), what is its activity after 25 minutes in the
body?
(A) 0.75 µCi
(B) 3.5 µCi
(C) 7.1 µCi
(D) 12 µCi
28. Propanone reacts with iodine in acid solution as shown in
this equation.
H+
CH3C(O)CH3 + I2  
→ CH3C(O)CH2I + HI
These data were obtained when the reaction was studied.
[CH3C(O)CH3], M [I2], M [H+], M Relative Rate
0.010
0.020
0.020
0.020
(C) spontaneous at temperatures below 66 °C.
(D) spontaneous at temperatures above 66 °C.
24. What is ∆Go for this reaction?
1/2N2(g) + 3/2H2(g) = NH3(g)
K p = 4.42 × 104 at 25 °C.
0.010
0.010
0.020
0.010
0.010
0.010
0.010
0.020
1
2
2
4
What is the rate equation for the reaction?
(A) rate = k[CH3C(O)CH3] [I2]
(A) –26.5 kJ·mol–1
(B) –11.5 kJ·mol–1
(B) rate = k[CH3C(O)CH3]2
(C) –2.2 kJ·mol–1
(D) –0.97 kJ·mol–1
(C) rate = k[CH3C(O)CH3] [I2] [H+]
(D) rate = k[CH3C(O)CH3] [H+]
29. A particular reaction rate increases by a factor of five
when the temperature is increased from 5 °C to 27 °C.
What is the activation energy of the reaction?
(A) 6.10 kJ·mol–1
(B) 18.9 kJ·mol–1
(C) 50.7 kJ·mol–1
(D) 157 kJ·mol–1
Not valid for use as an USNCO National Examination after April 16, 2000.
Page 5
30. Consider this reaction.
2H2(g) + 2NO(g) → N2(g) + 2H2O(g)
The rate law for this reaction is rate = k [H2] [NO]2.
Under what conditions could these steps represent the
mechanism?
Step 1.
2NO = N2O2
Step 2.
N 2O2 + H 2 → N 2O + H 2O
Step 3.
N 2O + H 2 → N 2 + H 2O
(A) These steps cannot be the mechanism under any
circumstances.
(B) These steps could be the mechanism if step 1 is the
slow step.
(C) These steps could be the mechanism if step 2 is the
slow step.
(D) These steps could be the mechanism if step 3 is the
slow step.
31.
A reaction has a forward rate constant of 2.3 × 106 s–1
and an equilibrium constant of 4.0 × 108. What is the rate
constant for the reverse reaction?
(A) 1.l × 10
–15
(C) 1.7 × 10 s
2
s
(B) 5.8 × 10 s
–1
–3
–1
(D) 9.2 × 1014 s–1
–1
32. For the reaction 2A(g) + 2B (g) = 3C(g) at a certain
temperature, K is 2.5 × 10–2. For which conditions will
the reaction proceed to the right at the same temperature?
[A], M
[B], M
[C], M
36. What is the conjugate acid of HPO42–?
(A) H3PO4(aq)
(B) H2PO4–(aq)
(C) H3O+(aq)
(D) PO43–(aq)
37. The amount of sodium
Acid
Ka
hydrogen carbonate,
H2CO3
2.5 × 10–4
NaHCO3, in an antacid
–
HCO3
2.4 × 10–8
tablet is to be determined
by dissolving the tablet in water and titrating the
resulting solution with hydrochloric acid. Which
indicator is the most appropriate for this titration?
(A) methyl orange, pKin = 3.7
(B) bromothymol blue, pKin = 7.0
(C) phenolphthalein, pKin = 9.3
(D) alizarin yellow, pK in = 12.5
38. How many moles of
Acid
Ka
NaOCl must be added to
HOCl
2.8 × 10–8
150 mL of 0.025 M HOCl
to obtain a buffer solution with a pH = 7.50?
(A) 2.6 × 10–5
(B) 1.1 × 10–3
(C) 3.3 × 10–3
(D) 2.2 × 10–2
39. If equal volumes of BaCl2
Substance K sp
and NaF solutions are
BaF 2
1.7 × 10–7
mixed, which of these
combinations will not give a precipitate?
(A)
0.10
0.10
0.10
(B)
1.0
1.0
1.0
(C)
1.0
0.10
0.10
(A) 0.0040 M BaCl2 and 0.020 M NaF
(D)
1.0
1.0
0.10
(B) 0.010 M BaCl2 and 0.015 M NaF
(C) 0.015 M BaCl2 and 0.010 M NaF
–
33. What is the Kb of a weak base that produces one OH per
molecule if a 0.050 M solution is 2.5% ionized?
(A) 7.8 × 10–8
(B) 1.6 × 10–6
(C) 3.2 × 10–5
(D) 1.2 × 10–3
34. What is the [OH–] of a
0.65 M solution of
NaOCl?
Acid
HOCl
Ka
2.8 × 10–8
(D) 0.020 M BaCl2 and 0.0020 M NaF
40. What takes place when zinc metal is added to a aqueous
solution containing magnesium nitrate and silver nitrate?
1. Zn is oxidized.
2. Mg2+ is reduced.
3. Ag+ is reduced.
4. No reaction takes place.
(A) 4.8 × 10–4 M
(B) 1.3 × 10–4 M
(A) 1 and 2 only
(B) 1 and 3 only
(C) 3.5 × 10–7 M
(D) 2.1 × 10–11 M
(C) 1, 2, and 3 only
(D) 4 only
35. Which acid is the strongest?
(A) H3BO3
(B) H3PO4
(C) H2SO3
(D) HClO3
Page 6
Not valid for use as an USNCO National Examination after April 16, 2000.
Questions 41, 42, and 43 should be answered with reference to
this information and diagram.
Ag+(aq) + e– → Ag(s)
Eo = 0.80 V
2+
–
Cu (aq) + 2e → Cu(s)
Eo = 0.34 V
45. How many unpaired electrons are in a gaseous Fe2+ ion in
the ground state?
(A) 0
(B) 2
(C) 4
(D) 6
46. Which element has the smallest first–ionization energy?
V
(A) Mg
(B) Al
(C) Si
(D) P
salt bridge
Ag
Cu
Ag+ (aq)
Cu2+ (aq)
47. Which set of orbitals is listed in the sequential order of
filling in a many-electron atom?
(A) 3s, 3p, 3d
(B) 3d, 4s, 4p
(C) 3d, 4p, 5s
(D) 4p, 4d, 5s
48. Which set is expected to show the smallest difference in
first–ionization energy?
41. What is the value for ∆G° when [Ag+] = [Cu2+] = 1.0 M?
(A) –44.4 kJ
(B) –88.8 kJ
(C) –243 kJ
(D) –374 kJ
42. Which expression gives the voltage for this cell if
[Cu2+] = 1.00 M and [Ag+] = 0.010 M?
(A) 0.46 V + 0.0591 V
(A) He, Ne, Ar
+
(B) B, N, O
2+
(C) Mg, Mg , Mg
(D) Fe, Co, Ni
49. When the atoms Li, Be, B, and Na are arranged in order
of increasing atomic radius, what is the correct order?
(A) B, Be, Li, Na
(B) Li, Be, B, Na
(C) Be, Li, B, Na
(D) Be, B, Li, Na
(B) 0.46 V + 2 × 0.0591 V
50. Which species has the same shape as the NO3– ion?
(C) 0.46 V – 0.0591 V
(D) 0.46 V – 2 × 0.0591 V
43. Which increases immediately if the surface area of the
silver electrode is increased?
(A) overall cell voltage
(A) SO3
(B) SO32–
(C) ClF3
(D) ClO 3–
51. What is the formal charge on the
central atom in N2O?
N N O
+
(B) rate of change of [Ag ]
(C) mass of Cu electrode
 mass of Cu 
(D) change in ratio of electrode masses; ∆ 

 mass of Ag 
44. In the galvanizing process, iron is coated with zinc. The
resulting chemical protection is most similar to that
provided when
(A) a magnesium bar is connected to an iron pipe.
(B) an iron can is plated with tin.
(A) +1
(B) 0
(C) –1
(D) –2
52. How many bonding pairs and lone pairs surround the
central atom in the I 3– ion?
Bonding Pairs
Lone Pairs
(A)
2
2
(B)
2
3
(C)
3
2
(D)
4
3
(C) copper pipes are connected using lead solder.
(D) a copper pipe is covered with epoxy paint.
Not valid for use as an USNCO National Examination after April 16, 2000.
Page 7
53. The nitrogen atoms in NH3, NH2–, and NH4+ are all
surrounded by eight electrons. When these three species
are arranged in order of increasing H–N–H bond angle,
what is the correct order?
–
+
(B) NH4 , NH2 , NH3
+
–
(D) NH2–, NH3, NH4+
(A) NH3, NH2 , NH4
(C) NH3, NH4 , NH2
+
57. Which is the formula for an alkyne?
(A) C 2H4
(B) C 3H6
(C) C 3H8
(D) C 4H6
–
58. How many isomers have the formula C3H8O?
(A) 2
54. What hybrid orbitals are
employed by carbon atoms
1,2, and 3, respectively, as
labeled in the compound
shown?
3
O
H3C C C N
1 2 3
(A) sp , sp, sp
2
2
(B) sp , sp , sp
3
3
2
2
(C) sp , sp , sp
(D) sp , sp , sp
(C) 4
(D) 5
59. Which type of organic compound is most resistant to
oxidation by acidified potassium dichromate?
(A) acid
(B) alcohol
(C) aldehyde
(D) alkene
2
55. In which pair, or pairs, is the stronger bond found in the
first species?
1. O2–, O2
2. N2, N2+
3. NO+, NO–
(A) 1 only
(B) 2 only
(C) 1 and 3 only
(D) 2 and 3 only
60. What product, in addition to water, is produced by this
reaction?
CH3OH + C6H5COOH →
O
C OH
(A)
H3C
(B)
56. What is the molecular
formula of this chemical
structure?
(B) 3
CH3
O
C CH3
(C)
CH3
(A) C 10H12
(B) C 10H14
(C) C 12H12
(D) C 12H14
CH3
(D)
O
C O CH3
END OF TEST
Page 8
Not valid for use as an USNCO National Examination after April 16, 2000.
US National Chemistry Olympiad – 2000
National Examination—Part I
SCORING KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Answer
C
D
C
A
B
B
B
C
A
B
D
C
B
C
D
A
B
D
C
B
Number
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
Answer
A
A
C
A
A
A
C
D
C
C
B
D
C
A
D
B
A
C
D
B
Property of the ACS Society Committee on Education
Number
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
B
D
B
A
C
B
C
D
A
A
A
B
D
C
D
C
D
B
A
D
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
Edward DeVillafranca (retired), Kent School, CT
Alice Johnsen, Bellaire High School, TX
John A. Krikau (retired), Lyons Township High School, IL
Patricia A. Metz, University of Georgia, GA
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
Diane D. Wolff, Western Virginia Community College, VA
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a
score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 16, 2000, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front
of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I. Answer all of the questions in order, and use both sides of
the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper,
and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 16, 2000.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
1.
(12%) An unknown metal, M, reacts with excess chlorine to give the metal chloride, MClx. When 0.396 g of the
chloride is dissolved in water and passed through an anion exchange column charged with hydroxide ions, the
solution requires 23.55 mL of 0.195 M HCl for neutralization.
a. Calculate the number of moles of HCl used in the titration.
b. Determine the mass of chlorine and the mass of metal in this sample of MClx.
c. Assuming that x in MClx is 1, 2 or 3, calculate possible atomic masses for M.
d. Use your knowledge of the Periodic Table to write formulas for the possible compounds between chlorine and
metals and identify those expected to be stable.
2.
(11%) The ionization constant for water is 1.14 × 10–15 at 0 °C and 9.6 × 10–14 at 60 °C.
a. Write the equation for the ionization of water and determine the pH of water at 60 °C.
b. Calculate each value.
i. ∆Hionization over this temperature range
ii. ∆G at 60 °C
iii. ∆S at 60 °C
c. State the significance of the sign of the sign of ∆S obtained in part 2b.iii, and explain how the process indicated
in 2a could lead to this sign.
3.
(15%) These are the reaction steps in a certain polymerization process, which may occur by either an uncatalyzed
or an acid-catalyzed pathway.
O
R
C
O
O
+ HA
H
k1
C+
R
k2
O
Group 1
O
R
A–
H
O
H
R
C
O
R'
O+
H
C+
O
H
A–
+
R'
O
k3
H
k4
H
Group 2
H
A–
H
O
R
C
R'
O+
H
O
O
H
k5
R
C
O
R'
+ H2O
+ HA
A–
Group 3
H
a. Write a balanced equation for the overall reaction.
b. Name the functional groups labeled [1], [2], and [3].
c. Given these data for the acid-catalyzed reaction, find the rate law and the value of k, specifying its units.
[RCOOH], M
[R´OH], M
[HA], M
Initial Rate, M·min –1
0.35
0.35
0.50
4.60
0.62
0.35
0.50
8.14
0.35
0.81
0.50
10.6
0.35
0.50
0.75
9.84
d. Identify the rate-determining step based on the rate law found in question 3c. Explain your answer.
e. The initial reaction rate can be followed spectrophotometrically by quenching the reaction and determining
the amount of ROH left by its reaction with dichromate ion, Cr2O72–.
i. Write a balanced equation for the reaction of Cr2O72– with R´OH in acid solution. Assume R´ is CH3CH2–
and the products of the reaction are Cr 3+ and CH3COOH.
ii. Describe the color change expected for the reaction written in question 3e, part i.
Page 2
Not valid for use as an USNCO National Examination after April 16, 2000.
4.
(12%) 25.00 mL of a solution of a weak monoprotic acid, HX, was titrated with a 0.0640 M solution of NaOH,
requiring 18.22 mL. The pH of the solution varied as a function of the percentage of HX titrated. These data
were collected.
% titrated
0
33.3%
66.7%
pH
3.39
5.14
5.74
a. Calculate the initial concentration of the weak acid in the 25.00 mL of solution.
b. Determine the value of K a for two of these three conditions.
c. Calculate the pH at the equivalence point of this titration and write an equation to account for this pH.
d. Calculate the number of moles of a salt, NaX, that must be added to produce a pH of 6.00 in 150.00 mL of the
original solution.
5.
(14%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the
reactants and products and omit formulas for all ions or molecules that do not take part in a reaction. Write
structural formulas for all organic substances. You need not balance the reactions. All reactions occur in aqueous
solution unless otherwise indicated.
a. Phosphorus is burned in excess oxygen.
b. Sulfur dioxide is bubbled into water.
c. Chlorine gas is bubbled through a sodium bromide solution.
d. Solutions of magnesium nitrate and potassium hydroxide are mixed.
e. A sodium thiosulfate solution is added to a suspension of silver chloride.
f. Bromine is added to a solution of ethylene in hexane.
g. Radium-226 emits an alpha particle.
7.
8.
(12%) Use the given phase diagram of water to answer these
questions. Note that the axis values are not drawn to scale.
a. Identify the physical state at points A, B, C, and D.
b. Calculate the volume of one mole of water in each of the phases
at the triple point, (At the triple point, the density of H2O(l) is
0.9998 g·mL–1 and the density of H2O(s) is 0.917 g·mL–1.)
c. Starting with point A, describe the pressure, temperature, and
phase changes that correspond to the rectangle around the triple
point.
P, mmHg
6. (12%) Nitrogen dioxide, NO2, can undergo reactions to form nitrite ion, NO2– , and nitronium ion, NO2+.
a. Draw Lewis structures for NO 2– and NO2+ including any resonance forms.
b. Predict the shape of each ion and account for each shape using a modern bonding theory.
c. Describe and account for the difference in the N–O bond lengths in NO2– and NO2+.
d. Determine the oxidation number and the formal charge of nitrogen in the NO2– ion. Outline your reasoning
and state the difference between formal charge and oxidation number.
760
4.58
D
A
C
B
0
0.01
T, °C
100
(12%) The behavior of elements can often be predicted based on their positions in the Periodic Table. Use your
knowledge about trends in the behavior of elements to answer the following questions about the recently isolated
elements 114, 116 and 118.
a. Give the names and symbols of the elements in the row above 114, 116, and 118 in the Periodic Table.
b. Predict the relative ionization energies of elements 114, 116, and 118 and describe how the ionization energy
of one of them is expected to compare with the ionization energy of the element above it, giving reasons for
your answers.
c. Predict the oxidation states expected for element 114 and indicate which oxidation state is expected to be most
stable, giving reasons for your answers.
d. Suggest a reason that elements 114, 116, and 118 have been made, but elements 113, 115, and 117 have not.
END OF PART II
Not valid for use as an USNCO National Examination after April 16, 2000.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K milli- prefix
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G mole
A frequency
ν Planck’s constant
N A gas constant
R pressure
°C gram
g rate constant
c hour
h second
C joule
J speed of light
E kelvin
K temperature, K
Ea kilo- prefix
k time
H liter
L volt
S measure of pressure mmHg volume
CONSTANTS
m
m
M
mol
h
P
k
s
c
T
t
V
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol –1
N A = 6.022 × 1023 mol –1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
– ∆H   1 
ln K = 
+c
 R  T
RT
E=E –
ln Q
nF
ο
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
Page 4
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 16, 2000.
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART II
1.
a.
b.
c.
2.
a.
0.195 mol
= 0.00459 mol H + = 0.00459 mol Cl –
L
35.45 g
0.00459 mol Cl – ×
= 0.163 g Cl –
mol
1
50.8 g
0.396 g MCl - 0.163 g Cl – = 0.233 g ×
=
0.00459 mol
mol
1
1
01.3 g
0.396 g MCl 2 - 0.163 g Cl – = 0.233 g ×
=
0.00230 mol
mol
1
152.3.3 g
–
0.396 g MCl 3 - 0.163 g Cl = 0.233 g ×
=
0.00153 mol
mol
–1
50.8 g·mol
most likely V
VCl unlikely to be stable
101.3 g·mol–1
most likely Ru RuCl2 stable
152.3 g·mol–1
most likely Eu EuCl3 stable
0.02355 L HCl ×
H2O → H+ + OH–
[H ] =
+
b. i.
9.6 × 10 –14
=
3.1 × 10–7; pH = 6.51
 k  ∆H  1
1
ln 2  =
 – 
R  T1 T2 
 k1 
 9.6 × 10 –14 
∆H  1
1 
ln
–
−15  =
 1.14 × 10  8.314  273 333 
ln 84.2 =
4.43 (8.314)
∆H
= ∆H
(0.003663 – 0.003003) ;
8.314
0.000660
∆H = 5.58 × 10 4 J or 55.8 kJ
ii.
iii.
c.
3.
∆G = – RT ln K = (–8.314)(333) ln(9.6 × 10 –14 ) = 8.30 × 10 4 J
8.30 × 10 4 − 5.58 × 10 4
= –81.7 J
–333
The negative sign corresponds to an increase in order. An increase in order can result
from the H+ and OH– ions structuring the H2O molecules around them.
∆G = ∆H – T∆S ; ∆S =
a.
RCOOH + R’OH → RCOOR’ + H2O
b.
Functional group I – carboxyl group (acid)
Functional group II – hydroxyl group (alcohol)
Functional group III – ester
8.14
0.62
= 1.77
= 1.77 RCOOH is first order.
4.60
0.35
10.6
0.81
= 2.30
= 2.30 R’OH is first order.
4.60
0.35
c.
Trials 1 and 4 can be used to determine the order with respect to HA. [RCOOH] is held
constant; however, both [R’OH] and [HA] vary. The effect of [R’OH] is:
0.50 M
6.57 M
= 1.43 and then (1.43)( 4.60) =
0.35 M
min
9.84
0.75
= 1.50
= 1.50 HA is first order
6.57
0.50
Rate = k[RCOOH] [R’OH] [HA]
4.60 M
= k (0.35M) (0.35M) (0.50M)
min
The effect of [HA] is
k = 75 M–2 ·min–1
d.
Rate determining step could be either:
Step 2 that involves R’OH and product from the reaction of RCOOH and HA, or
Step 3 that involves the breakup of product in step 2.
e. i.
2Cr2O72– + 3CH3CH2OH + 16H+ → 4Cr3+ + 3CH3COOH + 11H2O
The color will change from orange (Cr2O72–(aq)) to green (Cr3+(aq)).
ii.
4.
a.
0.0640 mol ⋅ L–1 × 18.22 mL
= 0.0466 M acid
25.00 mL
b.
With zero % titrated, pH = 3.39; [H+] = 4.07 × 10–4
Assuming negligible dissociation:
[H ][X ] = (4.07 × 10 )
=
+
Ka
Assuming significant dissociation:
–4 2
–
[HX]
= 3.56 × 10
(0.0466)
[H ][X ] = (4.07 × 10 ) = 3.59 × 10
=
[HX]
(0.0466 - 4.07 × 10 )
+
–6
Ka
–4 2
–
–6
–4
With 33.3 % titrated, pH = 5.14; [H+] = 7.24 × 10–6
Assuming negligible dissociation:
[H ][X ] = (7.24 × 10 )(0.0155)
=
+
Ka
Assuming significant dissociation:
[HX]
(0.0311)
[H ][X ] = (7.24 × 10 )(0.0155) = 3.61 × 10
=
[HX]
(0.0311 - 7.24 × 10 )
+
–6
–
= 3.62 × 10
–6
Ka
–
–6
–6
–6
With 66.7 % titrated, pH = 5.74; [H+] = 1.82 × 10–6
Assuming negligible dissociation:
Ka =
Assuming significant dissociation:
[H ][X ] = (1.82 × 10 )(0.0312)
= 3.64 × 10 –6 K a =
X– + H2O = HX + OH–
Kb =
+
[HX]
c.
(0.0155)
[H ][X ] = (1.82 × 10 )(0.0312) = 3.66 × 10
[HX]
(0.0155 -1.82 × 10 )
+
–6
–
)(25)
= 0.0270 M
[X ] = (0.(0466
43.22))
[OH ] ; [OH ] = 8.66 × 10
2.78 × 10 =
–
–6
–6
K w 1.00 × 10 –14
=
= 2.78 × 10 –9
Ka
3.60 × 10 –6
–
– 2
–9
–
0.0270
d.
pH = 6.0 and [H+] = 1.0
3.60 × 10 –6 =
–6
; pOH = 5.06; pH = 8.94
× 10–6
[ ] [ ]
(1.0 × 10 –6 ) X –
; X – = 0.168 M
0.0466
0.168 mol
Moles X– =
× 0.150 L = 0.0252 mol
L
–6
5.
Note: Balanced equations were not required.
a.
P4 + O2 → P4O10
b.
SO2 + H2O → H2SO3 or H+ + HSO3–
c.
Cl2 + Br– → Cl– + Br2
d.
Mg2+ + OH– → Mg(OH)2
e.
AgCl + S2O32– → Ag(S2O3)23– + Cl–
f.
H
C C
H
H
6.
H
→
H H
+ Br2 → Br C C Br
H H
g.
226
88 Ra
4
2 He
a.
For NO2–
b.
c.
NO2– will be bent due to the lone pair of electrons on N. NO2+ will be linear.
Nitrogen-to-oxygen bonds in NO2– will be longer than those in NO2+. The average bond
+
222
86 Rn
O N O
–
O N O
–
and for NO2+
O N O
+
order in NO2– is 11/2. The bond order in NO2+ is 2.
d.
The oxidation number of N in NO2– is +3. Formal charge is zero. The oxidation number is
obtained by assigning bonding electrons to the more electronegative atom. The formal charge is found by
dividing the bonding electrons evenly between atoms. The number of electrons left is compared with the
original number.
7.
a.
b.
A = liquid; B = gas, C and D = solid
1 mL
1 mL
18.0 g ×
= 18.0 mL liquid 18.0 g ×
= 19.6 mL solid
0.9998 g
0.917 g
3
nRT
1 (0.0821) (273.17) (760)
; V=
; V = 3.72 × 10 L
P
4.58
A → B pressure decreases, temperature constant, phase change from liquid to gas
PV = nRT ; V =
c.
B → C pressure constant, temperature decreases, phase change from gas to solid
C → D pressure increases, temperature constant, no phase change
D → A pressure constant, temperature increases, phase change from solid to liquid
8.
a.
The element above element 114 is lead, Pb; above element 116 is polonium, Po; above
element 118 is radon, Rn.
b.
The ionization energy is expected to increase from element 114 to 116 to 118. This is due to
increasing nuclear charge density thus resulting in greater attraction for outer electrons. The ionization
energy of elements 114, 116, and 118 are expected to be lower than those of the elements directly above
them due to the presence of electrons in higher energy levels, those with higher values of n.
c.
The oxidation states for element 114 are predicted to be +2 (loss of electrons in p orbitals) and +4
(loss of electrons in both s and p orbitals).The oxidation state +2 is likely more stable; lower oxidation
states are usually more stable.
d.
Even values of Z are usually more stable due to pairing of protons. Odd values of Z are less stable.
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Lucy Pryde Eubanks, Clemson University, Clemson, SC
Chair
Robert Becker, Kirkwood High School, Kirkwood, MO
Craig W. Bowen, Clemson University, Clemson, SC
Nancy Devino, ScienceMedia Inc., San Diego, CA
Sheldon L. Knoespel, Michigan State University, East Lansing, MI
Steve Lantos, Brookline High School, Brookline, MA
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Robert G. Silberman, SUNY-Cortland, NY
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 16, 2000.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC
2000 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
If anhydrous ammonium nitrate is added to water at room temperature, the temperature of the solution
decreases. However, if anhydrous calcium chloride is added to water at room temperature, the temperature of
the solution increases. Design and carry out an experiment to determine what mass of ammonium nitrate must
be added along with 10.0 g of CaCl2 to 100. mL of water so that the final temperature of the solution is the
same as the initial room temperature of the water. You will be asked to describe the method you developed to
solve this problem.
Lab Problem 2
Although bromocresol green is often used as an acid-base indicator, in this problem it is being used as a
reactant. When mixed with a dilute solution of household bleach, bromocresol green is gradually oxidized and
changes to a different color, a color that happens to match that of bromocresol green at a pH of 4. Design and
carry out an experiment to determine the kinetic order of this redox reaction with respect to bleach. You will
be asked to describe the method you developed to solve this problem.
Page 2
Not valid for use as an USNCO National Examination after April 16, 2000.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan. Include a sketch of the equipment you will use and the
steps you plan to take to solve this problem.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 16, 2000.
Examiner’s Initials:
Page 3
2. Record your data and other observations.
3. What is the mass of ammonium nitrate that must be added along with 10.0 g of calcium chloride to 100 mL
of water and so that the temperature of the resulting solution is the same as that of the water? Show your
methods clearly.
Mass of ammonium nitrate:
Page 4
Not valid for use as an USNCO National Examination after April 16, 2000.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name : ________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the
steps you plan to take to solve this problem.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 16, 2000.
Examiner’s Initials:
Page 5
2. Record your data and other observations.
3. Determine the order of the reaction with respect to bleach. Show your reasoning clearly.
Order of reaction with respect to bleach:
Page 6
Not valid for use as an USNCO National Examination after April 16, 2000.
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Instructions
Directions to the Examiner:
Thank you for administering the 2000 USNCO laboratory practical on behalf of your Local
Section. It is essential that you follow the instructions provided, in order to insure consistency of results
nationwide. There may be considerable temptation to assist the students after they begin the lab exercise.
It is extremely important that you do not lend any assistance or hints whatsoever to the students once
they begin work. As in the international competition, the students are not allowed to speak to anyone
until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab
station or table when the students enter the room. The equipment should be initially placed so that the
materials used for Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry
Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to
reason through a laboratory problem and communicate its results. Do not touch any of the equipment in
front of you until you are instructed to do so.
One of this year’s problems requires the use of a Styrofoam® cup calorimeter. These have been
assembled for you to preserve maximum time for your experimentation. There are extra cups available
for you to use as the inner cup.
Show a Styrofoam® cup calorimeter. (See picture of the set-up on page 3 of these instructions.)
This problem also requires you to use a balance, which is located _____________________________.
Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques
help to minimize the amount of materials you use, thereby increasing safety and minimizing waste.
Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and
reaction plates.
Show a 1-mL, 3-mL, or a 5-mL Beral-type pipet, and show a 6-well or a 12-well reaction plate.
You will be asked to complete two laboratory problems. All the materials and equipment you may want
to use to solve each problem has been set out for you and is grouped by the number of the problem. You
must limit yourself to this equipment for each problem. You will have one hour and thirty minutes to
complete the two problems. You may choose to start with either problem. You are required to have a
procedure for each problem approved for safety by an examiner. (Remember that approval does not
mean that your procedure will be successful–it is a safety approval.) When you are ready for an
examiner to come to your station for each safety approval, please raise your hand.
Page 1
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash
off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate
procedures for disposing of solutions at the end of this lab practical are:
____________________________________________________________________________________
____________________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. Are there any questions before we begin?
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. There is a
periodic table on the last page of the booklet. Allow students enough time to read the brief cover
directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out
all information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for
your cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the
lab practical with the students. They can learn about possible observations and interpretations and you
can acquire feedback as to what they actually did and how they reacted to the problems. After this
discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be
extremely useful to the Olympiad subcommittee as they prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron
sheets from Part I, and the “Blue Books” from Part II to this address:
ACS DivCHED Exams Institute
Clemson University
223 Brackett Hall
Box 340979
Clemson, SC 29634-0979
Monday, April 24, 2000, is the absolute deadline for receipt of the exam materials at the Examinations
Institute. Materials received after this deadline CANNOT be graded.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT
SCHEDULE FOR GRADING THIS EXAMINATION.
Page 2
EXAMINER’S NOTES
Lab Problem #1: Materials and Equipment
Each student will need access to a balance that is capable of weighing an object at least to the nearest
0.01 grams. One balance can serve 2-3 students, but one balance per student is highly desirable.
Each student will need:
One calorimeter such as shown in this
diagram. These should be preassembled for
the students to preserve maximum time for
their experimentation.
1 ring stand
2 clamps
1 thermometer, –10 to 110°C
6 370-mL (12 oz) foam cups
1 370-mL foam cup to serve as cover.
Trim off about top one-third of cup,
punch two holes of appropriate sizes
for the thermometer and stirring rod.
1 5-10 cm piece of wire or string to
hold up thermometer
1 30-35 cm piece of 20 gauge copper
wire. Bend in a circle on the bottom
to serve as a stirrer, hook on top for
handle.
Notes: The inner cup can be replaced as needed, which is why there are extra foam cups.
250 mL (8 oz) foam cups may also be used but 190 mL (6 oz) cups are not recommended. You
may choose to supply a ring clamp to stabilize the base of the cup assembly. Another option is
to place the cup assembly in an appropriately sized beaker that serves to support the cups and to
provide extra insulation.
2 plastic or metal scoops; can use cut Beral-style pipets
1 500-mL or larger wash bottle, labeled “distilled water” or “deionized water”
1 100-mL graduated cylinder
1 small glass or plastic vial with top, labeled 25.0 g ammonium nitrate
1 small glass or plastic vial with top, labeled 25.0 g calcium chloride
6 small plastic weighing boats, weighing papers, or small dry paper cups
4 Beral-style plastic pipets, 3 mL or 5 mL; eye droppers may be substituted.
1 plastic tub for disposal of liquid wastes (or easy access to sinks)
supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
1 pair plastic gloves (optional)
Page 3
Lab Problem #1: Chemicals
Each student will need:
25.0 g anhydrous ammonium nitrate Note: Be sure particles are not clumped.
25.0 g anhydrous calcium chloride Note: Do not use the dihydrate.
500 mL distilled or deionized water
Lab Problem #1: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not
comment on the plan other than looking for any potentially unsafe practices.
2. Consistent results will depend on starting with anhydrous salts and minimizing exposure to
moisture during transfer of the salts to the vials. Be sure the glass or plastic vials are tightly closed after
they are filled for each student.
3. The balances can be used as common equipment, although as noted earlier, it will be highly
desirable to have one balance per student. Only the first activity requires the use of a balance and this
should minimize delays if a number of students must share a balance. Please note any problems on the
Examiner’s report sheet if balance use negatively impacts the performance of your students.
4. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. Please know and follow all
safety procedures appropriate for your site. You will need to give students explicit directions for
handling spills and for disposing of waste materials, following approved safety practices for your
examining site.
Lab Problem #2: Materials and Equipment
Each student will need:
2 6-well or 1 12-well reaction plate. If reaction plates cannot be obtained or borrowed, 6 50-mL or
100-mL beakers can be substituted.
1 piece of white paper to place under the reaction plates or beakers
3 Beral-style plastic pipets, 1 mL. Eye droppers may be substituted.
1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water”
4 10-mL narrow-mouth plastic dropping bottles. Small glass bottles fitted with eye droppers may be
substituted.
One labeled “diluted household bleach solution”
One labeled “bromocresol green solution”
One labeled “pH 4.0 buffer”
One labeled “pH 7.0 buffer”
1 timer, stop watch, or access to classroom clock with a second hand
1 plastic tub for disposal of liquid wastes (or easy access to sinks)
supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
1 pair plastic gloves (optional)
Page 4
Lab Problem #2: Chemicals
Each student will need:
10 mL of pH 4.0 buffer
10 mL of pH 7.0 buffer
10 mL of diluted bleach solution.
Preparation: Obtain fresh commercial bleach containing 5.25% sodium hypochlorite solution.
Combine 5.0 mL of this solution with 295 mL distilled water to form the diluted bleach solution.
10 mL of diluted bromocresol green solution.
Preparation: Dilute 0.2 g of bromocresol green to form 500 mL of aqueous solution. Do not use
existing solutions of bromocresol green, for they may contain alcohol or NaOH commonly used
in their preparation.
50 mL of distilled or deionized water
Lab Problem #2: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not
comment on the plan other than looking for any potentially unsafe practices.
2. Simply combining the two reagents provides enough mixing for the reaction to take place.
However, you may wish to provide toothpicks or plastic stirring rods.
3. Premade buffer solutions of pH 4.0 and 7.0 may be used, or purchased from any one of several
suppliers. If preparing your own buffer solutions, a useful reference is Silberman, R. G., The Journal of
Chemical Education, 1992 (Vol 69, No. 2), p. A42. This article describes a system using boric acid,
citric acid monohydrate, and trisodium phoshate dodecahydrate. These directions are also given in ACS
Small-Scale Laboratory Assessment Activities, ACS Examinations Institute, 1996, p. C-14.
4. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. Please know and follow all
safety procedures appropriate for your site. You will need to give students explicit directions for
handling spills and for disposing of waste materials, following approved safety practices for your
examining site.
Page 5
Page 6
2000 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART III
Lab Problem 1
Plan:
State:
a. Add mass CaCl2 to recorded amount of H2O
b. Record initial and final temperatures to obtain ∆T for CaCl2 sample.
c. Add mass NH4NO3 to recorded amount of H2O.
d. Record initial and final temperatures to obtain ∆T for NH4NO3 sample
e. Repeat 2-3 times.
Sketch the apparatus as directed in the instructions.
Data: Record:
a. mass CaCl2 and volume of water.
b. Ti and Tf for CaCl2
c. mass NH4NO3 and volume of water
d. Ti and Tf for NH4NO3 Amounts of solids added should be large enough to get a
reasonable value for ∆T
e. results of a “verification” experiment to check prediction.
Replication: multiple trials
Calculations:
∆T
∆T
and
for the same amount of H 2 O or
Calculate:
a.
g CaCl 2
g NH 4 NO 3
q
q
and
g CaCl 2
g NH 4 NO 3
for the same amount of H 2 O
b. ratio of grams of NH4NO3 to CaCl2
c. amount of NH4NO3 needed for 100 grams of H2O
Sample Data:
Mass NH4NO3
Volume H2O
Initial temperature
Final temperature
∆T (NH4NO3)
Trial 1
5.005 g
50.0 mL
21.5 °C
14.8 °C
–6.7 °C
Trial 2
5.073 g
51.9 mL
21.5 °C
15.0 °C
–6.5 °C
mass CaCl2
volume H2O
Initial temperature
Final temperature
∆T (CaCl2)
5.064 g
50.0 mL
21.4 °C
33.7 °C
12.3 °C`
5.187 g
52.0 mL
21.0 °C
34.0 °C
13.0 °C
Sample Calculations:
For NH4NO3, note that the solution became cooler as the NH4NO3 dissolved. This means the dissolution
process is absorbing heat from the water. q NH 4 NO3 = – q H 2 O
Trial 1
q NH 4 NO3
Trial 2
mc∆T
=–
g
q NH 4 NO3 = –
mc∆T
g
q NH 4 NO3 = –
q NH 4 NO3 =
(50.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (–6.7 o C)
5.005 g
1.40 × 10 3 J 2.79 × 10 2 J
=
5.005 g
g
qaverage for NH 4 NO 3 =
q NH 4 NO3 = –
q NH 4 NO3 =
(51.9 g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (–6.5 o C)
5.073 g
1.41 × 10 3 J 2.78 × 10 2 J
=
5.073 g
g
2.79 × 10 2 J
or 279 J / g
g
This is the heat absorbed from the water as the NH4NO3 dissolves.
For CaCl2, note that the solution became warmer as the CaCl2 dissolved. This means the dissolution
process is releasing heat to the water. qCaCl2 = – q H 2 O
Trial 1
mc∆T
q=
g
Trial 2
mc∆T
q=
g
qCaCl2 = –
(50.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (12.3 o C)
5.064 g
qCaCl2 = –
(52.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (13.0 o C)
5.187 g
qCaCl2 = –
2.57 × 103 J
5.08 × 10 2 J
=–
5.064 g
g
qCaCl2 = –
2.83 × 10 3 J
5.45 × 10 2 J
=–
5.187 g
g
qaverage for CaCl 2 = –
5.27 × 10 2 J
or – 527 J / g
g
This is the heat released to the water as the CaCl2 dissolves.
1.89 g NH 4 NO 3
527 J / g CaCl 2
=
Ratio of heats: qCaCl2 = – q NH 4 NO3 so the ratio is:
279 J / g NH 4 NO 3
1.00 g CaCl 2
Conclusion:
For 10.0 g CaCl2:
10.0 g CaCl 2 ×
1.89 g NH 4 NO 3
= 18.9 g NH 4 NO 3
1.00 g CaCl 2
Lab Problem 2
Plan:
State:
a. Add a small amount of bromocresol green to a well of pH 4 buffer and to a well of pH 7 buffer to
determine color at each pH.
b. Measure known amounts of bromocresol green to well plates (by counting drops). Add various
amounts of bleach (by counting drops) and record time to color change.
c. Ideally, the reactions will be conducted at constant volume by adding either H2O or pH 7 buffer so
that concentration ratios are directly related to volumes of bleach used for each reaction.
Replications should be made for each trial.
Data: Record:
a. color change of bromocresol green from blue at pH = 7 to yellow at pH = 4
b. drops of bromocresol green added to each well
c. drops of buffer or water added to each well, if used in plan
d. drops of bleach added to each well
e. time required for color change
Replication: multiple trials
Calculations:
To determine the order of reaction, use one of these equations.
time1 [bleach2 ]
=
time2 [bleach1 ]x
x
rate1 [bleach1 ]
=
rate2 [bleach2 ]x
x
or
Sample Data:
The observed color change for bromocresol green is from blue at pH = 7 to yellow at pH = 4.
Data gathered if reactions are buffered at pH = 7
Trial
Drops of
Drops of
Bromocresol Green
Buffer
1
20
5
2
20
5
3
20
5
4
20
1.5
5
20
1.5
6
20
1.5
Drops of
Bleach
20
20
20
10
10
10
Data gathered if reactions are not buffered
Trial
Drops of
Drops of
Drops of
Bromocresol Green
Buffer
Bleach
1
20
5
20
2
20
5
20
3
20
5
20
4
20
1.5
10
5
20
1.5
10
6
20
1.5
10
Note: Students do not need to do both buffered and unbuffered reactions.
Sample Calculation:
Using data from the observations from the pH 7 buffered reactions:
Time for color
change to occur
14
16
17
28
27
28
Average time
for three trials
Time for color
change to occur
50
52
53
50
51
54
Average time
for three trials
15.7
27.7
51.7
51.7
time1 [bleach2 ]
=
and the concentration of bleach is directly proportional to the number of drops added.
time2 [bleach1 ]x
x
27.7 seconds (20 drops)
=
and x ≈ 1
15.7 seconds (10 drops) x
x
rate = k[bleach]1 if the reaction is carried out in the buffer.
Using data from the observations from the unbuffered reactions:
time1 [bleach2 ]
=
and the concentration of bleach is directly proportional to the number of drops added.
time2 [bleach1 ]x
x
51.7 seconds (20 drops)
=
and x ≈ 0
51.7 seconds (10 drops) x
x
rate = k[bleach]0 if the reaction is not carried out in the pH = 7 buffer.
Conclusion:
If buffered at pH = 7, then the reaction is first order with respect to bleach.
If unbuffered, then the reaction is zero order with respect to bleach.
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland, NY
Chair
Jo A. Beran, Texas A&M University-Kingsville, TX
Peter E. Demmin (retired), Amherst Central High School, NY
Edward DeVillafranca (retired), Kent School, CT
Dianne H. Earle, Paul M. Dorman High School, SC
Alice Johnsen, Bellaire High School, TX
Patricia A. Metz, United States Naval Academy, MD
Ronald O. Ragsdale, University of Utah, UT
Diane D. Wolff, Western Virginia Community College, VA
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 22, 2001, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 22, 2001.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
N A gas constant
R Planck’s constant
h
°C gram
g pressure
P
c heat capacity
C p rate constant
k
C hour
h retention factor
Rf
E joule
J second
s
Ea kelvin
K speed of light
c
H kilo- prefix
k temperature, K
T
S liter
L time
t
volt
V
E =E –
USEFUL EQUATIONS
 –∆H   1 
ln K = 
   +c
 R T
RT
lnQ
nF
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol –1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
 k 2  Ea  1 1 
=
−
 k1  R  T1 T2 
ln 
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
116
118
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(289)
(289)
(293)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as a USNCO National Exam after April 22, 2001.
DIRECTIONS
§ When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
§ You may write on the test booklet, but it will not be used for grading.
§ There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
§ Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which of these compounds is amphoteric?
I. Al(OH)3
II. Ba(OH)2
III. Zn(OH)2
7. What is the purpose of this
apparatus?
water out
(A) I only
(B) II only
(C) I and III only
(D) II and III only
2. Calcium hydride reacts with excess water to form
(A) CaO and H2
(B) Ca(OH)2 and O 2
(C) Ca(OH)2 only
(D) Ca(OH)2 and H 2
3. What is the most likely
boiling point of an
equimolar mixture of
hexane, C6H14, and
heptane, C7H16?
Boiling Point
C 6H14
69 °C
C 7H16
98 °C
(A) below 69 °C
(B) between 69 and 98 °C
(C) 69 °C
(D) 98 °C
4. Which element melts at the highest temperature?
(A) aluminum
(B) silicon
(C) phosphorus
(D) sulfur
5. Which substance participates readily in both acid-base
and oxidation-reduction reactions?
(A) Na 2CO3
(B) KOH
(C) KMnO4
(D) H2C 2O4
6. What mass of magnesium
hydroxide is required to
neutralize 125 mL of
0.136 M hydrochloric
acid solution?
Substance Molar Mass
Mg(OH)2
(A) 0.248 g
(B) 0.496 g
(C) 0.992 g
(D) 1.98 g
58.33 g·mol–1
water in
(A) distilling
(B) filtering
(C) refluxing
(D) titrating
8. Calculate the mass of
Substance
Molar Mass
ammonia that can be
produced from the
(NH4)2PtCl6
443.9 g·mol–1
decomposition of a
sample of (NH4)2PtCl6 containing 0.100 g Pt.
(A) 0.0811 g
(B) 0.0766 g
(C) 0.0175 g
(D) 0.00766 g
9. Assume 0.10 L of N2 and 0.18 L of H2, both at 50 atm
and 450 °C, are reacted to form NH3. Assuming the
reaction goes to completion, identify the reagent that is
in excess and determine the volume that remains at the
same temperature and pressure.
(A) H2, 0.02 L
(B) H2, 0.08 L
(C) N2, 0.01 L
(D) N2, 0.04 L
10. Concentrated hydrochloric acid is 12.0 M and is 36.0%
hydrogen chloride by mass. What is its density?
(A) 1.22 g·mL–1
(B) 1.10 g·mL–1
(C) 1.01 g·mL–1
(D) 0.820 g·mL–1
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 3
11.
C4H6O3 →
acetic
anhydride
C7H6O3 +
salicylic
acid
C 9H8O4 +
aspirin
C2H4O2
acetic
acid
What is the percent yield if
0.85 g of aspirin is formed
in the reaction of 1.00 g of
salicylic acid with excess
acetic anhydride?
Substance Molar Mass
C 7H6O3
C 4H6O3
C 9H8O4
C 2H4O2
138.12 g·mol–1
102.09 g·mol–1
180.15 g·mol–1
60.05 g·mol –1
(A) 65 %
(C) 85 %
(D) 91 %
(B) 77 %
12. The triple point of CO2 occurs at 5.1 atm and –56 °C. Its
critical temperature is 31 °C. Solid CO2 is more dense
than liquid CO2. Under which combination of pressure
and temperature is liquid CO2 stable at equilibrium?
(A) 10 atm and –25 °C
(B) 5.1 atm and –25 °C
(C) 10 atm and 33 °C
(D) 5.1 atm and –100 °C
13. The vapor pressure of
water at 20 °C is
17.54 mmHg. What will be
the vapor pressure of the
water in the apparatus
shown after the piston is
lowered, decreasing the
volume of the gas above
the liquid to one half of its
initial volume? (Assume
no temperature change.)
16. What is the average velocity of H2 molecules at 100 K
relative to their velocity at 50 K?
(A) 2.00 times the velocity at 50 K
(B) 1.41 times the velocity at 50 K
(C) 0.71 times the velocity at 50 K
(D) 0.50 times the velocity at 50 K
17. What type of semiconductor results when highly purified
silicon is doped with arsenic?
(A) n–type
(B) p–type
(C) q–type
(D) s–type
18. The heat of formation of
NO from its elements is
+90 kJ·mol–1. What is the
approximate bond
dissociation energy of the
bond in NO?
Bond
N N
O O
Bond Energy
941 kJ·mol –1
499 kJ·mol –1
(A) 630 kJ·mol –1
(B) 720 kJ·mol –1
(C) 765 kJ·mol –1
(D) 810 kJ·mol –1
water vapor
liquid water
(A) 8.77 mmHg
19. How much energy must be
supplied to change 36 g of
ice at 0 °C to water at room
temperature, 25 °C?
Data for Water, H2O
∆Hofusion 6.01 kJ·mol–1
C P, liquid 4.18 J·K –1·g –1
(A) 12 kJ
(B) 16 kJ
(C) 19 kJ
(D) 22 kJ
(B) 17.54 mmHg
(C) 35.08 mmHg
20. For a process that is both endothermic and spontaneous,
(D) between 8.77 and 17.54 mmHg
14. What is the density of propane, C3H8, at 25 °C and
740. mmHg?
(A) 0.509 g·L –1
(B) 0.570 g·L –1
(C) 1.75 g·L –1
(D) 1.96 g·L –1
15. An unknown gas effuses through a small hole one half as
fast as methane, CH 4, under the same conditions. What is
the molar mass of the unknown gas?
(A) ∆H < 0
(B) ∆G > 0
(C) ∆E = 0
(D) ∆S > 0
21. Consider the values for ∆Ho (in kJ·mol–1) and for ∆So
(in J·mol –1·K–1) given for four different reactions. For
which reaction will ∆Go increase the most (becoming
more positive) when the temperature is increased from
0 °C to 25 °C?
(A) ∆Ho = 50, ∆So = 50
(B) ∆Ho = 90, ∆So = 20
(A) 4 g·mol –1
(B) 8 g·mol –1
(C) ∆Ho = –20, ∆So = –50
(C) 32 g·mol –1
(D) 64 g·mol –1
(D) ∆Ho = –90, ∆So = –20
Page 4
Not valid for use as an USNCO National Examination after April 22, 2001.
22. A certain reaction is exothermic by 220 kJ and does
10 kJ of work. What is the change in the internal energy
of the system at constant temperature?
(A) +230 kJ
(B) +210 kJ
(C) –210 kJ
(D) –230 kJ
23. Fe2O3(s) + 3/2C(s) → 3/2CO2(g) + 2Fe(s) ∆Ho = +234.1 kJ
C(s) + O2(g) → CO2(g)
∆Ho = –393.5 kJ
Use these equations and ∆Ho values to calculate ∆Ho for
this reaction.
4Fe(s) + 3O2(g) → 2Fe2O3(s)
(A) –1648.7 kJ
(B) –1255.3 kJ
(C) –1021.2 kJ
(D) –129.4 kJ
24. A large positive value of ∆G corresponds to which of
these?
o
27. Consider this reaction.
2NO2(g) + O3(g) → N 2O5(g) + O2(g)
The reaction of nitrogen dioxide and ozone represented is
first order in NO2(g) and in O3(g). Which of these possible
reaction mechanisms is consistent with the rate law?
Mechanism I.
NO2 + O3 → NO3 + O2 slow
NO3 + NO2 → N2O5
fast
Mechanism II. O3 ¾ O2 + O
fast
NO2 + O → NO3
NO3 + NO2 → N2O5
slow
fast
(A) I only
(B) II only
(C) both I and II
(D) neither I nor II
28. When the temperature of a reaction is raised from 300 K
to 310 K, the reaction rate doubles. Determine the
activation energy, Ea , associated with the reaction.
(A) small positive K
(B) small negative K
(A) 6.45 kJ·mol–1
(B) 23.3 kJ·mol–1
(C) large positive K
(D) large negative K
(C) 53.6 kJ·mol–1
(D) 178 kJ·mol –1
25. What names apply to
chemical species
corresponding to locations
1 and 2 on this reaction
coordinate diagram?
2
1
29. Use the experimental data in the table to determine the rate
law for this reaction.
A +B → AB
These data were obtained when the reaction was studied.
[A], M
[B], M
∆[ AB] ∆t mol·L–1·s–1
0.100
0.200
0.300
Reaction Progress →
Location 1
Location 2
0.100
0.100
0.300
2.0 × 10–4
2.0 × 10–4
1.8 × 10–3
What is the rate equation for the reaction?
(A) activated complex
activated complex
(A) rate = k [A] [B]
(B) rate = k [A]2
(B) reaction intermediate
activated complex
(C) rate = k [B]
(D) rate = k [B]2
(C) activated complex
reaction intermediate
(D) reaction intermediate
reaction intermediate
30. Which of the reactions represented in these diagrams will
show the greatest increase in rate for a given increase in
temperature?
26. Gadolinium-153, which is used to detect osteoporosis,
has a half-life of 242 days. Which value is closest to the
percentage of the Gd-153 left in a patient's system after
2 years (730 days)?
(A) 33.0 %
(B) 25.0 %
(C) 12.5 %
(D) 6.25 %
Reaction I
Reaction II
(A) Reaction I forward
(B) Reaction I reversed
(C) Reaction II forward
(D) Reaction II reversed
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 5
Questions 31 and 32 should both be answered with reference
to this equilibrium system.
2NH3(g) ¾ N2(g) + 3H2(g)
K p = 80.0 at 250 °C
38.
31. What is Kp for this reaction?
1/2N2(g) + 3/2H2(g) ¾ NH3(g)
Eo = –1.66 V
Eo = +0.34 V
(A) 0.0125
(B) 0.112
What voltage is produced under standard conditions by
combining the half-reactions with these Standard
Electrode Potentials?
(C) 8.94
(D) 40.0
(A) 1.32 V
32. What is the expression for Kc at 250 °C for this reaction?
2NH3(g) ¾ N2(g) + 3H2(g)
(A) Kc =
(C) Kc
Kp
(RT )2
2
= K p (RT)
Kp
RT
(B) Kc
=
(D) Kc
= K p RT
HCOOH(aq) ¾ H+(aq) + HCOO–(aq) K a = 1.7 × 10 –4
The ionization of formic acid is represented.
Calculate [H+] of a solution initially containing
0.10 M HCOOH and 0.050 M HCOONa.
(B) 2.00 V
(C) 2.30 V
(D) 4.34 V
39. For which of these oxidation/reduction pairs will the
reduction potential vary with pH?
I. AmO22+/AmO2+ II. AmO22+/Am4+ III. Am4+/Am2+
40.
33.
Al3+(aq) + 3e– → Al(s)
Cu 2+(aq) +2e– → Cu(s)
(A) I only
(B) II only
(C) I and II only
(D) I, II, and III
2Ag+(aq) + Cu(s) → Cu2+(aq) + 2Ag(s)
The standard potential for this reaction is 0.46 V. Which
change will increase the potential the most?
(A) doubling the [Ag+]
(B) halving the [Cu2+]
(A) 8.5 × 10 M
(B) 3.4 × 10 M
(C) doubling the size of the Cu(s) electrode
(C) 4.1 × 10 M
(D) 1.8 × 10–2 M
(D) decreasing the size of the Ag electrode by one half
–5
–3
–4
34. Which are strong acids?
I. HClO3
II. H2SeO3
III. H3AsO4
(A) I only
(B) III only
(C) I and III only
(D) II and III only
35. Carbonic acid, H2CO3, is a diprotic acid for which
K 1 = 4.2 × 10 –7 and K2 = 4.7 × 10 –11. Which solution will
produce a pH closest to 9?
(A) 0.1 M H 2CO3
(B) 0.1 M Na2CO3
(C) 0.1 M NaHCO3
(D) 0.1 M NaHCO3 and
0.1 M Na2CO3
36. What is the pH of a saturated solution of magnesium
hydroxide, Mg(OH) 2 at 25 °C? (K sp = 6.0 × 10 –12 at 25˚C)
(A) 10.56
(B) 10.36
(C) 10.26
(D) 10.05
41. 10Cl–(aq) + 2MnO4–(aq) + 16H+(aq) →
5Cl2(g) + 2Mn2+(aq) + 8H2O(l)
The value of Eo for this reaction at 25 °C is 0.15 V. What
is the value of K for this reaction?
(A) 2.4 × 1025
(B) 4.9 × 1012
(C) 1.2 × 105
(D) 3.4 × 102
42. When water is electrolyzed, hydrogen and oxygen gas
are produced. If 1.008 g of H 2 is liberated at the cathode,
what mass of O2 is formed at the anode?
(A) 32.0 g
(B) 16.0 g
(C) 8.00 g
(D) 4.00 g
43. Which property of an element is most dependent on the
shielding effect?
(A) atomic number
(B) atomic mass
37.
P 4(s) + 3OH–(aq) + 3H2O(l) → PH 3(g) + 3H2PO2–(aq)
For this reaction, the oxidizing and reducing agents are,
respectively,
(A) P 4 and OH–
(B) OH– and P4
(C) P 4 and H 2O
(D) P 4 and P4
Page 6
(C) atomic radius
(D) number of stable isotopes
Not valid for use as an USNCO National Examination after April 22, 2001.
44. How many unpaired electrons are present in a ground
state gaseous Ni2+ ion?
(A) 0
(B) 2
(C) 4
(D) 6
45. When the elements carbon, nitrogen and oxygen are
arranged in order of increasing ionization energies, what
is the correct order?
(A) C, N, O
(B) O, N, C
(C) N, C, O
(D) C, O, N
(A) n = 2, l = 0
(B) n = 2, l = 1
(C) n = 3, l = 0
(D) n = 3, l = 1
47. Which element will exhibit the photoelectric effect with
light of the longest wavelength?
(B) Rb
(C) Mg
(D) Ca
48. All these elements have common allotropes except
(A) C
(B) O
(C) Kr
(A) Electrons are simultaneously attracted by more than
one nucleus.
(B) Filled orbitals of two or more atoms overlap one
another.
(C) Unoccupied orbitals of two or more atoms overlap
one another.
(D) Oppositely-charged ions attract one another.
46. Given this set of quantum numbers for a multi-electron
atom: 2, 0, 0, 1 /2 and 2, 0, 0, –1 /2. What is the next
higher allowed set of n and l quantum numbers for this
atom in its ground state?
(A) K
52. Which is the best description of a covalent bond?
(D) S
53. What is the formal charge on the chlorine atom in the
oxyacid HOClO2 if it contains only single bonds?
(A) –2
(B) –1
(C) +1
(D) +2
54. Which of these compounds is not adequately represented
by a valence bond model?
I. CO2
II. SO2
III. SiO2
(A) I only
(B) II only
(C) I and III only
(D) II and III only
55. Which compound is not correctly matched with its class
name?
(A) HCOOH, acid
(B) C 6H5CHO, aldehyde
(C) C 2H5COCH3, ether
49. How many sigma and pi
bonds are shown in this
compound?
H H O
H N C C C C C H
H
56. How many toluene derivatives have the formula C7H7Cl?
(A) 8 sigma and 7 pi
(B) 8 sigma and 3 pi
(C) 11 sigma and 3 pi
(D) 11 sigma and 4 pi
50. Which reaction involves a change in the electron-pair
geometry for the underlined atom?
(A) B F 3 + F– → B F 4
(B) NH3 + H+ → NH4+
(C) 2S O2 + O2 → 2 S O3
(D) H2O + H+ → H3O +
51. Which compound has the largest lattice energy?
(A) NaF
(B) CsI
(C) MgO
(D) CH3CHOHCH3, secondary alcohol
(D) CaS
(A) 1
(B) 2
(C) 3
(D) 4
57. When the compounds CH 3COOH, C2H5OH and C 6H5OH
are arranged in order of increasing acidity in aqueous
solution, which order is correct?
(A) C 2H5OH < CH3COOH < C 6H5OH
(B) C 6H5OH < CH3COOH < C 2H5OH
(C) CH3COOH < C 6H5OH < C2H5OH
(D) C 2H5OH < C6H5OH < CH3COOH
58. Which can be used as a catalyst in an esterification
reaction?
I. NaOH
II. H2SO4
(A) I only
(B) II only
(C) both I and II
(D) neither I nor II
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 7
59. Which term best describes a carbocation?
60. A racemic mixture has equal amounts of
(A) electrophilic
(B) free radical
(A) alkanes and alkenes.
(C) hydrophobic
(D) nucleophilic
(B) cis and trans isomers.
(C) functional group isomers.
(D) enantiomers.
END OF TEST
Page 8
Not valid for use as an USNCO National Examination after April 22, 2001.
U. S. National Chemistry Olympiad – 2001
National Examination—Part I
SCORING KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Answer
C
D
B
B
D
B
C
C
D
A
A
A
B
C
D
B
A
A
B
D
Number
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
Answer
C
D
A
A
B
C
C
C
D
B
B
A
B
A
C
B
D
B
B
A
Property of the ACS Society Committee on Education
Number
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
A
C
C
B
D
B
B
C
D
A
C
A
D
B
C
D
D
B
A
D
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland, NY
Chair
Jo A. Beran, Texas A&M University-Kingsville, TX
Peter E. Demmin (retired), Amherst Central High School, NY
Edward DeVillafranca (retired), Kent School, CT
Dianne H. Earle, Paul M. Dorman High School, SC
Alice Johnsen, Bellaire High School, TX
Patricia A. Metz, United States Naval Academy, MD
Ronald O. Ragsdale, University of Utah, UT
Diane D. Wolff, Western Virginia Community College, VA
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a
score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 22, 2001, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front
of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
Not valid for use as an USNCO National Exam after April 22, 2001.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper,
and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
1.
(10%) The mass percent of MnO2 in a sample of a mineral is determined by reacting it with a measured excess of As2O3 in acid
solution, and then titrating the remaining As2O3 with standard KMnO4. A 0.225 g sample of the mineral is ground and boiled
with 75.0 mL of 0.0125 M As2O3 solution containing 10 mL of concentrated sulfuric acid. After the reaction is complete, the
solution is cooled, diluted with water, and titrated with 2.28 × 10 –3 M KMnO4, requiring 16.34 mL to reach the endpoint.
Note: 5 mol of As2O3 react with 4 mol of MnO4–.
a. Write a balanced equation for the reaction of As2O3 with MnO2 in acid solution. The products are Mn2+ and AsO43–.
b. Calculate the number of moles of
i. As2O3 added initially.
ii. MnO4– used to titrate the excess As2O3.
iii. MnO2 in the sample.
c. Determine the mass percent of MnO2 in the sample.
d. Describe how the endpoint is detected in the KMnO4 titration.
2.
(15%) The presence of CO32–, HCO3– and CO2 in body fluids helps to stabilize the pH of these fluids despite the addition or
removal of H+ ions by body processes. Answer the following questions about solutions containing these species in varying
combinations at 25 °C. K1 and K2 for H2CO3 are 4.2 × 10 –7 and 4.7 × 10 –11, respectively.
a. Write balanced equations to represent the processes responsible for K1 and K2.
b. Calculate the [H+] and pH expected for
i. 0.033 M solution of H2CO3, which is the saturation point of CO 2 at 25 °C.
ii. 1:1 mixture of H2CO3 and HCO 3–.
iii. 1:1 mixture of HCO3– and CO32–.
iv. 0.125 M solution of CO32–.
c. The “normal” pH in blood plasma is 7.40. Identify the components that would provide the best buffer at this pH
and calculate their ratio.
d. The value of K 1 is based on the assumption that all of the CO 2 dissolved in water exists in the form of H2CO3.
However, recent evidence suggests that an additional equilibrium exists as represented by this equation.
CO2(aq) + H2O(l) ¾ H2CO3(aq)
When the “true” concentration of H2CO3(aq) is taken into account, K 1 =2 × 10 –4. Use this information to determine the percent
of dissolved CO2 that is actually present as H2CO3(aq) .
3.
(12%) Glucose, C 6H12O6, is readily metabolized in the body.
a. Write a balanced equation for the metabolism of C6H12O6 to CO2 and H 2O.
b. Calculate ∆Gometabolism for glucose. Given: The free energy of formation, ∆Gf , is –917 kJ·mol –1 for C6H12O6(s);
–394.4 kJ·mol–1 for CO 2(g); –237.2 kJ·mol–1 for H2O(l).
c. If ∆Ho for this process is –2801.3 kJ, calculate ∆So at 25 °C.
d. One step in the utilization of energy in cells is the synthesis of ATP4– from ADP3– and H 2PO4–, according to
this equation.
ADP3– + H2PO4– → ATP4– ∆Go = 30.5 kJ·mol –1
i. Calculate the number of moles of ATP4– formed by the metabolism of 1.0 g of glucose.
ii. Calculate the equilibrium constant, K, for the formation of ATP 4– at 25 ˚C.
o
4.
(11%) The corrosion of iron is an electrochemical process that involves the standard reduction potentials given here at 25 °C.
Fe2+(aq) + 2e– → Fe(s)
Eo = –0.44 V
+
–
O2(g) + 4H (aq) + 4e → 2H2O(l)
Eo = +1.23 V
a. Calculate the voltage for the standard cell based on the corrosion reaction. 2Fe (s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)
b. Calculate the voltage if the reaction in Part a occurs at pH = 4.00 but all other concentrations are maintained as they were in
the standard cell.
c. For the reaction Fe(OH)2(s) + 2e– → Fe(s) + 2OH–(aq) , Eo = –0.88 V. Use this information with one of the given standard
potentials to calculate the Ksp of Fe(OH)2.
Page 2
Not valid for use as an USNCO National Exam after April 22, 2001
d. An iron object may be protected from corrosion by coating it with tin. This method works well as long as the tin coating is
intact. However, when the coating is penetrated, the corrosion of the iron is actually accelerated. Use electrochemical
principles to account for both of these observations. The standard reduction potential for tin is:
Sn 2+(aq) + 2 e – → Sn (s)
Eo = –0.14 V
5.
6.
(14%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the reactants and
products. Omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic
substances. You need not balance the reactions. All reactions occur in aqueous solution unless otherwise indicated.
a. Solid calcium hydrogen carbonate is heated to a very high temperature.
b. Solid potassium sulfite is added to a solution of hydrochloric acid.
c. Solutions of barium hydroxide and sulfuric acid are mixed.
d. A tin (II) chloride solution is added to an acidic solution of potassium dichromate.
e. Concentrated hydrochloric acid is added to a solution of sodium hypochlorite.
f. Nitrogen-16 undergoes β– decay.
(13%) Answer these questions pertaining to chemical kinetics.
a. Determine the reaction rate at 10 seconds from the graph. Show
your work.
b. Using the same units for the reaction rate as in Part a, and assuming
concentrations in mol·L–1, give the units for the rate constant of a
reaction with an order of:
i. zero
ii. one
iii. two
c. Consider this reaction: 4HBr(g) + O2(g) → 2H 2O(l) + 2Br 2(g)
i. Express the reaction rates for HBr and Br 2 in this reaction relative to
that of O2.
ii. Explain why this reaction is unlikely to occur by direct collision of four HBr molecules with one O 2 molecule.
d. This mechanism has been suggested for the reaction in Part c:
HBr(g) + O2(g) → HOOBr (g)
Step 1
HOOBr(g) + HBr(g) → 2HOBr (g)
Step 2
HOBr(g) + HBr(g) → H 2O(g) + Br 2(g)
Step 3
Give the rate equation in terms of reactants expected for this reaction if the rate-determining step is:
i. Step 1
ii. Step 2
iii. Step 3
Assume in each case that the steps before the rate-determining step are in rapid equilibrium. Outline your reasoning in each
case.
7.
(13%) A certain element, X, forms the fluorides XF 3 and XF 5. Element X also reacts with sodium to form Na3X.
a. Give the symbol of an element that behaves in this way.
b. For both XF 3 and XF 5;
i. write Lewis electron dot structures.
ii. describe the electron pair and molecular geometries.
iii. give the hybridization of the X atom.
c. The bonds in XF5 are not all the same length. Identify the longer bonds and account for this behavior.
d. Another element, Y, in the same family as X, forms YF 3 but not YF5. Identify element Y and account for its inability
to form YF5.
8.
(12%) Account for each observation with appropriate atomic or molecular properties.
a. Carbon dioxide has a higher vapor pressure than sulfur dioxide at the same temperature.
b. Hydrogen chloride has a lower normal boiling point than either hydrogen fluoride or hydrogen bromide.
c. Calcium oxide has a much higher melting point (2580 °C) than potassium fluoride (858 °C).
d. Tin (II) chloride is an ionic compound (mp = 240 °C) while tin(IV) chloride is a covalent compound (bp = 114 °C).
END OF PART II
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
N A gas constant
R Planck’s constant
h
°C gram
g pressure
P
c heat capacity
C p rate constant
k
C hour
h retention factor
Rf
E joule
J second
s
Ea kelvin
K speed of light
c
H kilo- prefix
k temperature, K
T
S liter
L time
t
volt
V
E =E –
USEFUL EQUATIONS
 –∆H   1 
ln K = 
   +c
 R T
RT
lnQ
nF
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol –1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
 k 2  Ea
=
 k1  R
ln 
1 1
−
 T1 T2 
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
112
116
118
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(277)
(289)
(293)
Page 4
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 16, 2000.
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART II
1.
a.
2MnO2 + As2O3 + H2O → 2Mn2+ + 2AsO43– + 2H+
b. i.
0.0750 L ×
ii.
2.28× 10 –3 mol
= 3.73 ×10 −5 mol MnO 4–
L
5 mol As2O 3
mol MnO4 – ×
= 4.66× 10−5 mol As 2O 3 left
4 mol MnO 4–
0.01634 L ×
3.73× 10 −5
iii.
0.0125 mol
= 9.38× 10 −4 mol As 2O3
L
9.38 × 10–4 – 4.66 × 10–5 = 8.91 × 10–4 mol As2O3 react with MnO2
8.91× 10−4 mol As2 O3 ×
1.78 ×10 –3 mol MnO 2 ×
c.
86.94 g MnO2
= 0.155 g MnO2
mol MnO 2
0.155 g MnO2
× 100 = 68.9% MnO 2 in sample
0.225 g sample
mass % MnO 2 =
2.
2 mol MnO 2
= 1.78 ×10 −3 mol MnO 2
1 mol As2 O3
d.
The endpoint corresponds to a slight purple (pink) color due to excess MnO4–(aq).
a.
Process responsible for K1
H2CO3(aq) ¾ H+(aq) + HCO3–(aq)
Process responsible for K2
HCO3–(aq) ¾ H+(aq) + CO32–(aq)
[ H ] [ HCO ]
=
+
b. i.
ii.
[H 2CO3 ]
[H 2CO3 ] = [HCO 3– ]
K1
[H ]
=
+ 2
–
3
4.2× 10
–7
(0.033)
+
iii.
[HCO 3– ] = [CO32– ]
iv.
CO32–(aq) + H2O(l) ¾ HCO3–(aq) + OH–(aq)
[OH ] =
[
Kw
× CO 32
Ka
–
+
[ ] [ H ] = 4.2 ×10
= [H ] [ H ] = 4.7 ×10
K1 = H +
K2
[ H ] = 1.2 × 10
+
+
–4
pH = 3.93
–7
pH = 6.38
–11
pH =10.33
]
–
10
× 0.125
[OH ] = 1.0×
4.7 ×10
pOH = 2.28
[OH ] = 2.7× 10
[ OH ] = 0.0052
[H ] = 1.9× 10 pH = 11.72
Components of the best buffer: H CO
pH = 7.40 [ H ] = 4.0 × 10
–14
–
–11
–
–5
+
c.
–
–12
+
–8
2
4.2 × 10
–7
= 4.0 × 10
–8
×
HCO 3–
H 2CO 3
Key for 2001 USNCO National Exam, Part II
3
and HCO3–
HCO 3–
= 10.5
H 2CO3
Page 1
d.
3.
4.2 × 10 –7
= 2 ×10 –3 or 0.002
2 × 10 –4
Because the ratio of the two K values is 0.002, 0.2% of dissolved CO2 is actually H2CO3.
a.
C6H12O6 + 6O2 → 6CO2 + 6H2O
b.
o
∆Gometabolism = 6∆GCO
+ 6∆G oH2 O – ∆GCo 6 H12 O6
2
= 6 mol(−394.4 kJ⋅mol –1 )+ 6 mol(–237.2 kJ⋅ mol –1) – ( – 9 1 7kJ⋅ mol –1)
= –2366.4 k J – 1 4 2 3 . 2+k 9J 1 7 k J
= –2873kJ
∆Go = ∆Ho – T∆So
c.
( 2 9 8 K∆S) o = 72 kJ
– 2 8 7 3 k J= –2801.3 k J – 2 9 K∆S
8 o
d. i.
1.0 g C 6H12 O 6 ×
ii.
4.
a.
2872.6 kJ
= 16 kJ
mol
1 mol ATP
= 0.52 mol ATP formed
30.5 kJ
30.5 ×10 3J –8.314 J
=
(298 K)ln K
mol
mol ⋅ K
∆Go = –RT ln K
ln K = –12.31
2Fe(s) → 2Fe2+(aq) + 4e–
Eo = +0.44 V
O2(g) + 4H+(aq) + 4e– → 2H2O(l)
Eo = +1.23 V
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq)
b.
or 240 J/K
1 mol
= 5.6× 10 –3 m o l C6 H 12O6
180 g
5.6 ×10 –3 mol C 6H12 O6 ×
16 kJ ×
∆So = 0.24 kJ / K
[
K = 4.5 ×10 –6
Eo = +1.67 V
+ 2H2O(l)
]
2+ 2
Fe
RT
E =E –
ln
4
nF
H+ PO
2
o
[ ]( )
= 1.67V –
(8 . 3 1 4 J/ mol⋅ K) ( 2 9 8 K)
( 96,500 J / V) ( 4 mol)
ln
1
(1.0 ×10 –4 ) (1)
4
= 1.67 – (0.00642) (+36.84)
= 1.43 V
c.
Fe(OH)2(s) + 2e– → Fe(s) + 2OH–(aq)
Fe(s) → Fe2+(aq) + 2e–
Eo = –0.88 V
Eo = +0.44 V
Fe(OH)2(s) ¾ Fe2+(aq) + 2OH–(aq)
Eo = –0.44 V
∆Go = –nFEo = – RT ln Ksp
ln K sp =
Page 2
nFE o
RT
ln K sp =
(2 mol) (96,500J
⋅ V –1 ) (–0.44 V)
= –34.28
–1
( 8 . 3 1 4 J⋅ mol ⋅ K –1) (298 K)
Ksp = 1.30 × 10 –15
Key for 2001 USNCO National Exam, Part II
d. When iron is coated with Sn, the reaction Sn → Sn2+ + 2e– takes place. If the tin coating is broken, the
reaction Sn2+ + Fe → Sn + Fe2+ becomes spontaneous. Iron becomes the anode and is oxidized more
readily.
5.
6.
Note: Balanced equations were not required.
a.
Ca(HCO3)2 → CaO + H2O + 2CO2 (partial credit for CaCO3)
b.
K2SO3 + H+ → K+ + H2O + SO2 (partial credit for H2SO3)
c.
Ba2+ + OH– + H+ + SO42– → H2O + BaSO4
d.
Cr2O72– + Sn2+ + H+ → Cr3+ + Sn4+ + H2O (partial credit for SnCl4)
e.
H+ + Cl– + OCl–→ Cl2 + H2O (partial credit for HClO)
f.
16
7N
a.
Tangent to curve at 10 seconds:
b. i.
ii.
iii.
rate = k
rate = k [ ]
rate = k [ ]2
c. i.
Expressed in symbols:
ii.
d. i.
ii.
iii.
7.
a.
→
0
–1
+
16
8O
∆M –0.39M
=
= –0.012 M⋅ s–1
∆T
32 s
units are M.s–1
units are s–1
units are M–1.s–1
–d[ O 2 ] –d[ HBr ] d[ Br2 ]
=
=
dt
4dt
2dt
This shows that the rate of disappearance of HBr is 4 times that of O2 and the rate of production of
Br2 is twice the rate of disappearance of O2.
More than mono- or bi-molecular steps improbable.
rate = k [HBr] [O2]
Rate is proportional to reactants in the rate-limiting step.
rate = k [HBr]2 [O2]
[HOOBr] in the rate equation must be stated in terms of the
previous equilibrium.
rate = k [HBr]2 [O2]1/2
[HOBr] in the rate equation must be stated in terms of the
previous equilibria.
Phosphorus (P) and arsenic (As) might behave in this manner.
b. i.
F X F
F
ii.
iii.
c.
d.
F
F
F
X F
F
XF3
Electron pair geometry is tetrahedral; atom geometry is trigonal pyramidal.
XF5
Electron pair geometry and atom geometry are both trigonal bipyramidal.
XF3
X is sp3 hybridized.
XF5
X is dsp3 hybridized.
Axial bonds in XF5 are longer than the equatorial bonds. The axial bonds are p/d hybrids and the
equatorial bonds are s/p2 hybrids. Another acceptable explanation is that the axial bonds at 90o are
repelled more than the equatorial bonds at 120o.
Y could be N; N has no d orbitals. Another acceptable explanation is that N is too small to
accommodate five F atoms.
Y could be Bi; Bi can’t easily be oxidized to +5.
Key for 2001 USNCO National Exam, Part II
Page 3
8.
a.
CO2 is linear and therefore nonpolar.
O C O
S
SO2 is bent and therefore polar. O
b.
c.
d.
Page 4
O
The polar substance will bond more strongly and have the lower vapor pressure.
The boiling point of HCl is less than the boiling point of HF because HF forms hydrogen bonds
which are harder to break than van der Waals forces.
Both HCl and HBr are attracted by van der Waals forces. However, HBr has more electrons and
therefore has stronger van der Waals forces. As a result, the boiling point of HCl is less than the
boiling point of HBr.
Ca2+ and O2– ions are attracted about four times as strongly as K+ and F– ions. Ions with a +2
charge are attracted more strongly than ions with a +1 charge. In addition, the calcium-to-oxygen
distance is less than the potassium-to-fluoride distance, leading to an increased force of attraction
for the shorter bond.
Tin(II) chloride is ionic. Tin(IV) chloride is covalent. The +4 charge on tin causes it to attract
electrons more strongly from chloride ion, making the bonds covalent.
Key for 2001 USNCO National Exam, Part II
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Lucy Pryde Eubanks, Clemson University, Clemson, SC
Chair
Robert Becker, Kirkwood High School, Kirkwood, MO
Craig W. Bowen, US Naval Academy, Annapolis, MD
Nancy Devino, ScienceMedia Inc., San Diego, CA
Sheldon L. Knoespel, Michigan State University, East Lansing, MI
Steve Lantos, Brookline High School, Brookline, MA
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Robert G. Silberman, SUNY-Cortland, NY
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 1
2001 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
Design and carry out an experiment to determine the density of the plastic object you have been given. You
may use water and the alcohol solution provided at your lab station, as well as the equipment you will find
there, but you may not use a balance. You will be asked to describe the method you developed to solve this
problem.
Given:
density of water
= 1.00 g·mL–1
density of alcohol solution = 0.85 g·mL–1
Lab Problem 2
Design and carry out an experiment to determine the specific identity of the compound in each of eight
numbered vials. Each vial contains one of these ionic compounds.
BaCl2, CaCO3, Ca(OH)2, KI, NaCl, NaHCO3, Na2SO4, Pb(NO3)2
In addition to the equipment you will find at your lab station, you may use distilled water. You also have the
option of choosing ONE additional reagent from this list. You may do this either before or during your
experimentation.
6 M H2SO4, 6 M HCl, 6 M NaOH, phenolphthalein indicator solution
You will be asked to describe the method you developed to solve this problem.
Page 2
Not valid for use as an USNCO National Examination after April 22, 2001.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School: ________________________________________Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name : _______________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 3
2. Record your data and other observations.
3. What is the density of the plastic object? Show your methods clearly.
Page 4
Not valid for use as an USNCO National Examination after April 22, 2001.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School: ________________________________________Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name : _______________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
When you wish to request the optional reagent, return to the Examiner with this sheet.
I request this additional reagent: ______________
Examiner’s Initials:
2. Record your data and other observations.
Not valid for use as an USNCO National Examination after April 22, 2001.
Page 5
2. Record your data and other observations. (continued)
3. Identify the substance in each numbered vial, giving a brief justification for that choice.
Vial #
Contains
Justification
1.
2.
3.
4.
5.
6.
7.
8.
Page 6
Not valid for use as an USNCO National Examination after April 22, 2001.
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Directions
Thank you for administering the 2001 USNCO laboratory practical on behalf of your Local
Section. It is essential that you follow the instructions provided, in order to insure consistency of results
nationwide. There may be considerable temptation to assist the students after they begin the lab exercise.
It is extremely important that you do not lend any assistance or provide any hints whatsoever to the
students once they begin work. As is the case with the international competition, students should not be
allowed to speak to anyone until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab
station or table when the students enter the room. The equipment should be initially placed so that the
materials used for Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. Chemistry Olympiad
National Examination. In this part of the exam, we will be assessing your lab skills and your ability to
reason through a laboratory problem and communicate your results. Do not touch any of the equipment
in front of you until you are instructed to do so.
One of this year’s problems uses this type of plastic screw anchor as the object being investigated.
Show the type of plastic screw anchors being used at your site. (See picture on page 3 of these
instructions for approximate size of the anchors.)
Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques
help to minimize the amount of materials you use, thereby increasing safety and minimizing waste.
Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and
reaction plates.
Show a 1-mL, 3-mL, or a 5-mL Beral-type pipet, and show a 6-well or a 12-well reaction plate.
(If your testing site has substituted 50-mL or 100-mL beakers, show those instead.)
You will be asked to complete two laboratory problems. The materials and equipment needed to solve
each problem has been set out for you and is grouped by the number of the problem. You also may use
distilled (or deionized) water. You must limit yourself to this equipment and materials for each problem.
A balance may not be used for either problem. You may choose to start with either problem. You are
required to have a procedure for each problem approved for safety by an examiner. (Remember that
approval does not mean that your procedure will be successful–it is a safety approval.) When you are
ready for an examiner to come to your station for each safety approval, please raise your hand.
In the second problem, you have the option of choosing ONE additional reagent, either before or
during your experimentation. Again, when you are ready to make this choice, please write the formula
Examiner’s Directions, 2001 USNCO National Exam, Part III
Page 1
of the reagent being requested on your report sheet, and raise your hand. You will have one hour and
thirty minutes to complete both problems.
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash
off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate
procedures for disposing of solutions at the end of this lab practical are:
____________________________________________________________________________________
____________________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. Are there any questions before we begin?
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. Allow students
enough time to read the brief cover directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out
all information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for
your cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the
lab practical with the students. They can learn about possible observations and interpretations and you
can acquire feedback as to what they actually did and how they reacted to the problems. After this
discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be
extremely useful to the Olympiad Laboratory Practical subcommittee as they prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron
sheets from Part I, and the “Blue Books” from Part II to this address:
ACS DivCHED Exams Institute
Clemson University
223 Brackett Hall
Box 340979
Clemson, SC 29634-0979
Wednesday, April 25, 2001 is the absolute deadline for receipt of the exam materials at the
Examinations Institute. Materials received after this deadline CANNOT be graded.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT
SCHEDULE FOR GRADING THIS EXAMINATION.
Page 2
Examiner’s Directions, 2001 USNCO National Exam, Part III
EXAMINER’S NOTES
Lab Problem #1: Materials and Equipment
Note: Students will NOT be allowed to use a balance for this lab problem. Be sure none are
available in the testing area or secure them so they may not be used.
Each student will need:
2 10-mL graduated cylinders
2 small beakers (100 mL or 250 mL)
one labeled “water”, one labeled “alcohol solution”
2 1-mL Beral-style pipets (Eye droppers may be substituted.)
4 to 6 test tubes, 13 x 100 mm or larger
1 test tube rack
1 250-mL squeeze bottle, labeled “distilled water” or “deionized water”
2 plastic screw anchors (Check your local hardware
store for No. 8-10 x 7/8”. These are 2 cm long
with a maximum diameter of approximately
0.5 cm. Other sizes varying from No. 4-6 to 10-12
may be used but be sure to check that these sizes
fit easily into the test tubes being used.)
Lab Problem #1: Chemicals. Each student will need:
250 mL of distilled or deionized water
250 mL of 70% isopropyl alcohol
Note: This is sold as “rubbing alcohol” in most stores or pharmacies. You may wish to provide
each student with an unopened bottle to emphasize the use of a consumer product. Choose the
cheapest brand and check there are no additives such as dyes or perfumes that will change the
density. Do not purchase 91% or 99% isopropyl alcohol; these are often available as well. A less
desirable alternative, one that does not emphasize the use of a consumer chemical, is to prepare a
70% by volume solution from pure isopropyl alcohol and water, and provide the solution in a
250-mL labeled squeeze bottle.
Quick Check to be sure lab problem #1 will work for your examinees:
1) Does the screw anchor fit into the size test tubes being provided?
2) Does the screw anchor float in water, and sink in the alcohol solution?
3) Have you planned to prevent access to all balances in the working area?
Lab Problem #1: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not
comment on the plan other than looking for any potentially unsafe practices.
2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give
students explicit directions for handling spills and for disposing of waste materials, following approved
safety practices for your examining site. Please check and follow procedures appropriate for your site.
Examiner’s Directions, 2001 USNCO National Exam, Part III
Page 3
Lab Problem #2: Materials and Equipment. Each student will need:
8 numbered small vials with tops (30-mL plastic vials work well)
1 100-mL or larger wash bottle, labeled “distilled water” or “tap water”
1 24-well reaction plate or 2 12-well plates. If reaction plates cannot be obtained or borrowed,
6 50-mL or 100-mL beakers can be substituted.
stirring sticks such as wooden or plastic toothpicks, or coffee stirers
8 1-mL Beral-style pipets, cut to use as scoops (or 8 small spatulas or scoops)
2 1-mL Beral-style pipets (Eye droppers may be substituted.)
1 1-mL Beral-style pipet with label (This will contain 6 M H2SO4, 6 M HCl,
6 M NaOH, or phenolphthalein indicator solution)
1 container designated for disposal of heavy metal waste of Pb2+ and Ba2+
supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
Lab Problem #2: Chemicals. Each student will need:
1 set of filled, numbered vials. Each numbered vial will contain about 1 g of one of these dry solids.
Note: This is the order for filling the numbered vials. Please keep this list secure.
1. NaCl
4. BaCl2
7. Pb(NO3)2
2. CaCO3
5. NaHCO3
8. KI
3. Na2SO4
6. Ca(OH)2
Please have available 100 mL of each of these reagents: 6 M H2SO4, 6 M HCl, 6 M NaOH, and
phenolphthalein indicator solution. Students will be asked to choose ONE of these reagents for use with
lab problem #2, either before starting experimentation or during their work. You may find it convenient
to pre-fill a set of labeled Beral-style pipets for each student but they must not be supplied at the lab
station.
Supply of distilled water, if available; deionized water may also be used
Quick Check to be sure this lab problem will work for your examinees:
1) Are all the solids dry?
2) CaCO3 needs to be provided in powdered form, not as marble chips that are sometimes used.
3) Are you prepared to dispense H2SO4, HCl, NaOH, or phenolphthalein indicator solution quickly and
safely when the students have made their choice?
4) Are you prepared to collect all solutions containing Pb2+ and Ba2+ metal ions?
Lab Problem #2: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment
on the plan other than looking for any potentially unsafe practices. The examiner also will need to initial each
student’s choice of additional reagent.
2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give
students explicit directions for handling spills and for disposing of waste materials, following approved safety
practices for your examining site. Please check and follow procedures appropriate for your site.
Page 4
Examiner’s Directions, 2001 USNCO National Exam, Part III
2001 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART III
Lab Problem 1
Either of two general plans might be used to solve this problem. Multiple trials are expected for either plan used.
Plan A: Make a solution of water and alcohol in which the object is suspended. Using the measured volumes of the water
and alcohol, and their given densities, the density of the solution can be found. The density of the solution and the
object will be the same.
Plan A Sample Data:
Trial #1
Trial #2
Volume water to suspend object
2.21 mL
2.23 mL
Volume alcohol to suspend object
3.02 mL
3.03 mL
Plan A Calculations:
(
)
Mass of solution = VH2 O × densityH 2O + ( Valcohol × densityalcohol )
(
)
Volume of solution = VH2 O + (Valcohol ) (assuming volumes are additive)
Mass of solution
Density of solution =
Volume of solution
Plan A Sample Results:
Trial #1
Trial #2
Mass of solution (g)
4.78
4.81
Volume of solution (mL)
5.23
5.26
Density of solution (g/mL)
0.91
0.91
Plan B: Using Archimedes’ principle, determine the volume of the object by displacing the alcohol. Determine the mass of
the object by displacing water, and multiplying the volume of the water displaced by the density of water. Determine
the density of the object by dividing the mass of the object by its volume.
Plan B Sample Data:
Trial #1
Trial #2
Initial volume water
5.02 mL
5.14 mL
Final volume water
5.83 mL
5.94 mL
Initial volume alcohol
6.01 mL
Final volume alcohol
6.88 mL
Plan B Calculations:
Mass of object = VH 2 O displaced × density H2 O
Mass of object
Density of object =
Volume of object
Plan B Sample Results:
Trial #1
6.02 mL
6.89 mL
Trial #2
Mass of object (g)
0.81
0.80
Volume of object (mL)
0.87
0.87
Density of object (g/mL)
0.93
0.92
Conclusion: The determined density should be between 0.85 g/mL and 1.00 g/mL, because the object sinks in alcohol and
floats in water. The density determination must be supported by data gathered and calculations performed.
Page 1
2001 USNCO National Exam, Part III (Lab Practical)
Lab Problem 2
Credit was awarded for alternate, logical pathways that achieved identification of the compounds. The identifications depend
on developing a logical sequence of tests that will lead to the identifications. Some type of tabular form organizes
the data for clear presentation, even if a formal flow chart is not included.
Sample Plan: Many students started by adding water to each compound. Most often, H2SO4 was chosen as the extra reagent
Those tests were followed by adding selected solutions of the unknown to each other to help with identification.
Sample Data:
FIRST
#1
#2
#3
#4
#5
#6
#7
#8
TESTS
H2O
sol
insol
sol
sol
sol
insol
sol
sol
H2SO4
no rxn
dissolves,
no rxn
white ppt
bubbles
dissolves
white ppt
yellow
bubbles
The first set of tests allows identification of #2 as CaCO3, #5 as NaHCO3, and #6 as Ca(OH)2.
SECOND
TESTS
#1
#3
–
no rxn
–
#1
#3
#4
#4
no rxn
white ppt
–
#7
#8
#7
#8
white ppt
white ppt
no rxn
no rxn
white ppt
no rxn
–
yellow ppt
–
The second set of tests allows identification of #1 as NaCl, #7 as Pb(NO3)2, and #8 as KI. There is still ambiguity
about #3 and #4 at this point.
THIRD
TESTS
H2SO4
#3
no rxn
#4
white ppt
The third set of tests allows identification of #3 as Na2SO4 and #4 as BaCl2.
Conclusion:
Identification
Page 2
Substance
Supporting Evidence
#1
NaCl
Second tests: forms a white precipitate with #7, Pb(NO3)2
#2
#3
#4
CaCO3
Na2SO4
BaCl2
First tests: insoluble in water, dissolves and bubbles in acid
Third tests: no reaction with H2SO4
Third tests: white precipitate forms
#5
NaHCO3
First tests: soluble in water, dissolves and bubbles in acid
#6
#7
Ca(OH)2
Pb(NO3)2
First tests: insoluble in water, dissolves in acid
Second tests: forms three white precipitates and one yellow precipitate
#8
KI
Second tests: forms one yellow precipitate
2001 USNCO National Exam, Part III (Lab Practical)
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
Dianne H. Earle, Paul M. Dorman High School, SC
David W. Hostage, Taft School, CT
Alice Johnsen, Bellaire High School, TX
Elizabeth M. Martin, College of Charleston, SC
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 21, 2002, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after 1 hour, 30 minutes has elapsed, the
student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of 1 hour, 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer
sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 21, 2002.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
atomic number
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
M
G molar
A frequency
mol
ν mole
Z gas constant
h
R Planck’s constant
N A gram
pressure
P
g
°C heat capacity
k
C p rate constant
c hour
Rf
h retention factor
C joule
s
J second
E kelvin
speed
of
light
c
K
Ea kilo- prefix
temperature,
K
T
k
H liter
t
L time
S
volt
V
E =E –
USEFUL EQUATIONS
 –∆H   1 
ln K = 
   +c
 R T
RT
lnQ
nF
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol –1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
1 atm = 760 mmHg
 k 2  Ea  1 1 
=
−
 k1  R  T1 T2 
ln 
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(289)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as a USNCO National Exam after April 21, 2002.
DIRECTIONS
§ When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
§ You may write on the test booklet, but it will not be used for grading.
§ There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
§ Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which element commonly exhibits both +1 and +3
oxidation states?
(A) Al (Z = 13)
(B) Sc (Z = 21)
(C) Sn (Z = 50)
(D) Tl (Z = 81)
6. A weighed quantity of a gas is collected over water at
25 °C and 742 mmHg. The molar mass of the gas is to be
determined at standard temperature and pressure. If the
vapor pressure of water is ignored during the calculation,
what is the effect on the calculated pressure and
calculated molar mass of the gas?
pressure
2. Which procedure is best to extinguish burning
magnesium?
molar mass
(A)
low
low
(A) Add water to it.
(B)
low
high
(B) Blow nitrogen gas over it.
(C)
high
low
(C) Cover it with sand.
(D)
high
high
(D) Throw ice on it.
7. A 0.1 M solution of which substance is most acidic?
3. Which two sets of reactants best represent the
amphoterism of Zn(OH)2?
Set 1. Zn(OH)2(s) and OH–(aq)
Set 2. Zn(OH)2(s) and H2O(l)
Set 3. Zn(OH)2(s) and H+(aq)
Set 4. Zn(OH)2(s) and NH3(aq)
(A) Sets 1 and 2
(B) Sets 1 and 3
(C) Sets 2 and 4
(D) Sets 3 and 4
(A) NaHSO 4
(B) Na 2SO4
(C) NaHS
(D) NaHCO3
8. The mineral trona has the formula Na 2CO3. NaHCO3. 2H2O
and a formula mass of 226 g. mol –1. How many mL of
0.125 M HCl are needed to convert all the carbonate and
bicarbonate in a 0.407 g sample of trona into carbon
dioxide and water?
4. Which of these statements about sulfur is not correct?
(A) 43.2 mL
(B) 28.8 mL
(C) 21.6 mL
(D) 14.4 mL
(A) It exists in different allotropic forms.
(B) It can behave as either an oxidizing agent or a
reducing agent.
9. The percentages by mass of C, H, and Cl in a compound
are C 52.2%, H 3.7%, and Cl 44.1%. How many carbon
atoms are in the simplest formula of the compound?
(C) It can form up to six covalent bonds in compounds.
(A) 3
(B) 4
(C) 6
(D) 7
(D) It is a liquid at 25 °C and 1 atm pressure.
10.
5. A solution of sulfuric acid in water that is 25% H2SO4 by
mass has a density of 1.178 g·mL–1. Which expression
gives the molarity of this solution?
(A) 0.25 × 98 × 1178
(B)
0.25 ×1178
98
0.25
98 × 1178
(D)
1178
0.25 × 98
(C)
4KO2(s) + 2CO2(g) → 2K 2CO3(s) + 3O2(g)
What is the maximum volume of oxygen that can be
produced when 150. mL of CO2 is passed over 0.500 g
of KO2? Assume all gases are measured at 0 °C and
1 atm.
(A) 118 mL
(B) 157 mL
(C) 225 mL
(D) 475 mL
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 3
11. The first vertical line in the diagram represents a
thermometer with the boiling and freezing points for a
pure solvent. The numbered lines represent possible
boiling and freezing points for a solution of a nonvolatile
solute in the same solvent. Which line best represents the
boiling point and freezing point of a solution relative to
values for the pure solvent?
Note: The differences in temperatures are not to scale.
Solvent
1
2
3
4
16. When the temperature of a sample of H2S gas is lowered,
the pressure decreases more than predicted by the ideal
gas equation. To what is this deviation from expected
behavior due?
1. attractive forces between molecules
2. mass of the molecules
3. volume of the molecules
(A) 1 only
(B) 2 only
(C) 1 and 3 only
(D) 2 and 3 only
bp
fp
(A) 1
(B) 2
(C) 3
(D) 4
12. Equal masses of gaseous N2, NH3, and N2O are injected
into an evacuated container to produce a total pressure of
3 atm. How do the partial pressures of N 2, NH3, and N2O
compare?
17. This curve is
produced when a
pure substance is
heated. Which
characteristic of this
curve is related to the
value for the
enthalpy of fusion of
the substance?
F
D
B
C
A
Heat added
(A) length of AB
(B) length of BC
(C) slope of AB
(D) slope of CD
18. Which statement is correct?
(A) PN 2 = PNH3 = PN2 O
(B) PN 2 < PNH 3 < PN2 O
(A) In a coffee-cup calorimeter, q = ∆H.
(C) PNH 3 < PN 2 < PN2 O
(D) PN 2O < PN 2 < PNH3
(B) In a coffee-cup calorimeter, w = 0.
(C) In a bomb calorimeter, q = ∆S.
(D) In a bomb calorimeter, w > 0.
13. According to this phase
diagram, which phases can
exist at pressures lower
than the triple point
pressure?
19. Consider this reaction.
4PH3(g) + 8 O2(g) → P4O10(s) + 6H2O(g) ∆H o = –4500 kJ
Calculate ∆H of of P4O10(s) in kJ·mol–1.
Temperature →
(A) gas only
(B) solid and gas only
(C) liquid only
(D) solid and liquid only
14. 1.00 g of water is
Vapor Pressure, 50 °C
introduced into a 5.00 L
H2O
92.5 mmHg
evacuated flask at 50 °C.
What mass of water is present as liquid when equilibrium
is established?
(A) 0.083 g
(B) 0.41 g
(C) 0.59 g
(D) 0.91 g
15. Which substance has the greatest lattice energy?
(A) NaF
Page 4
(B) KCl
E
(C) MgO
Substance
PH3(g)
H2O(g)
∆H of , kJ·mol–1
+9.2
–241.8
(A) –5914 kJ
(B) –4751 kJ
(C) –4249 kJ
(D) –3012 kJ
20. For which substances and conditions can So = 0?
I. elements at 0 K
II. compounds at 0 K
III. gases at 298 K
(A) I only
(B) III only
(C) I and II only
(D) I and III only
(D) CaS
Not valid for use as an USNCO National Examination after April 21, 2002.
21. 50.0 mL of 0.10 M HCl is
Solution Values
mixed with 50.0 mL of
Cp
4.18 J·g–1·°C–1
0.10 M NaOH. The
solution temperature rises
density 1.0 g·mL–1
by 3.0 °C. Calculate the
enthalpy of neutralization per mole of HCl.
(A) –2.5 × 102 kJ
(B) –1.3 × 102 kJ
(C) –8.4 × 10 kJ
(D) –6.3 × 10 kJ
1
1
27. The rate of a reaction at 75 °C is 30.0 times that at 25 °C.
What is its activation energy?
(A) 58.6 kJ·mol–1
(B) 25.5 kJ·mol–1
(C) 7.05 kJ·mol–1
(D) 1.51 kJ·mol–1
28. 6I–(aq) + BrO3–(aq) + 6H+(aq) → 3I 2(aq) + Br –(aq) + 3H2O(l)
These data were obtained when this reaction was studied.
[I –], M
22. What can be
concluded about the
values of ∆H and
∆S from this graph?
+100
[BrO 3–], M
0.0010
0.0020
0.0020
0.0010
+50
0
–50
–100
0 100 200 300 400 500
Temperature, K →
(A) ∆H > 0, ∆S > 0
(B) ∆H > 0, ∆S < 0
(C) ∆H < 0, ∆S > 0
(D) ∆H < 0, ∆S < 0
23. The boiling point of chloroform, CHCl3, is 61.7 °C and
its enthalpy of vaporization is 31.4 kJ·mol–1. Calculate
the molar entropy of vaporization for chloroform.
(A) 10.7 J·mol–1·K–1
(B) 93.8 J·mol–1·K–1
(C) 301 J·mol–1·K–1
(D) 509 J·mol–1·K–1
(B) 1.1
(C) 0.86
(D) 4.5 × 10–6
0.010
0.010
0.010
0.020
Reaction rate,
mol·L–1·s–1
8.0 × 10–5
1.6 × 10–4
1.6 × 10–4
1.6 × 10–4
What are the units of the rate constant for this reaction?
(A) s–1
(B) mol·L –1·s–1
(C) L·mol–1·s–1
(D) L2·mol–1·s–1
29. Consider this gas phase reaction.
Cl2(g) + CHCl3(g) → HCl(g) + CCl4(g)
The reaction is found experimentally to follow this rate
law.
rate = k [CHCl3] [Cl 2]1/2
Based on this information, what conclusions can be
drawn about this proposed mechanism?
Step 1. Cl2(g) ¾ 2Cl(g)
Step 2.
Step 3.
24. ∆Go for a reaction at 25 °C is 30.5 kJ·mol–1. What is the
value of K?
(A) 2.2 × 105
0.0020
0.0020
0.0040
0.0040
[H +], M
Cl(g) + CHCl3(g) → HCl(g) + CCl3(g)
Cl(g) + CCl 3(g) → CCl 4(g)
(A) Step 1 is the rate-determining step.
(B) Step 2 is the rate-determining step.
(C) Step 3 is the rate-determining step.
25. This is the rate law for a reaction that consumes X.
rate = k [X]2
Which plot gives a straight line?
(A) [X] vs. time
(B) ln [X] vs. time
(C) 1 / [X] vs. time
(D) 1 / ln [X]2 vs. time
26. For a first order reaction, the concentration decreases to
30% of its initial value in 5.0 min. What is the rate
constant?
(A) 0.46 min –1
(B) 0.24 min –1
(C) 0.14 min –1
(D) 0.060 min –1
(D) The rate-determining step cannot be identified.
30. Determine the value of the equilibrium constant for this
reaction
2NOCl(g) + O2(g) ¾ 2NO2(g) + Cl2(g)
from the K values for these reactions.
2NOCl(g) ¾ 2NO(g) + Cl2(g)
K p = 1.7 × 10 –2
2NO2(g) ¾ 2NO(g) + O2 (g)
K p = 5.9 × 10–5
(A) 1.0 × 10–6
(B) 1.0 × 10–3
(C) 3.5 × 10–3
(D) 2.9 × 102
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 5
31. What is the pH of a 0.15 M
solution of hydrazine, N2H4?
(A) 3.41
(B)
6.82
Hydrazine
Kb
N2H4
1.0 × 10–6
(C) 10.59
(D) 11.00
38. How many moles of electrons must be removed from
each mole of toluene, C6H5CH3, when it is oxidized to
benzoic acid, C6H5COOH?
(A) 1
(B) 2
(C) 4
(D) 6
Questions 39 and 40 refer to the reaction represented by this
equation.
2Al(s) + 3Cu2+(aq) → 2Al3+(aq) + 3Cu(s)
39. What is the value of Eo for a voltaic cell based on this
reaction?
32. The rates of many
catalyzed reactions follow
the profile shown in the
graph. Why does the
reaction rate level off?
[Reactant], M →
(A) The reactant is used up.
(B) The reverse reaction becomes dominant.
Reaction
Cu (aq) + 2e–→ Cu(s)
Al3+(aq) + 3e–→ Al(s)
2+
Eo
+0.34 V
–1.66 V
(A) 1.32 V
(B) 2.00 V
(C) 2.30 V
(D) 4.34 V
(C) The catalyst decomposes as the reaction proceeds.
(D) The active sites on the catalyst are occupied.
Questions 33 and 34 refer to aqueous solutions of formic acid,
HCOOH, which has a Ka value of 1.9 × 10 –4 at 25 °C.
33. What is the percent ionization of a 0.10 M solution of
formic acid at 25 °C?
(A) 0.19%
(B) 1.4%
(C) 4.4%
(D) 14%
34. How many moles of sodium formate must be added to
1.0 L of a 0.20 M formic acid solution to produce a pH
of 4.00?
(A) 0.38
(B) 0.80
(C) 1.9
(D) 3.8
35. During the titration of a weak base with a strong acid,
one should use an acid-base indicator that changes color
in the
(A) acidic range.
(B) basic range.
(C) buffer range.
(D) neutral range.
36. What is the solubility
Substance
K sp
of calcium hydroxide
calcium hydroxide 4.0 × 10 –6
in mol·L–1?
(A) 1.6× 10
(B) 1.0 × 10
–2
(C) 2.0 × 10
–2
Page 6
(A) 6
(B) 5
(C) 3
(D) 2
41. Use the given standard reduction potentials to determine
the reduction potential for this half-reaction.
MnO4–(aq) + 3e– + 4H+ → MnO 2(s) + 2H2O(l)
Eo
Reaction
MnO4–(aq) + e–→ MnO 42–(aq)
+0.564 V
MnO42–(aq) + 2e– + 4H+ → MnO 2(s) + 2H2O(l) +2.261 V
(A) 1.695 V
(B) 2.825 V
(C) 3.389 V
(D) 5.086 V
42. How many Faradays are required to reduce all the
chromium in 0.150 L of 0.115 M of Cr 2O72– to Cr2+?
(A) 0.920 F
(B) 0.690 F
(C) 0.138 F
(D) 0.069 F
43. In which list are the elements arranged in order of
increasing first ionization energy?
(A) Li, Na, K
(B) S, O, F
(C) Na, Mg, Al
(D) F, Ne, Na
(D) 1.0 × 10–3
–3
37. What is the average oxidation number of tungsten in the
ion, W6O6Cl122–?
(A) 2.7
40. What value should be used for n in the Nernst equation to
determine the effect of changes in Al3+(aq) and Cu2+(aq)
concentrations in this reaction?
(B) 3.3
(C) 3.7
44. Which quantum number is associated with the shape of
an atomic orbital?
(A) n
(B) l
(C) ml
(D) ms
(D) 4.3
Not valid for use as an USNCO National Examination after April 21, 2002.
45. Consider the ions Li+, Na +, Be2+, and Mg2+. Which two
are closest to one another in size?
(A) Li+ and Na+
(B) Be2+ and Mg2+
(C) Be2+ and Li+
(D) Li+ and Mg2+
46. What is the electron configuration for a gas phase +3 ion
of iron (Z = 26)?
(A) [Ar] 3d
5
1
(C) [Ar] 4s 3d
2
(B) [Ar] 4s 3d
4
52. Which species has the strongest oxygen-oxygen bond
according to molecular orbital theory?
(A) O2
(B) O2–
(C) O22–
(D) O2+
53. How many atoms are covalently bonded to the chromium
atom in Cr(NH 3)4Cl3?
(A) 3
(B) 4
(C) 6
(D) 7
3
(D) [Ar] 4s2 3d 6
47. Magnesium (Z = 12) has isotopes that range from Mg–20
to Mg–31. Only Mg–24, Mg–25, and Mg–26 are not
radioactive. What mode of radioactive decay would
convert Mg–20, Mg–21, Mg–22, and Mg–23 into stable
isotopes most quickly?
(A) electron emission
(B) alpha particle emission
(C) gamma emission
(D) positron emission
48. Which oxides exist as individual molecules?
1. Al2O3
2. SiO 2
3. P 4O10
54. When the carbon-oxygen bonds in H3COH, H2CO, and
HCO2– are arranged in order of increasing length, what is
the correct order?
(A) H3COH, H2CO, HCO2–
(B) HCO2–, H3COH, H2CO
(C) H2CO, HCO2–, H3COH
(D) H3COH, HCO2–, H2CO
55. Which reaction is an oxidation? (Only the carboncontaining molecules are shown.)
(A) CH2CH2 → CH3CH2OH
(A) 2 only
(B) 3 only
(B) CH3CH2OH → CH 2CHO
(C) 1 and 3 only
(D) 2 and 3 only
(C) CH3CH2OH + HCOOH → CH3CH2OOCH
(D) 2 CH3CH2OH → CH3CH2OCH2CH3
49. How many sigma and pi
bonds are in this
compound?
O
C OH
(A) 9 sigma, 6 pi
(B) 10 sigma, 6 pi
(C) 10 sigma, 3 pi
(D) 15 sigma, 4 pi
Use this structure for the indigo
molecule to answer questions 56
and 57.
O
C
C
N
H
H
N
C
C
O
56. What is the molecular formula of indigo?
50. Which pair of ions has the same shape?
(A) CO32– and NO3–
(B) CO32– and SO32–
(C) NO3– and ClO3–
(D) CO32– and ClO3–
51. Which resonance form makes the greatest contribution to
the structure of N2O?
(A)
(C)
N
N
N
N
O
O
(B)
(D)
N
N
(A) C 8HNO
(B) C 16H2N2O2
(C) C 16H10N2O2
(D) C 16H22N2O2
57. What is the hybridization of the carbon atoms bonded to
oxygen?
(A) sp
(B) sp2
(C) sp3
(D) sp3d
N O
N O
58. Aniline, C 6H5NH2, does not dissolve well in water.
Which reagent could be used to increase its aqueous
solubility?
(A) 1 M HCl
(B) 1 M NaOH
(C) diethyl ether
(D) toluene
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 7
59. Which molecule reacts most rapidly with water?
60. Which of these elements is found in hemoglobin?
(A) CH3CH2CH2Cl
(B) CH3CHClCH3
(C) (CH3)2CHCH2Cl
(D) (CH3)3CCl
(A) Cr
(B) Fe
(C) Mg
(D) Ni
END OF TEST
Page 8
Not valid for use as an USNCO National Examination after April 21, 2002.
U. S. National Chemistry Olympiad – 2002
National Examination—Part I
SCORING KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Answer
D
C
B
D
B
C
A
A
D
A
D
D
B
C
C
A
B
A
D
C
Number
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
Answer
A
A
B
D
C
B
A
C
B
D
C
D
C
A
A
B
C
D
B
A
Property of the ACS Society Committee on Education
Number
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
A
C
B
B
D
A
D
B
D
A
B
D
C
C
B
C
B
A
D
B
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
Dianne H. Earle, Paul M. Dorman High School, SC
David W. Hostage, Taft School, CT
Alice Johnsen, Bellaire High School, TX
Elizabeth M. Martin, College of Charleston, SC
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a
score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 21, 2002, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after 1 hour, 45 minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front
of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. 1 hour, 45 minutes is allowed to complete this part. Be sure to print your name, the name of your
school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification
number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper.
Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you
complete Part II (or at the end of 1 hour, 45 minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do
not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 21, 2002.
Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A.
1.
2.
(12%) The percentages of NaHCO3 and Na2CO3 are to be determined in a
mixture of them with KCl. A 0.500 g sample of the mixture is dissolved
in 50.0 mL of deionized water and titrated with 0.115 M HCl, resulting in
this pH titration curve.
a. Write a balanced equation for the reaction that is responsible for the
equivalence point that occurs at about
i. pH = 9
ii. pH = 5
b. Calculate the total number of moles of acid used to reach each
equivalence point if the volumes are 9.63 mL and 34.27 mL,
respectively.
c. Determine the number of grams of Na2CO3 and NaHCO3 and their
percentages in the original mixture.
d. Sketch a titration curve for a solution of Na2CO3 by itself and describe
how it differs from the given curve.
(15%) Consider the formation of N2O5(g) by this reaction.
2NO2(g) + 1 /2O2(g) → N 2O5(g)
For this reaction, ∆H o = –55.1 kJ and ∆S o = –227 J·K–1
Additional data are given in the Table.
14
12
10
8
6
4
2
0
0
10.0
20.0
30.0
Volume HCl(aq), mL →
Type of Data
Substance
∆Hf o
NO2(g)
+33.2 kJ·mol –1
So
NO2(g)
239.7 J·mol–1·K–1
So
O2(g)
205.1 J·mol–1·K–1
a. Calculate these values.
40.0
Value
i. ∆Hf o of N2O5(g)
ii. S o of N2O5(g)
iii. ∆G o of the given reaction at 25 °C iv. K p of the given reaction at 25 °C
b. State and explain
i. whether this reaction is spontaneous at 25 °C.
ii. how the numerical value of K p would be affected by an increase in temperature.
iii. how the relative amounts of reactant and product molecules would be affected by an increase in temperature.
iv. why the S o values differ for NO2(g) and O2(g) at 25 °C.
3. (11%) Calcium ions form slightly soluble compounds with phosphate ions such as PO43–, HPO42–, and H2PO4–.
a. Write the formula and give the K sp expression for the compound formed by Ca2+ and each of these two ions.
i. PO43–
ii. H2PO4–
b. Calculate the equilibrium concentration of Ca 2+ in a saturated solution with each of the phosphate ions given in part a.
i. K sp for calcium phosphate equals 1.0 × 10–25.
ii. K sp for calcium dihydrogen phosphate equals 1.0 × 10–3.
+
c. Determine the [H ] needed to just prevent precipitation by H2PO4– in a 0.25 M H3PO4 solution that has [Ca2+] = 0.15 M.
The Ka1 of H 3PO4 is 7.1 × 10 –3.
4.
(14%) This reaction can be used to analyze for iodide
ion.
[I –], M
[IO3–], M
[H +], M
Reaction rate,
mol·L–1·s–1
IO3–(aq) + 5I–(aq) + 6H+(aq) → 3I2(aq) + 3H2O(l)
0.010
0.10
0.010
0.60
When the rate of this reaction was studied at 25 °C, the
0.040
0.10
0.010
2.40
results in the table were obtained.
0.010
0.30
0.010
5.40
a. Use these data to determine the order of the reaction
0.010
0.10
0.020
2.40
with respect to each of these species. Outline your
reasoning in each case.
i. I–
ii. IO 3–
iii. H +
b. Calculate the specific rate constant for this reaction and give its units.
c. Based on the kinetics, discuss the probability of this reaction occurring in a single step.
d. The kinetics of reactions are often studied under pseudo first-order conditions. Describe what is meant by the term pseudo first
order and illustrate how the reaction conditions above would be changed so that the [I–] would be pseudo first order.
e. The activation energy for this reaction was found to be 84 kJ·mol –1 at 25 °C. How much faster would this reaction proceed if the
activation energy were lowered by 10 kJ·mol–1 (for example, by using a suitable catalyst)?
Page 2
Not valid for use as an USNCO National Exam after April 21, 2002
5.
(12%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the reactants and
products. Omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic
substances. You need not balance the reactions. All reactions occur in aqueous solution unless otherwise indicated.
a. Concentrated hydrochloric acid is added to solid manganese(IV) oxide.
b. Solutions of magnesium sulfate and barium hydroxide are mixed.
c. Solid barium peroxide is added to water.
d. A piece of copper metal is added to a solution of dilute nitric acid.
e. A solution of sodium thiosulfate is added to a suspension of solid silver bromide.
f. 2-butanol is heated with a solution of acidified potassium dichromate.
6.
(12%) Chlorine trifluoride, ClF 3, is a vigorous fluorinating agent that has been used to separate uranium from the fission
products in spent nuclear fuel rods.
a. Write a Lewis dot structure for ClF3.
b. Sketch and describe clearly the geometry for the ClF3 molecule.
c. Sketch one other possible geometry and explain why it is not observed.
d. Identify the hybrid orbitals that are considered to be used by the chlorine atom in ClF3.
e. The electrical conductance of liquid ClF3 is only slightly lower than that of pure water. This behavior is attributed to the
self-ionization of ClF3 to form ClF2+ and ClF4–. Sketch and describe the expected structures of ClF2+ and ClF4–.
7.
(12%) Answer these questions about the voltaic cell based on these half-reactions.
MnO2(s) + 4H+(aq) + 2e– → Mn2+(aq) + 2H2O(l)
E o = +1.23 V
2+
+
–
TiO (aq) + 2H (aq) + 4e → Ti(s) + H2O(l)
E o = –0.88 V
a. Write the equation for the reaction that produces a positive standard potential and then calculate that potential.
b. Identify the half-reaction that occurs at the cathode. Explain.
c. Specify the conditions that produce the standard potential.
d. State whether each of the changes listed in parts i - iii will affect the standard potential calculated in part a for the assembled
cell. For each change, state whether the potential will increase, decrease, or remain the same. Outline your reasoning or show
your calculations in each case.
i. The [Mn2+] is doubled.
ii. The size of the Ti(s) electrode is doubled.
iii. The pH of both compartments is increased by the same amount.
8.
(12%) Three common allotropic forms of carbon are diamond, graphite, and buckminsterfullerene (C60).
a. Describe or sketch clearly the structure of each allotrope.
b. Compare diamond and graphite with regard to their hardness and electrical conductivity and account for any differences in
behavior on the basis of the structures in part a.
c. Graphite is more stable than diamond (by 2.9 kJ·mol–1) at room temperature and pressure. Explain why the diamonds in
jewelry do not change readily into graphite.
d. Use this phase diagram for carbon to determine which has the greater density,
106 diamond
diamond or graphite. Explain your reasoning and suggest a means of
liquid
converting graphite into diamond.
104
graphite
102
vapor
0
2000 4000 6000
Temperature, oC →
END OF PART II
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
atomic number
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
M
G molar
A frequency
mol
ν mole
Z gas constant
h
R Planck’s constant
N A gram
pressure
P
g
°C heat capacity
k
C p rate constant
c hour
Rf
h retention factor
C joule
s
J second
E kelvin
speed
of
light
c
K
Ea kilo- prefix
temperature,
K
T
k
H liter
t
L time
S
volt
V
E =E –
USEFUL EQUATIONS
 –∆H   1 
ln K = 
   +c
 R T
RT
lnQ
nF
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol –1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
1 atm = 760 mmHg
 k 2  Ea  1 1 
=
−
 k1  R  T1 T2 
ln 
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(289)
Page 4
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 21, 2002.
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART II
1.
a. i.
ii.
b.
c.
pH = 9 CO32– + H+ → HCO3–
pH = 5 HCO3– + H+ → H2CO3 or HCO3– + H+ → H2O + CO2
0.115 mol
9.63× 10 −3 L ×
= 1.107× 10−3 mol HCl to titrate CO32–
L
0.115 mol
3.427× 10 −3 L ×
= 3.941× 10 −3 mol HCl to titrate HCO3–
L
105.99 g Na 2 CO3
1.107 ×10 −3 mol CO3 2– ×
= 1.17× 10 –1 g Na2CO 3
mol
84.01 g NaHCO3
2.834 × 10−3 mol HCO3 – ×
= 2.38× 10 –1 g NaHCO3
mol
1.17 ×10 –1 g Na 2 CO3
× 100 = 23.4% Na 2 CO3
5.00 ×10 –1 g mixture
2.38× 10 –1 g NaHCO3
× 100 = 47.6% NaHCO 3
5.00 × 10–1 g mixture
d. The total volume required to reach the second
equivalence point is twice that required to reach the
first equivalence point because the number of moles
of HCO3– is equal to the number of moles of CO32–.
14
12
10
8
6
4
2
0
0
10.0
20.0
30.0
40.0
Volume HCl(aq), mL →
2.
a.i. ∆Hfo of N2O5(g)
o
∆Hrxn
= ∆H of (N 2 O5 ) − 2 ∆H of (NO 2 )
–55.1 kJ = ∆H of (N 2 O5 ) – 2 mol (33.2 kJ·mol–1)
∆H of (N 2 O5 ) = +11.3 kJ·mol–1
ii. S o of N2O5(g)
(
o
∆Srxn
= So(N 2O 5 ) – 2S o (NO2 ) + S o (O2 )
)
–227.0 J·K–1 = S o(N 2 O 5 ) – [2(–239.7 J·mol–1·K–1) + 1/2(205.1 J·mol–1·K–1)]
S o(N 2 O 5 ) = 355.4 J·mol–1·K–1
iii. ∆G o at 25 °C
∆Go = ∆Ho – T∆So
o
∆G = –55.1 kJ – (298 K)(–0.227 kJ·K–1)
∆Go = 12.5 kJ
iv. Kp at 25 °C
∆Go = – RT ln K p
12500 J =
–8.314 J
( 2 9 8 K ) lKn p
mol⋅ K
Key for 2002 USNCO National Exam, Part II
l nK p = –5.045 and K p = 6.44 ×10
–3
Page 1
b. i. This reaction is not spontaneous at 25 °C. The value of ∆G o is positive.
ii. An increase in temperature will cause ∆G o to become more positive because the value of ∆S o is negative.
Therefore, the numerical value of Kp will decrease.
iii. An increase in temperature will cause the relative amount of reactants to increase and products to
decrease. This can be explained by noting that the value of ∆Hrxn is negative, which means that adding
heat will shift the reaction to the left. Another argument is that as the temperature increases, the value of
the equilibrium constant Kp will decrease, also shifting the reaction to the left.
iv. The S o values for NO2(g) and O2(g) at the same temperature are not the same because NO2, with 3 atoms
per molecule, has more possible arrangements than O2, with only 2 atoms per molecule. This leads to a
higher value for entropy, although not very much higher. The molar mass of NO2 is also higher than
for O2.
3.
a.i. Ca3(PO4)2
ii. Ca(H2PO4)2
b.i. 1.0 × 10 –25
[ ] [PO ]
K = [Ca ] [ H PO ]
= [Ca ] [ PO ] Then, let 3x = [Ca ] and 2 x = [ PO ]
3
Ksp = Ca 2+
3– 2
4
2+
2
sp
2+ 3
4
– 2
3– 2
4
2+
3 –
4
1.0 × 10 –25 = [3 x ]3 [2x ]2
1.0 × 10 –25 = 108x 5 and x 5 = 9.26 ×10 –28 and x = 3.9 ×10 –6
[Ca ] = 3(3.92 ×10
2+
[
–6
) =1.2 ×10 –5 M
][
1.0 × 10 –3 = Ca2 + H 2 PO 4–
]2
[
]
[
Then, let x = Ca 2+ and 2x = H2 PO 4 –
]
1.0 × 10 –3 = [ x ] [2x ]2
1.0 × 10 –3 = 4x3 and x = 6.3 ×10 –2
[Ca ] = 6.3 ×10
2+
c.
–2
M
Ca(H2PO4)2(s) ¾ Ca2+(aq) + 2H2PO4–(aq)
[
][
1.0 × 10 –3 = Ca2 + H 2 PO 4–
[H 2PO 4– ] =
]2
1.0 ×10–3
= 6.67× 10–3 = 8.2× 10 –2 M
0.15
H3PO4(aq) ¾ H+(aq) + H2PO4–(aq)
[ H ][ H PO ]
+
7.3 × 10 –3 =
7.3 × 10 –3 =
2
[ H3 PO 4]
4
–
[ x] [0.082 + x ] and x = 1.7 ×10–2 M
[0.25 − x ]
[ ]
Therefore, to prevent precipitation, H + must be greater than = 1.7 × 10 –2 M
4.
a. i. First order in I–. Compare the results of experiments 1 and 2 to see that the rate went up by a factor of 4
when the concentration of I– went up by 4.
ii. Second order in IO3–. Compare the results of experiments 1 and 3 to see that the rate went up by 9
when the concentration of IO3– went up by 3.
Page 2
Key for 2002 USNCO National Exam, Part II
iii. Second order in H+. Compare the results of experiments 1 and 4 to see that the rate went up by 4 when
the concentration of H+ went up by 2.
[ ] [IO ] [H ]
b. rate = k I –
1
+ 2
– 2
3
1
2
2
rate = k [0.01] [0.10 ] [ 0.01]
0.60
k=
= 6.0× 10 7 mol–4 ⋅ L4 ⋅ s–1
1.0 ×10 –8
c. Reaction is very unlikely to occur in one step. That would require the simultaneous collision of five
particles.
d. Pseudo-first order refers to carrying out a reaction under conditions such that only one reactant changes
concentration. For this reaction, pseudo-first order kinetics can be established by having a large excess of
[IO3–] and [H+].
–E a 2
Ea1
 Ea2

RT
k
Ae
k2
–
RT –
RT 
e. 2 =
and
after
cancelling
A
and
combining
exponents,
=
e
– Ea1
k1
k1
RT
Ae
Ea
Ea
k
Taking the natural log and rearranging yields this expression.
ln 2 = 1 – 2
k1 RT RT
ln
k2
8.4 ×10 4 J ⋅ mol–1
7.4 × 104 J ⋅mol –1
=
–
–1
–1
k1 (8.314 J⋅ mol ⋅ K ) ( 2 9 8 K ) (8.314 J ⋅ mol –1 ⋅ K –1 ) ( 2 9 8 K )
ln
k
k2
= 3 3 . 9 0 –29.87 = 4 . 0 3 a n d 2 = e4.03
k1
k1
k2
= 56.3
k1
5. Note: Balanced equations were not required.
a. H+ + Cl– + MnO2 → Cl2 + Mn2+ + H2O
b. Mg2+ + SO42– + Ba2+ + OH– → BaSO4 + Mg(OH)2
c. BaO2 + H2O → Ba2+ + OH– + O2H–
d. Cu + H+ + NO3– → Cu2+ + NO + H2O
e. AgBr + S2O32– → Ag(S2O3)23– + Br–
f.
H H OH H
H H O H
H C C C C H + Cr2O72– + H+ → H C C C C H + Cr3+ + H2O
H H H H
H H
H
F
6. a.
F
Cl F
b. T-shaped F
F
Cl
F
F
c. F Cl
F
or F Cl
F
F
These structures are not favored because they provide less volume for the lone pairs and therefore do not
minimize all repulsions.
Key for 2002 USNCO National Exam, Part II
Page 3
+
F
d. sp3d
Cl
e. bent
–
F
square planar
F
7.
F
Cl
F
F
a. Ti + 2MnO2 + 6H+ → 2Mn2+ + 3H2O + TiO2+
Eo = 1.23 V + 0.88 V = 2.11 V
b. MnO2 + 4H+ + 2e– → Mn2+ + 2H2O
Reduction occurs at the cathode.
c. The conditions for standard potential are 25°C, 1 atm pressure, and 1 M concentrations.
d. i. Doubling [Mn2+] will decrease the potential because Mn2+ is a product. An increase in Mn2+ will shift the
equilibrium to the left.
ii. Doubling the size of the electrode has no effect. The electrode does not appear in the equilibrium
expression nor in the Nernst equation.
iii. Increasing pH lowers [H+]. Because H+ appears on the left of the balanced equation, decreasing [H+]
will lower the potential. The reaction shifts to the left.
8.
a.
diamond
sp3 hybridization
3-dimensional tetrahedral
network solid
graphite
sp2 hybridization
2-dimensional sheets
trigonal planar covalent
half-filled p orbital
hexagonal rings
C60
spherical shape made up
of hexagonal and pentagonal rings
“soccer ball” design
C
C
C
C
C
C
C
C
C
C
b. Diamond is the hardest. Diamond has 4 covalent bonds per C atom, making a very strongly bonded 3-D
network solid. Graphite’s sheets have only weak forces between the sheets, allowing one to slide by the other.
This makes graphite much “softer” than diamond.
All valence electrons in diamond are involved in sigma bonds, resulting in a nonconducting material.
Graphite has delocalized electrons in the half-filled p orbitals (pi-bonding), allowing for electron movement
from one atom to the next when an electromotive force is applied. Graphite is a good conductor.
c. Although equilibrium favors graphite at room temperature, the rate of the reaction from diamond to
graphite is extremely slow because of a very high activation energy barrier.
d. The graph shows that as pressure is increased at a fixed temperature of 1000 °C, graphite is converted into
diamond. Since increasing pressure should favor increasing density, one could conclude that diamond is denser
than graphite. Since graphite is composed of sheets with considerable empty space between the carbon layers,
converting graphite to the tetrahedral form decreases the empty space and increases density.
To prepare diamond from graphite, the graph indicates that by carrying out the process at 0 °C, a pressure
of only 104 atm would be needed. However, since the rate of change would be very slow, this might not be the
most ideal set of conditions. An alternate method might be to heat graphite to 2500 °C and apply a pressure of
105 atm, which should increase the rate of the conversion. Another alternate method would be to apply a
pressure of 103 atm, then heat to 5000 °C to allow for liquefication, increase pressure to 106 atm and then cool
to the lower temperature.
Page 4
Key for 2002 USNCO National Exam, Part II
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Lucy Pryde Eubanks, Clemson University, Clemson, SC
Chair
Robert Becker, Kirkwood High School, Kirkwood, MO
Nancy Devino, Coordinating Board for Higher Education, Jefferson City, MO
Sheldon L. Knoespel, Michigan State University, East Lansing, MI
Steve Lantos, Brookline High School, Brookline, MA
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Robert G. Silberman, SUNY-Cortland, NY
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 1 hour, 30 minutes allotted to this part of the test. Students do not need to
stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by
the examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 1 hour, 30 minutes. There is no need to stop between tasks
or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are
required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that
problem. You are permitted to use a non-programmable calculator. At the end of 1 hour, 30 minutes, all answer sheets should be
turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions
from your examiner for safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 1
2002 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw
conclusions from your experimental work. You will be graded on your experimental design, on your skills in
data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are
also components of successful solutions to these problems, so make your written responses as clear and
concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
Design and carry out an experiment to investigate a relationship between the surface area of a piece of raw
potato and the rate of decomposition of hydrogen peroxide. You may use only those materials available at
your experimental station. You will be asked to describe the method you developed to carry out this
investigation.
Lab Problem 2
Design and carry out an experiment to determine the equilibrium constant, Keq, for this reaction at room
temperature.
urea(s) + H2O(l) ¾ urea(aq)
You will be asked to describe the method you developed to solve this problem.
Given:
Page 2
molar mass of urea, CO(NH2)2
molarity of pure H2O
= 60.0 g·mol–1
= 55.5 mol·L–1
Not valid for use as an USNCO National Examination after April 21, 2002.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School: ________________________________________Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name : _______________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 3
2. Record your data and other observations.
3. What relationship did you have find between the surface area of a raw potato and the rate of decomposition
of hydrogen peroxide? Support your conclusion with your experimental evidence.
Page 4
Not valid for use as an USNCO National Examination after April 21, 2002.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School: ________________________________________Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name : _______________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
Not valid for use as an USNCO National Examination after April 21, 2002.
Page 5
2. Record your data and other observations.
3. What value did you calculate for the equilibrium constant? Show your methods clearly.
Page 6
Not valid for use as an USNCO National Examination after April 21, 2002.
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Directions
Thank you for administering the 2002 USNCO laboratory practical on behalf of your Local
Section. It is essential that you follow the instructions provided, in order to insure consistency of results
nationwide. There may be considerable temptation to assist the students after they begin the lab exercise.
It is extremely important that you do not lend any assistance or provide any hints whatsoever to the
students once they begin work. As is the case with the international competition, students should not be
allowed to speak to anyone until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab
station or table when the students enter the room. The equipment should be initially placed so that the
materials used for Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. Chemistry Olympiad
National Examination. In this part of the exam, we will be assessing your lab skills and your ability to
reason through a laboratory problem and communicate your results. Do not touch any of the equipment
in front of you until you are instructed to do so.
Both of this year’s problems use some small-scale chemistry equipment. Small-scale chemistry
techniques help to minimize the amount of materials you use, thereby increasing safety and minimizing
waste. Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets.
Show the Beral-type pipets being used at your site. If you substituted droppers in the first
problem, show those as well.
You may be unfamiliar with the graduated centrifuge tubes available in both parts of this lab practical
exam. This is the type we will be using.
Show the type of centrifuge tube being used at your site.
The watertight caps will be an advantage in at least one of the problems.
You will be asked to complete two laboratory problems. The materials and equipment needed to solve
each problem has been set out for you and is grouped by the number of the problem. You also may use
distilled (or deionized) water. You must limit yourself to this equipment and materials for each problem.
A balance is not needed for either problem. You may choose to start with either problem. You are
required to have a procedure for each problem approved for safety by an examiner. (Remember that
approval does not mean that your procedure will be successful–it is a safety approval.) When you are
ready for an examiner to come to your station for each safety approval, please raise your hand.
You will have one hour and thirty minutes to complete both problems.
Examiner’s Directions, 2002 USNCO National Exam, Part III
Page 1
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash
off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate
procedures for disposing of solutions at the end of this lab practical are:
____________________________________________________________________________________
____________________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. Are there any questions before we begin?
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. Allow students
enough time to read the brief cover directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out
all information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for
your cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the
lab practical with the students. They can learn about possible observations and interpretations and you
can acquire feedback as to what they actually did and how they reacted to the problems. After this
discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be
extremely useful to the Olympiad Laboratory Practical subcommittee as they prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron
sheets from Part I, and the “Blue Books” from Part II to this address:
ACS DivCHED Exams Institute
Clemson University
223 Brackett Hall
Clemson, SC 29634-0979
Wednesday, April 24, 2002 is the absolute deadline for receipt of the exam materials at the
Examinations Institute. Materials received after this deadline CANNOT be graded.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT
SCHEDULE FOR GRADING THIS EXAMINATION.
Page 2
Examiner’s Directions, 2002 USNCO National Exam, Part III
EXAMINER’S NOTES
Lab Problem #1: Materials and Equipment. Each student will need:
1 stopwatch, timer, or access to clock with second hand
6 25-mL or 15-mL graduated cylinders, with bases
Note: If providing this many graduated cylinders per student is not possible,
6 13 x 100 test tubes and a test tube rack may be substituted.
2 small beakers (100 mL or 250 mL); one labeled “water”, one labeled
“3% hydrogen peroxide”
2 1-mL Beral-style pipets (eye droppers may be substituted)
2 15-mL graduated centrifuge tubes (see lab problem #2 for specifications; caps
not needed for this lab problem)
1 6-in plastic ruler
1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water”
1 25-mL dropping bottle labeled “liquid detergent”
1 kitchen cutting board (or suitable clean hard surface on lab bench)
1 sharp kitchen paring knife, non-serrated edge
1 plastic container (such as a margarine tub or a Deli salad container); capable of holding approximately
200 mL
1 plastic tub for disposal of liquid wastes (or easy access to sinks)
supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
Lab Problem #1: Chemicals . Each student will need:
1 8-fluid oz (237 mL) bottle of 3% hydrogen peroxide
Note: Hydrogen peroxide antiseptic is sold in the first aid section in most supermarkets and drug stores.
Provide each student with an unopened bottle to emphasize the use of a consumer product. The cheapest
brand, so long as it is fresh and unopened, will work.
1 white potato (Russet potatoes work well and are generally available; provide potato whole and unpeeled)
10 mL liquid detergent
100 mL of distilled or deionized water
Quick Check to be sure this lab problem will work for your examinees:
1) Are the bottles of hydrogen peroxide fresh and unopened?
2) Is the knife capable of making clean cuts in the potato?
3) Have all detergent or soap residues been removed from the glassware?
Lab Problem #1: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment
on the plan other than looking for any potentially unsafe practices.
2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical.
A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit
directions for handling spills and for disposing of waste materials, following approved safety practices for your
examining site. Please check and follow procedures appropriate for your site.
Examiner’s Directions, 2002 USNCO National Exam, Part III
Page 3
Lab Problem #2: Materials and Equipment. Each student will need:
4 15-mL centrifuge tubes with 0.1 mL or 0.5 mL graduations.
The centrifuge tubes should have conical bottoms and screw
caps. 2 tubes labeled “~9.0 mL”; 2 tubes labeled “4.0 g urea”
Note: Polystyrene, polypropylene, or Pyrex® centrifuge
tubes are widely used in biology, chemistry, and
biochemistry departments. The dome-seal screw cap prevents
loss of any liquid, important for this experiment. The
conical-bottom tubes are preferred to the round bottom tubes.
It is not necessary to use the far more expensive borosilicate
centrifuge tubes.
1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water”
1 10-mL graduated Beral-style pipet
1 small beaker (100 mL or 250 mL), labeled “water”
1 plastic tub for disposal of liquid wastes (or easy access to sinks)
supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
Lab Problem #2: Chemicals. Each student will need:
2 4.0 g samples of urea, (NH2)2CO, provided in closed, labeled conical-bottom graduated centrifuge tubes.
The samples should be prepared in advance.
100 mL of distilled or deionized water
Quick Check to be sure this lab problem will work for your examinees:
1) Have the correct centrifuge tubes been obtained and labeled?
2) Have two 4.0 g urea samples been prepared and placed in the labeled centrifuge tubes for each student?
3) Have two other centrifuge tubes been labeled “~9.0 mL H2O” for each student? These should not be
filled in advance.
Lab Problem #2: Notes
1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment
on the plan other than looking for any potentially unsafe practices.
2. Be sure that the labels on the centrifuge tubes do not obscure any graduations.
3. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical.
A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit
directions for handling spills and for disposing of waste materials, following approved safety practices for your
examining site. Please check and follow procedures appropriate for your site.
Page 4
Examiner’s Directions, 2002 USNCO National Exam, Part III
2002 U. S. NATIONAL
CHEMISTRY OLYMPIAD
KEY for NATIONAL EXAM—PART III
Lab Problem 1
Part 1. Experimental Plan
A good plan included a detailed description of a method to observe the rate of reaction. It would also include a plan
for varying the surface area of the potato. Finally, the plan needed to account for the importance of controlling the
volume of peroxide.
For example, a good plan might consist of these steps.
1) Cut the potato into different size cubes, such as 2 cubes of 0.5 cm on a side and 2 cubes of 1.0 cm on a side.
2) Measure 2.0 mL of H2O2 into a test tube or graduated cylinder and add 3 drops of detergent..
3) Shake to generate a small amount of foam for a starting point. Measure the height of the foam column or read the
volume of the foam directly if using a graduated cylinder.
4) Drop in one cube of potato and start a timer.
5) At appropriate intervals, measure the height of the foam column or read the volume of the foam directly if using a
graduated cylinder.
6) Repeat steps 2-5 for other potato pieces.
An average plan was either missing one of these three components or had less detail in two or more of these
components.
A weak plan had minimal detail about how the experiment would be conducted.
Part 2. Experiments and Observations
A good experimental section included these points.
1) Appropriate measurements of
a. reaction rate or progress.
b. dimensions of the potato pieces used.
c. time.
d.volume of H2O2 used.
2) Multiple (at least two) trials for each different surface area of the potato piece.
3) Appropriate quantitative detail such as
a. precision in the trials.
b.averaging trial data.
c. description of any calculation methods used, such as for determining the surface area of the potato piece.
An average experimental section was either missing one of these three components or had less detail in two or
more of these components.
A weak experimental section had minimal detail about how the experiment was conducted and what observations
were made.
Part 3. Discussion
A good discussion included these points.
1) Calculations or graphical determinations of the relationship between surface area and rate of decomposition of
hydrogen peroxide.
2) An appropriate description of the scientific reasoning utilized.
3) An accurate conclusion supported by the experimental observations and data reported.
Note: Points were not deducted for correct and reasoned discussion of an experiment with seemingly anomolous
results.
Page 1
2002 USNCO National Exam, Part III (Lab Practical)
Lab Problem 2
Part 1. Experimental Plan
A good plan recognized that it was necessary to determine how much water was required to completely dissolve
4.0 g of urea.
For example, a good plan might consist of these steps.
1) Add water in small increments to 4.0 g of urea in the graduated centrifuge tube.
2) Cap the tube and shake after each addition.
3) If any solid remains, add another small portion of water. Cap and shake the tube.
4) Continue to add water until all the urea is dissolved.
5) Record the total volume of solution and/or the total volume of water added.*
6) Repeat with the second 4.0 g sample of urea.**
An average plan was either missing one of these components or had less detail in two or more of these components.
A weak plan had minimal detail about how the experiment would be conducted.
Part 2. Experiments and Observations
Sample Data
Trial 1
Total Volume of
Solution*
7.3 mL
Mass Urea**
4.0 g
Trial 2
7.5 mL
4.0 g
Many students also observed that as the urea dissolved, the tube felt cool to the touch. Some allowed time for
the tube to return to room temperature before making final observations of volume.
Part 3. Calculation of Equilibrium Constant
Sample calculations for Trial 1
1 mol urea
1) Moles of urea = 4.0 g urea ×
= 0.067 mol urea
60.0 g urea
0.067 mol urea
= 9.2 M
0.0073 L solution
3) Calculation of Keq if assume that the concentration of water is in standard state.
2) Molarity of urea solution =
urea(s) + H2O(l) ¾ urea(aq)
Keq = [urea (aq) ]
Keq = 9.2
* A superior plan recognized that because there is a high concentration of urea in the saturated solution, the solution cannot
be treated as a dilute solution in which the concentration of the solvent is the same as that of pure water. Points were awarded
to students who realized this and adjusted their experimental approach. This requires knowing the volume of water added, not
just the volume of the resulting solution. For example, if 4.3 mL of water was added, the final volume of the solution was
reported as 7.3 mL. The molarity of water in the saturated solution can then be calculated, as shown in this example.
4.3 mL H2 O
(9.2 M)
[urea (aq)]
and Keq =
× 5 5 M= 32 M and Keq =
= 16
7.3 mL solution
(1)(32 M /55.5 M)
[ urea (s)] [H 2O(l )]
** Some students elected to obtain more samples by dividing the given mass of urea. The mass of smaller samples of urea
was estimated by calculating the density of urea, given the known mass and the volume markings on the graduated centrifuge
tube.
Page 2
2002 USNCO National Exam, Part III (Lab Practical)
2003 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
David W. Hostage, Taft School, CT
Alice Johnsen, Bellaire High School, TX
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
Amy Rogers, College of Charleston, SC
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 27, 2003, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 27, 2003.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
ABBREVIATIONS AND SYMBOLS
n Faraday constant
F molal
A formula molar mass
M molar
atm free energy
G molar mass
u frequency
ν mole
A gas constant
R Planck’s constant
N A gram
g pressure
°C heat capacity
C p rate constant
c hour
h retention factor
C joule
J second
E kelvin
K speed of light
Ea kilo– prefix
k temperature, K
H liter
L time
S
measure of pressure mmHg volt
K
milli– prefix
m
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
c
T
t
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
EQUATIONS
E = Eo −
1
1A
1
H
k  E  1 1 
ln 2  = a  − 
 k1  R  T1 T2 
 −∆H  1 
lnK = 
  + c
 R  T 
RT
ln Q
nF
PERIODIC TABLE OF THE ELEMENTS
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
1.008
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Not valid for use as an USNCO National Examination after April 27, 2003
DIRECTIONS
! When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
! There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
1. In an experiment to determine the percentage of water in
a solid hydrate by heating, what is the best indication that
all the water has been removed?
(A) The solid melts.
(B) The solid changes color.
(C) Water vapor no longer appears.
6. According to the
solubility curve
shown, how many
grams of solute can be
recrystallized when 20
mL of a saturated
solution at 60 ˚C are
cooled to 0 ˚C?
(D) Successive weighings give the same mass.
60
40
20
10
K
o
Temperature, C
2. The curve shown results
when a liquid is cooled.
What temperature is closest
to the freezing point of the
liquid?
Solubility (g solute / 100 mL soln)
! Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
(A) 7.0
M
L
N
(B) 12
30
50
Temperature ( o C)
(C) 25
70
(D) 35
7. Which would produce the largest change in the H2O level
when added to water in a 25 mL graduated cylinder?
(A) 10.0 g of Hg (d = 13.6 g·mL-1)
Time, min
(A) L
(B) M
(C) L + M
2
(D) M + N
2
3. What is the proper way to dispose of a two milliliter
sample of hexane after completing experiments with it?
(A) Return it to the solvent bottle.
(B) Place it in a waste bottle with compatible organic
materials.
(C) Flush it down the drain with large quantities of
water.
(D) Pour it on a solid absorbent so it can be thrown away
with solid waste.
4. Which anion can undergo both oxidation and reduction?
(A) Cr2O72(B) NO3(C) OCl -
(D) S 2-
5. The mass percentages in a compound are carbon 57.48%,
hydrogen 4.22% and oxygen 38.29%. What is its
empirical formula?
(A) C 2H2O
(B) C 4H3O2
(C) C 5H4O2
(D) C 8H7O4
(B) 7.42 g of Al (d = 2.70 g·mL-1)
(C) 5.09 g of iron pyrite (d = 4.9 g·mL-1)
(D) 2.68 g of oak (d = 0.72 g·mL -1)
8. Diborane, B2H6, can be prepared by the reaction;
3NaBH 4 + 4BF3 r 3NaBF4 + 2B2H6
If this reaction has a 70 percent yield, how many moles
of NaBH4 should be used with excess BF3 in order to
obtain 0.200 mol of B2H6?
(A) 0.200 mol
(B) 0.210 mol
(C) 0.300 mol
(D) 0.429 mol
9. What volume of 6.0 M H2SO4 should be mixed with
10. L of 1.0 M H2SO4 to make 20. L of 3.0 M H2SO4
upon dilution to volume?
(A) 1.7 L
(B) 5.0 L
(C) 8.3 L
(D) 10. L
10. An aqueous solution that is 30.0% NaOH by mass has a
density of 1.33 g.mL-1. What is the molarity of NaOH in
this solution?
(A) 8.25
Not valid for use as an USNCO National Examination after April 27, 2003
(B) 9.98
(C) 16.0
(D) 33.2
Page 3
(A) 1 only
(B) 2 only
(C) both 1 and 2
(D) neither 1 nor 2
12. Benzene melts at 5.50 ˚C and has a freezing point
depression constant of 5.10 ˚C. m-1. Calculate the freezing
point of a solution that contains 0.0500 mole of acetic
acid, CH 3COOH, in 125 g of benzene if acetic acid forms
a dimer in this solvent.
(A) 3.46 ˚C
(B) 4.48 ˚C
(C) 5.24 ˚C
(D) 6.01 ˚C
(B) C 4H10
(C) C 5H12
Temperature
(A) decrease both the melting and boiling points
(B) increase both the melting and boiling points
(C) increase the melting point and decrease the boiling
point
(D) decrease the melting point and increase the boiling
point
13. A 200. mL sample of a gaseous hydrocarbon has a
density of 2.53 g.L-1 at 55 ˚C and 720 mmHg. What is its
formula?
(A) C 2H6
18. According to the phase
diagram, what would be the
effect of increasing the
pressure on this substance?
Pressure
11. Which change
1. an increase in water temperature
increases the
solubility of a gas in 2. a decrease in gas pressure
water?
(D) C 6H6
19. When the substances below are arranged in order of
increasing entropy values, S˚, at 25 ˚C which is the
correct order?
(A) CO2(s) < CO2(aq) < CO 2(g)
14. A liquid has a vapor pressure of 40 mmHg at 19.0 ˚C and
a normal boiling point of 78.3 ˚C. What is its enthalpy of
vaporization in kJ . mol -1?
(A) 42.4
(B) 18.4
(C) 5.10
15. Sulfur and fluorine form SF 6 and
S 2F 10, both of which are gases at
30 ˚C. When an equimolar
mixture of them is allowed to
effuse through a pinhole, what is
(D) 1.45
Molar Mass g.mol-1
SF 6
S 2F 10
146
254
the ratio SF 6/S2F10 in the first sample that escapes?
(A) 1.32/1
(B) 1.74/1
(C) 3.03/1
(C) CO2(s) < CO2(g) < CO2(aq)
(D) CO2(g) < CO2(s) < CO2(aq)
20. When 50. mL of 0.10 M HCl is mixed with 50. mL of
0.10 M NaOH the temperature of the solution increases
by 3.0 ˚C. Calculate the ∆Hneutralization per mole of HCl.
(The solution has a density = 1.0 g.mL-1 and
C p = 4.2 J . g-1. ˚C-1)
(A) 1.3 × 103 kJ
(B) -1.3 × 102 kJ
(C) -2.5 × 102 kJ
(D) -1.3 × 103 kJ
(D) 3.48/1
16. The volumes of real gases often exceed those calculated
by the ideal gas equation. These deviations are best
attributed to the
(A) attractive forces between the molecules in real
gases.
(B) dissociation of the molecules in real gases.
(C) kinetic energy of the molecules in real gases.
(D) volumes of the molecules in real gases.
17. The electrical conductivity of a solid is slight at 25 ˚C
and much greater at 125 ˚C. The solid is most likely a(n)
(A) ionic compound
(B) insulator
(C) metal
(D) semiconductor
Page 4
(B) CO2(g) < CO2(aq) < CO 2(s)
21. The combustion of 0.200 mol of
liquid carbon disulfide, CS2, to
give CO 2(g) and SO2(g) releases
215 kJ of heat. What is ∆Hf˚ for
CS2(l) in kJ. mol -1?
(A) 385
(B) 87.9
∆ Hf˚
kJ.mol-1
CO2(g)
SO2(g)
-393.5
-296.8
(C) -385
(D) -475
22. For the reaction: 2NO2(g) r N2O4(g) ∆H < 0. What
predictions can be made about the sign of ∆S and the
temperature conditions under which the reaction would
be spontaneous?
∆ Srxn
Temperature Condition
(A) negative
low temperatures
(B) negative
high temperatures
(C)
positive
high temperatures
(D)
positive
low temperatures
Not valid for use as an USNCO National Examination after April 27, 2003
23. As ∆G˚ for a reaction changes from a large negative
value to a large positive value, K for the reaction will
change from
(A) a large positive value to a large negative value.
(B) a large positive value to a small positive value.
29. For the reaction
2A + 2B r Product
the rate law is Rate = k[A][B]2. Which mechanism is
consistent with this information?
(A) B + B s C
(C) a large negative value to a large positive value.
(D) a large negative value to a small negative value.
C + A r Product (slow)
(C) A + A s C
24. ∆E˚ is measured at constant volume and ∆H˚ is measured
at constant pressure. For the reaction;
2C (s) + O2(g) r 2CO(g) ∆H˚ < 0 kJ
How do the ∆E˚ and ∆H˚ compare for this reaction?
(A) ∆E˚ < ∆H˚
(B) ∆E˚ > ∆H˚
(C) ∆E˚ = ∆H˚
(D) Impossible to tell
from this information.
25. Which statement about second order reactions is correct?
(A) Second order reactions require different reactants.
(B) Second order reactions are faster than first order
reactions.
(C) Second order reactions are unaffected by changes in
temperature.
(D) The half-life of a second order reaction depends on
the initial reactant concentration.
26. A first order reaction has a rate constant of 0.0541 s-1 at
25 ˚C. What is the half-life for this reaction?
(A) 18.5 s
(B) 12.8 s
(C) 0.0781 s
(D) 0.0375 s
27. The reaction between NO and I2 is second-order in NO
and first-order in I 2. What change occurs in the rate of the
reaction if the concentration of each reactant is tripled?
(A) 3-fold increase
(B) 6-fold increase
(C) 18-fold increase
(D) 27-fold increase
28. For the rate equation,
Rate = k[A][B]2,
what are the units for the rate constant, k, if the rate is
given in mol . L-1. sec-1?
(A) L. mol . sec
(C) L2. mol -2. sec-1
(B) L. mol -1. sec-1
(D) L3. mol -3. sec-2
(B) A + B r C (slow)
C + B r product
(D) A + B s C
B+BsD
B + C r D (slow)
C + D r Product (slow)
D + A r product
30. Which straight line gives the activation energy for a
reaction?
(A) rate constant vs T
-1
(C) rate constant vs T
(B) ln (rate constant) vs T
(D) ln (rate constant) vs T-1
31. Based on the equilibrium constant for the reaction below,
2SO3(g) s 2SO2(g) + O2(g)
K = 1.8 × 10-5
what is the equilibrium constant for the reaction
SO2 (g) + 1/2O2 (g) s SO3(g)
K=?
(A) 2.1 × 10-3
(B) 4.2 × 10-3
(C) 2.4 × 102
(D) 5.6 × 104
32. CO(g) + Cl2(g) s COCl(g) + Cl(g) Keq = 1.5 × 10-39
If the rate constant, k, for the forward reaction above is
1.4 x 10 -28 L . mol-1. sec-1 what is k (in L. mol-1. sec-1 ) for
the backward reaction?
(A) 2.1 × 10-67
(B) 1.0 × 10-11
(C) 9.3 × 1010
(D) 7.1 × 1027
33. Calculate the [H +] in a 0.25 M solution of methylamine,
CH3NH2 (Kb = 4.4 × 10-4).
(A) 1.1 × 10-4
(B) 1.0 × 10-2
(C) 9.1 × 10-11
(D) 9.5 × 10-13
34. A 0.010 M solution of a weak acid, HA, is 0.40%
ionized. What is its ionization constant?
(A) 1.6 × 10-10
(B) 1.6 × 10-7
(C) 4.0 × 10-5
(D) 4.0 × 10-3
Not valid for use as an USNCO National Examination after April 27, 2003
Page 5
35. 1.0 L of an aqueous solution in which
[H 2CO3] = [HCO3-] = 0.10 M, has [H+] = 4.2 × 10-7. What
is the [H+] after 0.005 mole of NaOH has been added?
(A) 2.1 ×10-9 M
(B) 2.2 × 10-8 M
(A) 0.629 V
(B) 0.689 V
(C) 3.8 × 10 M
(D) 4.6 × 10 M
(C) 0.748 V
(D) 0.866 V
-7
-7
36. A solution of Pb(NO3)2 is added dropwise to a second
solution in which [Cl-] = [F-] = [I -] = [SO42-] = 0.001 M.
What is the first precipitate that forms?
37.
41. The voltage for the cell
Fe ❘ Fe2+(0.0010 M) ❘ ❘ Cu2+(0.10 M) ❘ Cu
is 0.807 V at 25 ˚C. What is the value of E˚?
(A) PbCl2
(Ksp = 1.5 × 10-5)
(B) PbF2
(K sp = 3.7 × 10-8)
(C) PbI2
(K sp = 8.5 × 10-9)
(D) PbSO4
(K sp = 1.8 × 10-8)
(A) 0.39 M
(B) 3
(C) 4
(D) 6
38. Use the standard reduction potentials;
Sn 2+(aq) + 2e– r Sn(s) E˚ = -0.141 V
Ag+(aq) + e– r Ag(s)
E˚ = 0.800 V
to calculate E˚ for the reaction;
Sn(s) + 2Ag +(aq) r Sn2+(aq) + 2Ag(s)
(B) 0.941 V
(C) 1.459 V
(D) 1.741 V
39. Which of the
processes happen
during the
discharging of a lead
storage battery?
(B) microwave
(C) ultraviolet
(D) visible
44. All of the following possess complete d shells EXCEPT
(B) 2 only
(C) 1 and 3 only
(D) 2 and 3 only
(B) -585 kJ
(C) -390 kJ
(D) -195 kJ
Page 6
(B) Cu 2+
(C) Ga 3+
(D) Zn2+
(A) 6s
(B) 5p
(C) 5d
(D) 4d
46. Which set of quantum numbers (n, l, ml , ms) is
permissible for an electron in an atom?
(A) 1 only
(A) -1170 kJ
(D) 0.89 M
45. Which orbital fills completely immediately before the 4f?
1. H2(g) is produced
2. PbO2 is converted to PbSO4
3. The density of the electrolyte
solution decreases
40. What is the value of ∆G˚ for the reaction?
2Al(s) + 3Cu 2+(aq)
r 2Al3+(aq) + 3Cu(s)
(C) 0.78 M
(A) infrared
(A) Ag+
(A) 0.659 V
(B) 0.46 M
43. Which region of the electromagnetic spectrum is capable
of inducing electron transitions with the greatest energy?
Cl2 + OH- r Cl- + ClO3What is the coefficient for OH- when this equation is
balanced with the smallest integer coefficients?
(A) 2
42. A current of 2.0 A is used to plate Ni(s) from 500 mL of
a 1.0 M Ni2+(aq) solution. What is the [Ni2+] after 3.0
hours?
(A) 1, 0, 0, -1/2
(B) 1, 1, 0, +1/2
(C) 2, 1, 2, +1/2
(D) 3, 2, -2, 0
47. When the elements Li, Be, and B, are arranged in order
of increasing ionization energy, which is the correct
order?
(A) Li, B, Be
(B) B, Be, Li
(C) Be, Li, B
(D) Li, Be, B
48. Which forms the most alkaline solution when added to
water?
(A) Al2O3
(B) B 2O3
(C) CO2
(D) SiO 2
E˚ = 2.02 V
49. What is the total number of valence electrons in the
peroxydisulfate, S2O82-, ion?
(A) 58
(B) 60
(C) 62
(D) 64
Not valid for use as an USNCO National Examination after April 27, 2003
50. For which species are both
bonds of equal length?
1. ClO22. NO2-
(A) 1 only
(B) 2 only
(C) both 1 and 2
(D) neither 1 nor 2
51. Which compound has the highest melting point?
(A) MgO
(B) KCl
(C) NaCl
(D) CaO
52. Which molecular geometry is least likely to result from a
trigonal bipyramidal electron geometry?
57. Which substance will react most rapidly with Br 2(aq)?
(A) benzene
(B) chloropropane
(C) propanone
(D) propene
58. Which compound includes a carbon atom with an sp
hybridized orbital?
(A) benzene
(B) butyne
(C) methyl chloride
(D) phenol
59. Which compound has the highest vapor pressure
at 25 ˚C?
(A) trigonal planar
(B) see-saw
(A) CH3CH2CH2CH2OH
(B) CH3CH2CH2OCH3
(C) linear
(D) t-shaped
(C) CH3CH2CH2CH2NH2
(D) (CH3) 3COH
53. Which diatomic species has the greatest bond strength?
(A) NO
(B) NO+
(D) O2-
(C) O2
60. Which of the molecules can exist as optical isomers?
54. During the complete combustion of methane, CH4, what
change in hybridization does the carbon atom undergo?
(A) sp3 to sp
(B) sp3 to sp 2
(C) sp2 to sp
(D) sp2 to sp 3
O
(A)
C
Br
H
O
OH
C
H
(C)
55. What is the formula for the
compound?
C
Br
OH
C
H
O
(B)
CH3
O
(D)
CH3
CH3
(A) C 8H10
(B) C 8H12
(C) C 8H14
(D) C 8H16
H
OCH3
C
OH
C
Br
C
Br
C
Br
H
H
56. Which is most likely to react by an SN1 mechanism?
(A) CH3Cl
(B) CH3CHClCH3
(C) (CH3) 3CCl
(D) C 6H5Cl
Not valid for use as an USNCO National Examination after April 27, 2003
END OF TEST
Page 7
CHEMISTRY OLYMPIAD 2003
NATIONAL EXAM
PART 1— KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
D
B
B
C
D
A
B
D
C
B
D
B
C
A
A
D
D
B
A
C
B
A
B
A
D
B
D
C
D
D
Not valid for use as an USNCO National Examination after April 27, 2003
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
C
C
D
B
C
D
D
B
D
A
C
C
C
B
A
A
A
A
C
C
A
A
B
A
A
C
D
B
B
B
2003 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Peter E. Demmin (retired), Amherst Central High School, NY
David W. Hostage, Taft School, CT
Alice Johnsen, Bellaire High School, TX
Jerry D. Mullins, Plano Senior High School, TX
Ronald O. Ragsdale, University of Utah, UT
Amy Rogers, College of Charleston, SC
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a
score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2003, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front
of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper,
and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 27, 2003.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
1.
(12%) 0.1152 g of a compound containing carbon, hydrogen, nitrogen and oxygen are burned in excess oxygen. The gases
produced are treated further to convert nitrogen-containing products into N2. The resulting mixture of CO2, H2O and N2 and
excess O2 is passed through a CaCl2 drying tube, which gains 0.09912 g. The gas stream is bubbled through water where the CO 2
forms H2CO3. Titration of this solution to the second endpoint with 0.3283 M NaOH requires 28.81 mL. The excess O2 is
removed by reaction with copper metal and the N2 is collected in a 225.0 mL measuring bulb where it exerts a pressure of 65.12
mmHg at 25 ˚C. In a separate experiment the molar mass of this compound is found to be approximately 150 g. mol -1.
a. Calculate the number of moles of
i.
H2O
ii.
CO2
iii.
N2
b. Determine the mass in the original compound of
i.
C
ii.
H
iii.
N
iv.
O
c. Find the empirical formula of the compound.
d. Find the molecular formula.
2.
(13%) The enthalpy of combustion of liquid octane, C8H18(l) to gaseous products, is -5090 kJ.mol -1. Use this value to answer the
questions below, assuming a temperature of 100 ˚C.
a. Write a balanced equation for the complete combustion of liquid octane C 8H18(l).
b. Determine the molar enthalpy of formation, ∆Hf˚, for liquid octane, C8H18(l).
[∆H f˚ kJ. mol -1; CO2(g) -393.5, H2O(g) -241.8]
c. Calculate the value of the internal energy change, ∆E˚, for the combustion reaction.
d. If ∆G˚ for the combustion is -5230 kJ. mol -1 of octane, calculate the value of ∆S˚. Comment on the sign of ∆S˚ relative to the
equation written above.
e. State whether the heat associated with the combustion of liquid octane in a bomb calorimeter represents ∆H˚ or ∆E˚. Explain
your reasoning.
3.
(13%) Phosphoric acid, H3PO4, ionizes according to the equations,
s H+(aq) + H2PO4–(aq)
K1 = 7.1 x 10-3
H2PO4–(aq) s H+(aq) + HPO42-(aq)
K2 = 6.2 x 10-8
HPO42-(aq) s H+(aq) + PO43-(aq)
K3 = 4.5 x 10-13
H3PO4(aq)
a. Write the equilibrium expression for the ionization of H3PO4 and find the pH of a 1.5 M solution of H3PO4.
b. A student is asked to prepare a phosphate buffer with a pH of 7.00. Identify the species that should be used in this solution and
calculate their ratio.
c. Assume that 50.0 mL of the buffer solution in b. are available in which the more abundant buffer species has a concentration
of 0.10 M. Determine the [H +] in this solution after 2.0 x 10-3 mol of NaOH are added.
d. Determine the [H+] in a 0.20 M solution of Na3PO4.
4.
(13%) An electrochemical cell is constructed with a piece of copper wire in a 1.00 M solution of Cu(NO3)2 and a piece of
chromium wire in a 1.00 M solution of Cr(NO3) 3.
The standard reduction potentials for Cr3+(aq) and Cu2+(aq) are:
Cr3+(aq) + 3e– ---> Cr(s)
Cu 2+(aq) + 2e– ---> Cu(s)
–0.744 V
0.340 V
a. Write a balanced equation for the spontaneous reaction that occurs in this cell and calculate the potential it produces.
b. Sketch a diagram for this cell.
i. Label the anode.
ii. Show the direction of electron flow in the external circuit.
iii. Show the direction of movement of nitrate ions. Explain.
c. The cell is allowed to operate until the [Cu 2+] = 0.10 M.
i. Find the [Cr3+].
ii. Calculate the cell potential at these concentrations.
Not valid for use as an USNCO National Examination after April 27, 2003.
Page 2
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Water is added to magnesium nitride.
b. Excess carbon dioxide is bubbled through a solution of calcium hydroxide.
c. Acidic solutions of potassium dichromate and iron(II) chloride are mixed.
d. Solutions of lead acetate and sulfuric acid are mixed.
e. Excess concentrated sodium hydroxide is added to a solution of zinc nitrate.
f. Fluorine-18 undergoes positron emission.
6.
(14%) Account for the following observations about chemical kinetics.
a. Reactions involving molecular chlorine often have nonintegral rate laws.
b. The rates of exothermic reactions increase when their temperatures are increased.
c. Two reactions, A and B, have rate constants that are equal at 25˚C but the rate constant for reaction A is much greater than
that for reaction B at 35˚C.
d. The rates of reactions catalyzed by complex molecules, such as enzymes, increase with an increase in temperature up to a
certain point above which they decrease again.
7.
(12%) The atoms C, N and O can be arranged in three different orders to form negative ions, i.e. CNO–, CON–, and NCO–. Salts
of one of these ions are stable. Salts of one are explosive and salts of one are unknown at this time.
a. Write Lewis electron dot structures for each atomic arrangement.
b. For each arrangement,
i. find the formal charge on each atom.
ii. use these formal charges to identify the most and least stable arrangements. Explain your reasoning.
c. Predict the geometry of the most stable atom arrangement and identify the type of hybridization used by the central atom in
this structure.
8.
(11%) Three compounds; X, Y, Z, have the formula C3H8O.
a. Write structural formulas for three different compounds with the formula C3H8O.
b. The boiling points of the compounds are; X 10.8˚C, Y 82.4˚C, Z 97.4˚C. Assign each boiling point to one of the structures in
part a. and account for these boiling points on the basis of the molecular structures and the types of intermolecular forces in
each.
c. A fourth compound with the formula C2H4O2, has the same molar mass as the three compounds above and boils at 117.9˚C.
Propose a structure for this compound and account for its higher boiling point.
END OF PART II
Not valid for use as an USNCO National Examination after April 27, 2003.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
N A gas constant
R Planck’s constant
h
°C gram
g pressure
P
c heat capacity
C p rate constant
k
C hour
h retention factor
Rf
E joule
J second
s
Ea kelvin
K speed of light
c
H kilo- prefix
k temperature, K
T
S liter
L time
t
volt
V
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
E = Eο –
 k2  Ea  1 1 
=  − 
 k1  R  T1 T2 
– ∆H   1 
ln K = 
+c
 R  T
RT
ln Q
nF
ln
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(277)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 27, 2003.
Page 4
2003 U.S. NATIONAL
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART II
1.
Facts from the problem.
• 0.1152 g sample contains C, H, N, O.
• 0.0912 g H2 O are recovered.
• Conversion of carbon dioxide to carbonic acid has 1 to 1 stoichiometry,
•
1 CO2 ! 1 H2 CO3 .
• Carbonic acid is titrated with 28.81 mL of 0.3283 M NaOH
•
H2 CO3 + 2 NaOH ! Na2 CO3 + 2 H2O
• Volume of N2 collected is 225.0 mL at 65.12 torr and 25 o C.
• Molar mass of the substance is 150 g/mol
a.
i. H2 O
(1 pt)
moles H2O = 0.09912 g ×
ii.
1 mol H2O
= 5.501 × 10−3 mol H2O
18.02 g H2O
CO2 (2 pts)
moles CO2 =
0.3283 mol NaOH
1 mol CO2
× 0.02881 L soln ×
= 4.729 × 10−3 mol CO2
L soln
2 mol NaOH
iii. N2 (1 pt)
moles N2 =
b.
i.
pV
(65.12 torr)(0.2250 L )
=
= 7.879 × 10−4 mol N2
RT (62.4 L ⋅ torr ⋅ mol-1 ⋅ K -1 )(298 K)
C (1 pt)
g C = 4.729 × 10 -3 mol CO 2 ×
ii.
1 mol C
12.01 g C
×
= 0.05680 g C
1 mol CO 2 1 mol C
H (1 pt)
g H = 5.501 × 10 -3 mol H 2O ×
2 mol H
1.008 g H
×
1 mol H 2O 1 mol H
= 0.01109 g H
iii. N (1 pt)
g N = 7.879 × 10 -4 mol N 2 ×
2 mol N
14.01 g N
×
1 mol N 2 1 mol N
iv. O (1 pt)
g O = g sample – (g C + g H + g N)
= 0.1152 g – (0.05680 g + 0.01109 g + 0.02208 g)
= 0.1152 g – 0.08997 g = 0.02523 g O
c. Find the empirical formula of the compound. (2 pts)
Key for 2003 USNCO National Exam, Part II
= 0.02208 g N
C : 0.05680 g C ×
1 mol C
12.01 g C
= 4.729 × 10−3 / 1.576 × 10−3 = 3.001 ≈ 3
H : 0.01109 g H ×
1 mol H
1.008 g H
= 1.100 × 10−2 / 1.576 × 10−3 = 6.980 ≈ 7
N : 0.02208 g N ×
1 mol N
14.01 g N
= 1.576 × 10−3 / 1.576 × 10−3 = 1.000 = 1
1 mol O
= 1.577 × 10−3 / 1.576 × 10−3 = 1.001 ≈ 1
16.00 g O
Therefore the empirical formula is C3H7 NO
O : 0.02523 g N ×
d. Find the molecular formula. (2 pts)
The molar mass of the empirical formula is 73.10
Compare this molar mass to the measured molar mass.
150 g ⋅ mol
= 2.05 ≈ 2
73.1 g ⋅ mol
Therefore the molecular formula is C6 H14N2 O2
2.
a. Write balanced equations to represent the processes responsible for K1 and K2.
2 C8 H18 (l) + 25 O2 (g) ! 16 CO2 (g) + 18 H2O (g) (1 pt)
b. Determine the molar enthalpy of formation, ∆Hf˚, for liquid octane, C8H18(l). (4 pts)
∆H rxn = 8∆H f (CO 2 ) + 9∆H f (H 2O) − ∆H f (C 8H18 )
−5090 kJ = 8(−393.5) kJ + 9(−241.8) kJ − ∆H f (C 8H18 )
∆H f (C 8H18 ) = −3148 − 2176.2 + 5090 kJ
= -234.2 kJ
c. Calculate the value of the internal energy change, ∆E˚, for the combustion reaction.
(4 pts)
∆H = ∆E + ∆nRT so ∆E = ∆H − ∆nRT
= –5090000 J – (4.5 mol)(8.314 J. mol-1 . K-1 )(373 K)
= –5090000 J – 13955 J
= –5104000 J = -5104 kJ /mol C8 H18(l)
-1
d. If ∆G˚ for the combustion is -5230 kJ. mol of octane, calculate the value of ∆S˚. Comment on the sign of ∆S˚ relative to the
equation written above . (3 pts)
∆G° = ∆H ° − T∆-1S °
-5230 kJ . mol = -5090 kJ . mol-1 – 373 K (∆So )
so
−5090 kJ ⋅ mol-1 + 5230 kJ ⋅ mol-1
∆S° =
373 K
= 0.375 kJ.mol-1.K-1
The increase in ∆S˚ is consistent with the formation of more moles of gas during the reaction.
Key for 2003 USNCO National Exam, Part II
e. State whether the heat associated with the combustion of liquid octane in a bomb calorimeter represents ∆H˚ or ∆E˚. Explain
your reasoning. (1 pt)
Heat in a bomb calorimeter is ∆E˚ (q at constant volume) - no credit unless there is a discussion of zero
work under constant volume.
3.
a.
Write the equilibrium expression for the ionization of H3PO4 and find the pH of a 1.5 M solution of H3PO4.
(1 pt for
equation, 1 pt for solution)
[H ][H PO ]
=
+
Ka
–
2
4
[H 3PO 4 ]
for a 1.5 M solution, assuming no initial concentration of reactants and that the amount of phosphoric
acid that reacts is small compared to the original volume,
x2
1.5
Solving for x, x = 1.5( 7.1 × 10−3 ) = 0.103.
Use successive approximations to check that the amount of reacting phosphoric acid doesn’t change the
answer… plug in 1.397 (1.5-0.103=1.397) rather than 1.5,
7.1 × 10−3 =
x = 1.397( 7.1 × 10−3 ) = 0.100 the change is small enough to accept this answer.
b. A student is asked to prepare a phosphate buffer with a pH of 7.00. Identify the species that should be used in this solution and
calculate their ratio. (1 pt for correct identification of salts, 1 pt for solution of ratio)
it would be Ka2. Thus the species that would
To obtain a pH of 7 the Ka should be close– to 7. In this case
2be present in this buffer should be H2 PO4 (aq) and HPO4 (aq).
The ratio can be found by the equation,
2–
 HPO 2– 
HPO 4
4
0.62
−8
−7 

6.2 × 10 = 1.0 × 10
and the ratio is,
= 0.62 or
–
–
 H 2PO 4 
1
H 2PO 4


[
[
c.
]
]
[
[
]
]
Assume that 50.0 mL of the buffer
solution in b. are available
in which the more abundant buffer species has a concentration
+
-3
of 0.10 M. Determine the [H ] in this solution after 2.0 x 10 mol of NaOH are added. (1 pt for correct calculation of
each concentration)
–
H2 PO4 (aq) is the more abundant species in the buffer from part b.
–
–
0.050 L × 0.10 mol ⋅ L-1 H 2PO 4 = 0.0050 mol H 2PO– 4 initially
–
0.0020 mol of OH is added, so the amount
of H2 PO4 (aq) left is (1 to 1 stoichiometry)
–
0.0050 – 0.0020 = 0.0030 mol– H2 PO4 (aq)
If the concentration of 2-H2 PO4 (aq) is 0.10 M the ratio calculated earlier indicates that the initial
concentration of HPO4 (aq) must be 0.62 × 1 = 0.062 M, so
2–
2–
0.050 L × 0.062 mol ⋅ L-1 HPO 4 = 0.0031 mol HPO 4 initially
2and 0.0020 mol are formed in the reaction so there is 0.0051 mol HPO4 (aq) present.
Calculating the amount of hydrogen ion present,
0.0051
6.2 × 10−8 = [H + ]
, solving for [H+] gives, 3.65 × 10-8 M
0.0030
Key for 2003 USNCO National Exam, Part II
d.
+
Determine the [H ] in a 0.20 M solution of Na3PO4. (1
pt for determining Kb, 1 pt for noting the need for successive
approximation in the calculation of phosphate ion, 1 pt for final solution)
PO43– + H2O ! HPO4 2– + OH– and
K
1.0 × 10−14
Kb = w =
= 0.0222
K a 3 4.5 × 10−13
Obtain concentration of hydroxide, initially assume no reaction of phosphate and let x = [OH–],
x2
0.0222 =
so x = [OH – ] = 6.67 × 10−2 now for phosphate, 0.20 – 0.0667 = 0.1333
0.20
x2
0.0222 =
so x = [OH – ] = 5.44 × 10−2 now for phosphate, 0.20-0.0544 = 0.1456
0.1333
x2
0.0222 =
so x = [OH – ] = 5.70 × 10−2 now for phosphate, 0.20-0.0570 = 0.143
0.1456
x2
0.0222 =
so x = [OH – ] = 5.63 × 10−2 which is acceptably close to the previous iteration, now use
0.143
Kw to calculate [H+],
1.0 × 10−14
+
H
=
[ ] 5.63 × 10−2 = 1.77 × 10−13
4.
a. Write a balanced equation for the spontaneous reaction that occurs in this cell and calculate the potential it produces. (3 pts)
2Cr(s) + 3 Cu2+(aq) r 2 Cr3+(aq) + 3 Cu(s)
o
E = Eox + Ered = 0.744 + 0.340 = 1.084 V
b. Sketch a diagram for this cell. (5 pts for sketch with proper labels. Points taken of for incorrectly labeled components, etc.)
i. Label the anode.
ii. Show the direction of electron flow in the external circuit.
iii. Show the direction of movement of nitrate ions. Explain.
volts
electrons
_
Cr
node
Cu
NO3
r +
Key for 2003 USNCO National Exam, Part II
u
+
2+
c. The cell is allowed to operate
until the [Cu ] = 0.10 M.
3+
i. Find the [Cr ]. (2 pts)
ii. Calculate the cell potential at these concentrations. (3
pts)
[Cu2+] goes from 1.0 M to 0.10 M, so ∆[Cu2+] is –0.90
∆[Cr3+] = 0.90 × 2/3 = 0.60 so, [Cr3+] = 1.60
plug these values into the equation,

3+ 2 
RT  [Cr ] 
o
E=E −
nF  [Cu 2+ ] 3 



0.0257 (1.60) 2 
E = 1.084 −

 = 1.05 V
6  (0.10) 3 
5. (Note that balanced chemical equations are note required.)
(1 pt for each reactant and 2 points for products (usually 1 for each product) then the total point value was
multiplied by 2/3 to scale to 12 pts)
a.
b.
c.
d.
e.
f.
Mg3 N2 + H2O ! Mg(OH)2 + NH3
CO2 + OH– ! HCO3–
Cr2 O7 2– + Fe2+ + H+ ! Cr3+ + Fe3+ + H2O
Pb2+ + C2 H3 O2 – + H+ + SO4 2– ! PbSO4 + HC2H3 O2
Zn2+ + OH– ! Zn(OH)4 2–
18
→188 O + +10β
9F 
6.
a. (3 pts) Cl2 often dissociates to Cl atoms, which react individually. If these atoms participate in the rate
determining step the overall rate equation is proportional to [Cl•] = [Cl2]1/2
b. (3 pts) All reactions increase in rate with an increase in temperature due to an increase in the collision
frequency and the increase in fraction of species with high velocities. Higher velocity particles impart
more energy to collisions in which they participate so those collisions are more likely to exceed Ea. The
exothermicity or endothermicity of a reaction has no bearing on its kinetics or the effect of temperature.
c. (4 pts) Reactions A and B must have different activation energies. When the log (or ln) of the rate is
constant for a reaction is plotted versus 1/T the slope is the Ea. Reactions A and B have different slopes
which cross at 25 o C. The Ea for reaction A is greater because an increase in temperature affects it’s rate
constant more.
d. (4 pts) Reaction rates increase with higher temperature because of an increase in collision rate and an
increase in the fraction of molecules that have the necessary energy to react. At still higher temperatures
the enzyme is denatured so that it no longer is an effective catalyst.
7.
(1 pt for each correct Lewis structure (3 total)
1 pt for correct formal charges on each Lewis structure (3 total)
1 pt for correct identification of most stable structure / 1 pt for correct reasoning
1 pt for correct identification of least stable structure / 1 pt for correct reasoning
1 pt for correctly identifying the structure as linear
1 pt for correct hybridization.)
Key for 2003 USNCO National Exam, Part II
a.
–
–
C
N
C
O
O
N
b.
C
O
–
–
i.
–
N
C
N
O
C
-1
+1
-1
-1
O
+2
–
N
N
C
O
-2
0
0
-1
ii. N C O is the most stable arrangement because the formal charges are the lowest in this structure.
C O N is the least stable structure because the formal charges are the greatest
in this structure
c. N C O as drawn would be linear because there are two charge centers around the
central atom. The central atom will be sp hybridized.
8.
a. The structural formulas would be, (1 pt for each)
O
o
X (b.p. 10.8 C)
CH3
CH2
CH3
CH
CH3
OH
CH3
Y (b.p. 82.4 o C)
Z (b.p. 97.4 oC) CH3 CH2 CH2
OH
Compound Y and Z would have hydrogen bonding where compound X would not. Therefore X would
have the weakest intermolecular forces and the lowest boiling point (of 10.8 oC) (2 pts)
b. The linear shape of compound Z would allow for stronger dispersion forces than compound Y.
Therefore Y should have a lower boiling point than Z. (4 pts)
c. The possible Lewis structures include,
O
O
CH3
C
C
OH
H
O
OH
CH2
CH3
C
C
O
OH
OH
H
H
C
H
All of these would have increased hydrogen-bonding relative to the three compounds in part a.
Key for 2003 USNCO National Exam, Part II
2003 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Brookline High School, Brookline, MA
Chair
Nancy Devino, ScienceMedia Inc., San Diego, CA
Lucy Pryde Eubanks, Clemson University, Clemson, SC
Sheldon L. Knoespel, Michigan State University, East Lansing, MI
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
Linda Weber, Natick High School, Natick, MA
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 27, 2003.
Page 1
2003 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
Sani-Flush®, a commercial toilet bowl cleaner, contains sodium bisulfate (sodium hydrogen sulfate) and
sodium carbonate as its active ingredients. Other ingredients include sodium chloride, sodium lauryl sulfate,
talc and some fragrance. Given the information from the manufacturer that the sodium bisulfate is
substantially in excess compared to the sodium carbonate, carry out an experiment to determine the percent by
weight of the sodium carbonate in a sample of the product.
Lab Problem 2
You have been provided with eight vials, each of which is labeled with a number from 1 to 8. Each vial
contains one of the following chemicals:
Na3PO4, NH4Cl, ZnCl2, KNO3, Mg(OH)2, Pb(NO3)2, CaCO3, Na2SO3.
You are allowed to use distilled water, test tubes or well plates, and only TWO additional reagents from the
following choices:
6 M H2SO4, 6 M HCl, 6 M AgNO3, phenolphthalein indicator solution
You must designate your choice of reagents prior to the start of your testing.
Page 2
Not valid for use as an USNCO National Examination after April 27, 2003.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 27, 2003.
Examiner’s Initials:
Page 3
2. Record your data and other observations.
3. Show your calculations.
Page 4
Not valid for use as an USNCO National Examination after April 27, 2003.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name : ________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
When you wish to request the optional reagents, return to the Examiner with this sheet.
I request these additional reagents:
Not valid for use as an USNCO National Examination after April 27, 2003.
Examiner’s Initials:
Page 5
2. Record your data and other observations.
3. Identify the substance in each numbered vial.
Vial #
Contains
1.
2.
3.
4.
5.
6.
7.
8.
Page 6
Not valid for use as an USNCO National Examination after April 27, 2003.
4. Explain clearly how you arrived at the identity of each of the vial’s contents.
Not valid for use as an USNCO National Examination after April 27, 2003.
Page 7
2003 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Instructions
Directions to the Examiner:
Thank you for administering the 2003 USNCO laboratory practical on behalf of your Local
Section. It is essential that you follow the instructions provided, in order to insure consistency of results
nationwide. There may be considerable temptation to assist the students after they begin the lab exercise.
It is extremely important that you do not lend any assistance or hints whatsoever to the students once
they begin work. As in the international competition, the students are not allowed to speak to anyone
until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab
station or table when the students enter the room. The equipment should be initially placed so that the
materials used for Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry
Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to
reason through a laboratory problem and communicate its results. Do not touch any of the equipment in
front of you until you are instructed to do so.
One of this year’s problems requires the use of a plastic syringe with a Luer-lock® tip cap.
Show a syringe and Luer-lock® tip cap.
This problem also requires you to use a balance, which is located _____________________________.
Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques
help to minimize the amount of materials you use, thereby increasing safety and minimizing waste.
Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and
reaction plates.
Show a 5-mL Beral-type pipet, and show a 24-well reaction plate or small test tubes.
You will be asked to complete two laboratory problems. All the materials and equipment you may want
to use to solve each problem has been set out for you and is grouped by the number of the problem. You
must limit yourself to this equipment for each problem. You will have one hour and thirty minutes to
complete the two problems. You may choose to start with either problem. You are required to have a
procedure for each problem approved for safety by an examiner. (Remember that approval does not
mean that your procedure will be successful–it is a safety approval.) When you are ready for an
examiner to come to your station for each safety approval, please raise your hand.
Page 1
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash
off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate
procedures for disposing of solutions at the end of this lab practical are:
____________________________________________________________________________________
____________________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. Are there any questions before we begin?
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. There is a
periodic table on the last page of the booklet. Allow students enough time to read the brief cover
directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out
all information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for
your cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the
lab practical with the students. They can learn about possible observations and interpretations and you
can acquire feedback as to what they actually did and how they reacted to the problems. After this
discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be
extremely useful to the Olympiad subcommittee as they prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron
sheets from Part I, and the “Blue Books” from Part II to this address:
ACS DivCHED Exams Institute
Department of Chemistry
University of Wisconsin – Milwaukee
US Postal Service:
P.O. Box 413
Milwaukee, WI 53201
FedEx or UPS:
3210 N Cramer Street
Milwaukee, WI 53211
Tuesday, April 29, 2003, is the absolute deadline for receipt of the exam materials at the Examinations
Institute. Materials received after this deadline CANNOT be graded.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT
SCHEDULE FOR GRADING THIS EXAMINATION.
Page 2
EXAMINER’S NOTES
Lab Problem #1: Materials and Equipment
Each student should have available the following equipment and materials:
•
•
•
•
•
•
•
•
•
•
•
60-mL plastic syringe with Luer-lock® tip cap
Small plastic cap to hold sample inside syringe
Electronic balance, 0.01 grams. One balance can serve 3-4 students. Please do not
substitute a milligram balance. Student processing skills, not precision, are being evaluate
here.
Scoopula or spatula
400-mL or 600-mL beaker
Weighing boats or weighing paper
Capped vial or small beaker with film covering, at least 50 mL capacity
Easy access to sinks
Supply of paper towels
1 pair safety goggles
1 lab coat or apron (optional)
Lab Problem #1: Chemicals
Each student will need:
Approximately 5 grams solid sample of Sani-Flush® Note: Solid Sani-Flush® is sold in the
cleaning supply section of most supermarkets and convenience stores.
Bottle of distilled water, at least 100 mL
Examiner access to room temperature and barometric pressure
Lab Problem #1: Notes
1.
Note that the examiner will need to initial each student’s experimental
plan to be sure that safety is considered.
2.
You will need to provide students with the room Celsius temperature and barometric
pressure (in mm Hg or atm). These data can be recorded on the board and also
announced at the start of the experiment.
3.
The small plastic cap can be from a soda bottle and should be able to fit with room
inside the syringe without interfering with plunger operation.
4.
The Sani-Flush should be in solid crystalline form.
5.
Be sure that the syringe caps fit tightly over the tip of the syringe.
6.
Check that the plastic bottle cap fit easily and completely inside the syringe and does the
plunger slide easily when pushed.
7.
It is your responsibility to ensure that students wear their safety goggles during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. You will
also need to give students explicit directions for handling spills and for disposing of
waste materials following approved safety practices for your examining site. Please
check and follow procedures appropriate for your site.
Page 3
Lab Problem #2: Materials and Equipment
Each student will need:
• Eight vials, with caps, 10-25 mL capacity each
• 24-well plate or several spot plates or approximately fifteen 13 x 100mm test tubes
• tube rack (if test tubes are used)
• A disposal container for chemical waste (designated for heavy metal waste)
• 3-4 Beral-style pipets, ungraduated 5-mL capacity
• 400-mL beaker to hold vials
• paper towels
• easy access to sink
• one pair safety goggles
• one pair lab coat or apron (optional)
Note: The 24-well plates and Beral-style pipets can be purchased through Educational
Innovations or Micro Mole Scientific among other vendors.
Lab Problem #2: Chemicals
3-4 grams each of Na3PO4, NH4Cl, ZnCl2, KNO3, Mg(OH)2, Pb(NO3)2, CaCO3, and Na2SO3 in
vials labeled 1 – 8.
Since this is an identification experiment, do not identify the contents of each vial! The order
should be as follows:
Vial 1
NH4Cl
Vial 2
CaCO3
Vial 3
Vial 4
Vial 5
Pb(NO3)2 Mg(OH)2 KNO3
Vial 6
Na2SO3
Vial 7
Na3PO4
Vial 8
ZnCl2
Bottle of distilled water, at least 100 mL
Additionally, the examiner will prepare 250 mL of 6 M H2SO4, 250 mL of 6 M HCl, 250 mL of 1
M AgNO3, 250 mL of 0.1% phenolphthalein. It is suggested that these reagents be placed in 400
mL beakers. Fill the number of 5 mL Beral-type pipets with each reagent as the number of
students present. Student may select only two of these four reagents to use in their experiment.
When they request the two solutions, be sure to indicate these on their answer sheets. They may
not change their choices once they’ve begun using these two solutions, nor may they use more
than one pipet of each of the two solutions selected.
Lab Problem #2: Notes
1.
Note that the examiner will need to initial each student’s experimental plan to be sure
that safety is considered.
2.
Make sure that the vials are NOT labeled with their contents.
3.
Do not let students know which reagents might be better choices to select. Remind them
that they may only select one pipet of each of two of these additional reagents, without
refills.
4.
The hydrated form of several of these salts (Na3PO4, ZnCl2) should be used, not the
anhydrous form.
5.
The phenolphthalein solution is a 0.1% solution dissolved in ethanol. Most pre-prepared
solutions are the standard 0.1% solution.
Page 4
2003 U.S. NATIONAL
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART III
Problem 1. Sani-Flush®, a commercial toilet bowl cleaner, contains sodium bisulfate and sodium carbonate as
its active ingredients and other ingredients such as sodium chloride, sodium lauryl sulfate, talc and some
fragrance. Given the information from the manufacturer that the sodium bisulfate is substantially in excess
compared to the sodium carbonate, carry out an experiment to determine the percent by weight of the sodium
carbonate in a sample of the product.
Experimental Plan:
2HSO4 – + CO32– ! CO2 + H2O + 2SO4 2–
A good plan consisted of weighing a sample of Sani-Flush® and determining the volume of CO2 produced when
the sample is added to water.
For example, a good plan might include these steps,
1. Weigh approximately 1 g of Sani-Flush®.
2. Add to syringe.
3. Draw water into the syringe and cover with cap.
4. Determine the volume of gas (CO2 ) produced.
5. Use ideal gas equation to calculate moles of CO 2
6. Convert moles of CO2 to moles of sodium carbonate and then mass of Na2 CO3
7. Divide mass of Na2 CO3 by sample mass to find percentage.
8. Repeat with second sample.
An average plan was either missing one of these components or had a procedure based solely on the difference in
mass. Such a procedure is subject to greater error due to the smaller change in mass relative to the change in
volume.
A weak plan had minimal detail about how the experiment would be conducted.
Data and other Observations:
An example of good work on observations and data recording would be,
1.45 g sample placed in syringe.
Set plunger at 5.0 mL.
Drew in 14.0 mL of H2O.
Gas was evolved and the solution was blue.
Plunger reached 47.8 mL after the reaction was complete.
mass Sani-Flush®
Trial 1
Trial 2
2.05 g
1.60 g
initial plunger
level
5.0 mL
5.0 mL
Temperature is 23 oC and pressure is 755 mmHg
Key for 2003 USNCO National Exam, Part III
H2 O volume
added
14.0 mL
15.0 mL
final plunger
level
47.8 mL
39.0 mL
Sample Calculations:
Sample 1: 47.8 – 14 = 33.8 mL
removing water volume: 33.8 mL – 5.0 mL = 28.8 mL gas evolved.
 755 mmHg


-1 (0.0288 L )
760 mmHg ⋅ atm 
PV 
mole of gas: PV = nRT therefore n =
=
RT
(0.0821L ⋅ atm ⋅ mol-1 ⋅ K -1 )(296 K)
= 0.00118 mol CO2
106 g
convert to mass: 0.00118 mol CO2 = 0.00118 mol Na 2CO 3 ×
= 0.125 g Na 2CO 3
1 mol
calculate percentage:
%=
0.125 g
× 100 = 6.1%
2.05 g
Problem 2: You have been provided with eight vials, each of which is labeled with a number from 1 to 8. Each
vial contains one of the following chemicals:
Na3 PO4, NH4Cl, ZnCl2 , KNO3 , Mg(OH)2 , Pb(NO3 )2 , CaCO3 , Na2 SO3.
You are allowed to use distilled water, test tubes or well plates, and only TWO additional reagents from the
following choices:
6 M H2SO4, 6 M HCl, 6 M AgNO3 , phenolphthalein indicator solution
You must designate your choice of reagents prior to the start of your testing.
Experimental Plan
A good plan involved stating that samples of each of the eight unknowns would be placed in wells of the spot plates,
distilled water would be added and observations would be made regarding the solubility of the salts. Then, each of th
SELECTED reagents would be added and specific identifying tests would be detailed. It should be noted that several
combinations of reagents could have been used to identify the unknowns.
For example, a good test might include these tests.
1. Place each sample into each of three wells in the spot plate. Add distilled water to each and record solubility r
2. Two of the samples will not dissolve. Add HCl to these. One sample should dissolve and the other should fiz
former is magnesium hydroxide, the latter is calcium carbonate.
3. Add HCl to the other solutions. Lead (II) nitrate will form a precipitate. Sodium sulfite will release sulfur dio
which can be identified by odor.
4. Add silver nitrate to the other four solutions. The sample which does not react is potassium nitrate. The yello
precipitate is produced by sodium phosphate and the two white solids are zinc chloride and ammonium chlorid
5. Prepare additional solutions of the two chlorides and the identified sodium phosphate. Add the latter to the
chlorides. The zinc chloride will form a precipitate while the ammonium chloride will not react.
An average plan did not detail the expected results of specific tests.
A weak plan did not include solution formation or specific results of tests.
Key for 2003 USNCO National Exam, Part III
Observations and results:
An example of good work on observations and results would be:
Reagent
Sample
1
Sample
2
Sample
3
Sample
4
Sample
5
water
soluble
soluble
fizz
No Rx
AgNO3
No Rx
White
ppt
soluble
White
ppt
insoluble
HCl
insoluble
dissolve
s
Na3PO4
No Rx
Key for 2003 USNCO National Exam, Part III
No Rx
Sample 6
Sample 7
Sample
8
soluble
Sharp
odor
soluble
soluble
No Rx
Yellow
Ppt
No Rx
White
ppt
White
ppt
2004 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM Part I
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
David W. Hostage, Taft School, CT
William Bond, Snohomish High School, WA
Alice Johnsen, Bellaire High School, TX
Peter E. Demmin (retired), Amherst Central High School, NY
Marian Dewane, Centennial High School, ID
Adele Mouakad, St. John’s School, PR
Ronald O. Ragsdale, University of Utah, UT
Dianne Earle, Boiling Springs High School, SC
Jacqueline Simms, Sandalwood Sr. High School, FL
Michael Hampton, University of Central Florida, FL
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 19, 2004, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
ABBREVIATIONS AND SYMBOLS
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G molar mass
A frequency
ν mole
N A gas constant
R Planck’s constant
°C gram
g pressure
c heat capacity
C p rate constant
C hour
h retention factor
E joule
J second
Ea kelvin
K temperature, K
H kilo– prefix
k time
S
liter
L volt
K
milli– prefix
m
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
T
t
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
1 atm = 760 mmHg
EQUATIONS
E = Eo −
1
1A
1
H
k  E  1 1 
ln 2  = a  − 
 k1  R  T1 T2 
 −∆H  1 
ln K = 
  + constant
 R  T 
RT
ln Q
nF
PERIODIC TABLE OF THE ELEMENTS
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
1.008
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
DIRECTIONS
! When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
! There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
! Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which element is obtained commercially from seawater?
(A) bromine
(B) gold
(C) iron
(D) oxygen
2. Which solution can serve as both reactant and indicator
when it is used in redox titrations?
(A) FeNH4(SO4)2
(B) KMnO4
(C) H2C2O4
(D) Na 2S 2O3
(A) N2 and H2O
(B) N2O and H2O
(C) NO and H2
(D) N2, H2 and O2
4. Which method should be used to extinguish burning
magnesium metal?
(A) Blanket it with CO2
(B) Blow on it.
(C) Dump sand on it.
(D) Pour water on it.
A
C
B
D
(A) A
(B) B
(C) C
(A) Ca
(B) Mn
(C) Ni
(D) Zn
8. What is the coefficient for OH- after the equation
_ Br2 + _ OH- r _ Br- + _ BrO3- + _ H2O
is balanced with the smallest integer coefficients?
3. What is formed when a solution of NH4NO2 is heated
gently?
5. Which letter indicates
where a thermometer
should be placed to
determine the boiling point
of a distillate?
7. A 1.871 gram sample of an unknown metallic carbonate
is decomposed by heating to form the metallic oxide and
0.656 g of carbon dioxide according to the equation
MCO3(s) r MO(s) + CO2(g)
What is the metal?
(A) 3
(B) 6
(C) 12
(D) 18
9. An ionic compound contains 29.08% sodium, 40.56%
sulfur and 30.36% oxygen by mass. What is the formula
of the sulfur-containing anion in the compound?
(A) S 2O32-
(B) S 2O42-
(C) S 2O52-
(D) S 2O62-
10. A solution is prepared
Vapor pressure (mmHg)
containing a 2:1 mol ratio of
C 2H4Br2
173
dibromoethane (C2H4Br2) and
dibromopropane (C 3H6Br2).
C 3H6Br2
127
What is the total
vapor pressure over the solution assuming ideal
behavior?
(A) 300 mmHg
(B) 158 mmHg
(C) 150 mmHg
(D) 142 mmHg
(D) D
6. A 50 mL sample of gas is collected over water. What will
be the effect on the calculated molar mass of the gas if
the effect of the water vapor is ignored? It will be
(A) high because of the mass of water in the collection
flask.
(B) high because of omitting the vapor pressure of the
water in the calculation.
(C) low because of the mass of water in the collection
flask.
11. A solution of magnesium chloride that is 5.10%
magnesium by mass has a density 1.17 g/mL. How many
moles of Cl- ions are in 300. mL of the solution?
(A) 0.368
(B) 0.627
(C) 0.737
(D) 1.47
12. Which aqueous solution has a freezing point closest to
that of 0.30 M C12H22O11?
(A) 0.075 M AlCl3
(B) 0.15M CuCl2
(C) 0.30 M NaCl
(D) 0.60 M C6H12O6
(D) low because of omitting the vapor pressure of the
water in the calculation.
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
Page 3
13. An unknown gas is
placed in a sealed
container with a fixed
volume. Which of the
characteristics listed
change(s) when the
container is heated
from 25 ˚C to 250 ˚C?
I The density of the gas
II The average kinetic energy
of the molecules
III The mean free path between
molecular collisions
(B) II only
(C) III only
(D) I and II only
14. Which gas has the same density at 546 ˚C and 1.50 atm
as that of O2 gas at STP?
(B) NH3
(C) SO2
(D) SO3
15. Which plot involving vapor pressure (VP) and absolute
temperature results in a straight line?
(A) VP vs T
(B) VP vs T-1
(C) ln VP vs T
(D) ln VP vs T-1
16. For a substance with the values of ∆Hvap and ∆Svap
given below, what is its normal boiling point in ˚C?
(∆Hvap = 59.0 kJ . mol -1; ∆Svap = 93.65 J. mol -1 .K-1)
(A) 357
(B) 630
(C) 1314
(D) 1587
17. What is the order of the boiling points (from lowest to
highest) for the hydrogen halides?
(A) HF < HCl < HBr < HI
(B) HI < HBr < HCl < HF
(C) HCl < HF < HBr < HI
(D) HCl < HBr < HI < HF
18. Of the three types of cubic lattices, which have the
highest and lowest densities for the same atoms?
Highest
(A) S˚200K is smaller because entropy decreases as
temperature increases.
(B) S˚200K is smaller because the surroundings are more
disordered at higher temperatures.
(A) I only
(A) N2
20. Which is the best description of the relationship between
the absolute entropies, S˚, of solid water at 100 K and at
200 K?
Lowest
(C) S˚100K = S˚200K = because water is in the solid phase
at both temperatures.
(D) S˚200K is larger because the vibration of the
molecules increases as temperature increases.
21. For the reaction,
CH4 + Cl2
r CH3Cl + HCl
which expression
gives ∆H?
Bond dissociation
energies
C-H
C-Cl
Cl-Cl
H-Cl
kJ . mol-1
(A) ∆H = (413 + 328) - (242 + 431)
(B) ∆H = (413 - 328) - (242 - 431)
(C) ∆H = (413 - 242) - (328 - 431)
(D) ∆H = (413 + 242) - (328 + 431)
22. Which phase change for water has positive values for
both ∆H˚ and ∆G˚?
(A) (l) r (s) at 250 K
(B) (l) r (s) at 350 K
(C) (l) r (g) at 350 K
(D) (l) r (g) at 450 K
23. When solid CuSO4 dissolves in water to make a 1M
solution, the temperature of the system increases. When
solid NH4NO3 dissolves in water to make a 1 M solution,
the temperature of the system decreases. Which
statement(s) must be correct for these dissolving
processes?
I ∆H˚ values for both processes have the same sign.
II ∆G˚ values for both processes have the same sign.
(A) simple cubic
body-centered cubic
(B) face-centered cubic
simple cubic
(C) body-centered cubic
face-centered cubic
(A) I only
(B) II only
D) face-centered cubic
body-centered cubic
(C) Both I and II
(D) Neither I nor II
19. For which reaction is ∆H (enthalpy change) most nearly
equal to ∆E (internal energy change)?
24. Which set of relationships could apply to the same
electrochemical cell?
(A) H2(g) + 1/2O2(g) r H2O(g)
(A) ∆G˚ > 0; E˚ = 0
(B) ∆G˚ < 0; E˚ = 0
(B) Cl2(g) + F2(g) r 2ClF(g)
(C) ∆G˚ > 0; E˚ > 0
(D) ∆G˚ < 0; E˚ > 0
(C) H2O(l) r H2O(g)
(D) 2SO3(g) r 2SO2(g) + O 2(g)
Page 4
413
328
242
431
25. The rate constant for
a reaction is affected
by which factors?
I increase in temperature
II concentration of the reactants
III presence of a catalyst
(A) I and II only
(B) I and III only
(C) II and III only
(D) I, II and III
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
26. The rate data given were obtained for the reaction,
2NO(g) + 2H2(g) r N2(g) + 2H2O(g)
What is the rate law for this reaction?
NO pressure (atm) H2 pressure (atm) Rate (atm. sec-1)
0.375
0.500
6.43 × 10-4
0.375
0.250
3.15 × 10-4
0.188
0.500
1.56 × 10-4
(A) Rate = k PNO
2
(B) Rate = k PNO
(C) Rate = k PNO PH2
2
2
(D) Rate = k PNO
PH 2
time
1 / concentration
concentration
log (concentration)
27. What is the order of a reaction that produces the graphs
shown?
31. H2S(aq) s H+(aq) + HS-(aq)
K = 9.5 × 10-8
+
2–
HS (aq) s H (aq) + S (aq)
K = 1.0 × 10-19
Given the equilibrium constants provided, what is the
equilibrium constant for the reaction;
S 2–(aq) + 2H+(aq) s H2S(aq)
K=?
(A) 9.5 × 10-27
(B) 9.7 × 10-14
(C) 9.5 × 1011
(D) 1.0 × 1026
32. Calculate the hydronium ion concentration in 50.0 mL of
0.10 M NaH2AsO4.
(K 1 = 6.0 × 10-3, K2 = 1.1 × 10-7 K3 = 3.0 × 10-12)
(A) 2.4 × 10-2
(B) 1.6 × 10-3
(C) 1.0 × 10-4
(D) 2.5 × 10-5
33. When the acids; HClO3, H3BO3, H3PO4, are arranged in
order of increasing strength, which order is correct?
(A) H3BO3 < H3PO4 < HClO 3
time
time
(B) HClO3 < H3BO3 < H3PO4
(A) zero order
(B) first order
(C) H3PO4 < HClO 3 < H3BO3
(C) second order
(D) some other order
(D) H3BO3 < HClO 3 < H3PO4
28. What is the rate law for the hypothetical reaction with the
mechanism shown?
2A
s intermediate 1 fast equilibrium
intermediate 1 + B r intermediate 2 slow
fast
intermediate 2 + B r A2B2
(A) Rate = k[A]2
(C) Rate = k[A][B]
(B) Rate = [B]2
2
(D) Rate = k[A] [B]
29. According to the Arrhenius equation: k = Ae -Ea/RT , a plot
of ln k against 1/T yields
(A) Ea as the slope and A as the intercept
(B) Ea/R as the slope and A as the intercept
(C) Ea/R as the slope and ln A as the intercept
(D) -Ea/R as the slope and ln A as the intercept
(A) II only
(B) I and II only
(C) I and III only
(D) I, II and III
35. A solution is 0.10 M in Ag+, Ca2+, Mg2+, and Al3+ ions.
Which compound will precipitate at the lowest [PO43-]
when a solution of Na3PO4 is added?
(A) Ag3PO4 (Ksp = 1 × 10-16)
(B) Ca3(PO4)2 (Ksp = 1 × 10-33)
(C) Mg3(PO4)2 (Ksp = 1 × 10-24)
(D) AlPO4 (Ksp = 1 × 10-20)
36. Which salt is significantly more soluble in a strong acid
than in water?
Reaction rate
30. Curves with the shape
shown are often observed
for reactions involving
catalysts. The level portion
of the curve is best
attributed to the fact that
34. A buffer solution results from mixing equal volumes of
which solutions?
I 0.10 M HCl and 0.20 M NH3
II 0.10 M HNO 2 and 0.10 M NaNO2
III 0.20 M HCl and 0.10 M NaCl
(A) PbF2
Reactant concentration
(A) product is no longer being formed.
(B) the reaction has reached equilibrium.
(C) all the catalytic sites are occupied.
(D) all the reactant has been consumed.
(B) PbCl2
(C) PbBr 2
(D) PbI2
37. What is the standard cell potential for the reaction,
2Cr(s) + 3Sn2+(aq) r 3Sn(s) + 2Cr3+(aq)
given the E˚ values shown?
Cr3+(aq) + 3e- r Cr(s) -0.744 V
Sn 2+(aq) + 2e- r Sn(s) -0.141 V
(A) 0.945 V
(B) 0.603 V
(C) -0.603 V
(D) -0.945 V
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
Page 5
38. How many electrons are needed in the balanced halfreaction for the oxidation of ethanol to acetic acid?
C 2H5OH r CH3COOH
(A) 1
(B) 2
(C) 3
(B) Cu 2+(aq)
(C) H+(aq)
(D) Zn2+(aq)
(A) N
(B) 22
(C) 6.1 × 10
(D) 3.8 × 109
41. When an aqueous solution of potassium fluoride is
electrolyzed, which of the following occurs?
(A) O2 and H+ are produced at one electrode and H2 and
OH- are formed at the other.
(B) O2 and OH- are produced at one electrode and H2
and H+ are formed at the other.
(C) Metallic K is formed at one electrode and O2 and H+
are formed at the other.
(D) Metallic K is produced at one electrode and
elemental F2 is produced at the other.
42. A CuSO4 solution is electrolyzed for 20. minutes with a
current of 2.0 ampere. What is the maximum mass of
copper that could be deposited?
(A) 0.20 g
(B) 0.40 g
(C) 0.79 g
(D) 1.6 g
43. Which experimental evidence most clearly supports the
suggestion that electrons have wave properties?
(A) diffraction
(A) 0
(D) deflection of cathode rays by a magnet
44. Which quantum number determines the number of
angular nodes in an atomic orbital?
(C) ml
(D) ms
45. Which element exhibits the greatest number of oxidation
states in its compounds?
(A) Ca
Page 6
(B) V
(C) 4
(D) 6
12
7N
18
8O
(B)
(C)
20
9F
(D)
20
10 Ne
49. According to the Lewis dot
O
C N
structure shown, what are the
formal charges of the O, C and N atoms, respectively, in
the cyanate ion?
(A) 0, 0, 0
(B) -1, 0, 0
(C) -1, +1, -1
(D) +1, 0, -2
50. The hybridization of As in AsF5 is best described as
(A) sp3
(B) sp4
(C) dsp3
(D) d2sp3
51. In which species do the atoms NOT lie in a single plane?
(A) BF3
(B) PF3
(C) ClF3
(D) XeF4
52. For which compound does the reaction,
MCO3(s) r MO(s) + CO2(g)
occur most readily?
(A) BeCO3
(B) MgCO3
(C) CaCO3
(D) BaCO3
53. The color of Co(H2O)62+ is best attributed to electronic
transitions
(A) between different n levels in the metal.
(B) between the metal's d orbitals.
(D) during ionization.
(C) photoelectric effect
(B) l
(B) 2
(C) from the Co2+ ion to water molecules.
(B) emission spectra
(A) n
(D) Cl
48. Which species is most likely to lose a positron (β+)?
40. The standard potential for the reaction
Cl2(g) + 2Br -(aq) ---> Br2(l) + 2Cl-(aq)
is 0.283 volts. What is the equilibrium constant for this
reaction at 25 ˚C?
4
(C) S
47. How many unpaired electrons are in a gaseous Fe2+ ion in
its ground state?
(A)
(A) 1.6 × 10-5
(B) P
(D) 4
39. Which is the weakest oxidizing agent in a 1 M aqueous
solution?
(A) Ag+(aq)
46. Of the elements given, which has the lowest ionization
energy?
(C) Cu
54. When the carbon-oxygen bonds in the species; CH3OH,
CH2O and CHO2- are arranged in order of increasing
length, which is the correct order?
(A) CH3OH < CH2O < CHO2(B) CH2O < CH3OH < CHO2(C) CHO - < CH OH < CH2O
2
3
(D) CH2O < CHO2- < CH3OH
(D) As
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
55. How many different trichlorobenzenes, C6H3Cl3, can be
formed?
(A) 1
(B) 2
(C) 3
(D) 4
56. What organic product is formed from the mild oxidation
of a secondary alcohol?
(A) acid
(B) aldehyde
(C) ether
(D) ketone
57. The compound with the formula, H 2NCH2CH2COOH, is
best classified as a(n)
(A) amide
(B) amino acid
(C) fatty acid
(D) nucleic acid
58. The reaction between which pair of reactants occurs the
fastest for [OH-] = 0.010 M?
(A) CH3CH2CH2CH2Cl + OH (B) (CH3)3CCl + OH(C) CH3CH2CH2CH2Br + OH(D) (CH3)3CBr + OH59. What is the major organic product formed from the
reaction of CH3CH=CH2 and HCl?
(A) CH3CHClCH3
(B) CH3CH2CH2Cl
(C) CH3CHClCH2Cl
(D) CH2ClCH=CH2
60. Fats and oils are formed from the combination of fatty
acids with what other compound?
(A) cholesterol
(B) glucose
(C) glycerol
(D) phenol
END OF TEST
Not valid for use as an USNCO Olympiad National Exam after April 19, 2004.
Page 7
National Olympiad 2004
Part 1
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
A
B
A
C
A
D
D
B
A
B
D
A
B
C
D
A
D
B
B
D
D
C
B
D
B
D
C
D
D
C
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Property of the ACS DivCHED Examinations Institute
Answer
D
D
A
B
D
A
B
D
D
D
A
C
A
B
B
C
C
A
D
C
B
A
B
D
C
D
B
D
A
C
2004 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
David W. Hostage, Taft School, CT
William Bond, Snohomish High School, WA
Alice Johnsen, Bellaire High School, TX
Peter E. Demmin (retired), Amherst Central High School, NY
Marian Dewane, Centennial High School, ID
Adele Mouakad, St. John’s School, PR
Ronald O. Ragsdale, University of Utah, UT
Dianne Earle, Boiling Springs High School, SC
Jacqueline Simms, Sandalwood Sr. High School, FL
Michael Hampton, University of Central Florida, FL
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a
score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 19, 2004, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front
of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper,
and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO National Exam after April 19, 2004.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
1.
2.
(12 %) A solution of copper(II) sulfate that contains 15.00% CuSO4 by mass has a density of 1.169 g/mL. A 25.0 mL portion of
this solution was reacted with excess concentrated ammonia to form a dark blue solution. When cooled, filtered and dried,
6.127 g of a dark blue solid were obtained. A 0.195 g sample of the solid was analyzed for ammonia by titrating with 0.1036 M
hydrochloric acid solution, requiring 30.63 mL to reach the equivalence point. A 0.150 g sample was analyzed for copper (II) by
titrating with 0.0250 M EDTA, (which reacts with Cu2+ in a 1:1 ratio). The endpoint was reached after 24.43 mL of the EDTA
were added. A 0.200 g sample was heated at 110 ˚C to drive off water, producing 0.185 g of the anhydrous material.
a. Determine the molarity of Cu2+ ions in the original solution.
b. Find the number of moles of Cu2+ in the 25.0 mL portion.
c. Calculate the percentages by mass in the prepared compound of;
i.
NH3
ii.
Cu 2+
iii.
H2O
iv.
SO42d. Use the results in part c. to determine the formula of the compound.
e. Assuming that Cu2+ is the limiting reactant in the synthesis, determine the percent yield.
(16%)
2NO2(g) + O 3(g) r N2O5(g) + O2(g)
∆H˚ = -198 kJ ∆S˚ = -168 J . K-1
Ozone reacts with nitrogen dioxide according to the equation above.
a. Calculate ∆H f˚ for NO2(g) in kJ.mol -1. [∆H f˚ kJ. mol -1; O3(g) 143, N2O5(g) 11]
b. Account for the sign of ∆S˚.
c. Calculate the value of ∆G˚ at 25 ˚C.
d. State and explain how the spontaneity of this reaction will vary with increasing temperature.
e. Use the rate data below to determine the rate law for the reaction of NO2(g) and O3(g)
NO2(g), M
O3(g), M
Rate M. s-1
0.0015
0.0025
4.8×10-8
0.0022
0.0025
7.2×10-8
0.0022
0.0050
1.4×10-7
f. Calculate the specific rate constant and give its units.
g. The following mechanisms have been proposed for this reaction. Discuss the suitability of each to account for the rate law
obtained.
Mechanism I
Mechanism II
NO2 + NO2 r NO3 + NO
slow
NO2 + O3 s NO3 + O2
fast
NO3 + NO2 r N2O5
fast
NO3 + NO2 r N2O5
slow
NO + O3
r NO2 + O2
fast
h. Describe and account for any change expected in the rate of this reaction as the temperature is increased.
3.
(10%) A popular lecture demonstration involves the sequential precipitation and dissolution of several slightly soluble silver
compounds beginning with a [Ag+ ] = 0.0050 M. Use the information below to answer the following questions about this
demonstration.
[Ksp values; AgCl 1.8×10-10, AgBr 5×10-13, AgI 8.3×10-17, Ag2SO4 1.4×10-5]
a. What must the [SO42-] be in order to start precipitation in a solution in which [Ag+] = 0.005 M?
b. State the order in which the halide ions should be added to a concentration of 0.10 M so that each precipitate will form from
the [Ag+] in equilibrium with the previous precipitate. Support your answer with appropriate calculations.
c. As a way of making this demonstration more striking, one of the silver halides in this series is dissolved by adding aqueous
ammonia before precipitating the next silver halide. Which silver halide(s) dissolve in 0.60 M NH3? Support your answer with
calculations. [Kf Ag(NH3)2+ 1.7×107]
Page 2
Not valid for use as an USNCO National Exam after April 19, 2004
4.
(12%) Aluminum metal is obtained commercially by electrolyzing Al2O3 mixed with cryolite (Na3AlF 6).
a. Explain why electrolysis is used rather than heating the Al 2O3 either directly or in the presence of C (as is done to extract
Fe or Zn from their ores).
b. State the purpose of the Na 3AlF 6.
c. Write the two half-reactions that occur during electrolysis and indicate which of the two occurs at the cathode
d. How many moles of electrons must pass through the cell to produce 5.00 kg of Al? (Assume 100% efficiency.)
e. Determine the current required (in amperes) if the aluminum in d. is produced in 10.0 hours.
f. Calculate the volume of gas formed in the process in d. at 25 ˚C and 720 mmHg.
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Concentrated nitric acid is added to iron(II) sulfide.
b. Acetic acid is added to solid calcium phosphate.
c. Small pieces of aluminum metal are added to concentrated sodium hydroxide solution.
d. Solutions of chromium(III) sulfate and barium hydroxide are mixed.
e. Concentrated hydrochloric acid is added to an aqueous solution of cobalt(II) nitrate.
f. Bromine gas is added to propene.
6.
(16%) Account for the following statements or observations in terms of atomic-, ionic- or molecular-level explanations.
a. Magnesium exists as +2 ions rather than +1 ions in all of its compounds despite the fact that the second ionization energy of
a magnesium atom is more than twice as great as the first ionization energy.
b. Titanium forms ions with different charges (+2, +3 and +4). The first two of these ions are colored while the last is colorless.
c. Carbon dioxide (CO 2) is a gas at room temperature but silicon dioxide (SiO2) is a high melting solid.
d. Nitrogen forms NF 3 but not NF5 whereas phosphorus forms PF3 and PF5. The trifluorides are both trigonal pyramidal and
the pentafluoride is trigonal bipyramidal.
7.
(12%) Explain these observations about nuclei.
a. Elements with even atomic numbers tend to have more stable isotopes than elements with odd atomic numbers.
b. Carbon-14 decays with the loss of a β- particle while carbon-10 decays with the loss of a β+.
c. Carbon-14 (t1/2 = 5730 years) can be used to determine the age of organic materials that died between approximately 500
and 50,000 years ago.
8.
(10%) A student is asked to determine the molar mass of an unknown monoprotic carboxylic acid by freezing point depression.
The student dissolves 0.029 g of the unknown acid in 10. mL of water and obtains the cooling curve shown.
o
Temperature, C
0.00
-0.02
-0.04
a. Account for the shape of the curve including the
i. downward slope of the final portion of the curve.
Time, min
ii. depression before the final portion of the curve.
b. Give the value of the freezing point of the solution and calculate the molality of the solute. [Kf = -1.86 ˚C. m-1]
c. Determine the molar mass of the acid.
d. When the student titrates a solution of the acid in water, a molar mass of 120. g.mol -1is determined. Compare these results
with those in part c. and offer an explanation for this behavior.
END OF PART II
Not valid for use as an USNCO National Examination after April 19, 2004.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
N A gas constant
R Planck’s constant
h
°C gram
g pressure
P
c heat capacity
C p rate constant
k
C hour
h retention factor
Rf
E joule
J second
s
Ea kelvin
K speed of light
c
H kilo- prefix
k temperature, K
T
S liter
L time
t
volt
V
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
 k2  Ea  1 1 
=  − 
 k1  R  T1 T2 
– ∆H   1 
ln K = 
+c
 R  T
RT
E = Eο –
ln Q
nF
ln
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(277)
Page 4
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 19, 2004.
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART II
1.
g
0.1500 g Cu
1 mol
×
×
= 1.099 M
L 1.000 g solution 159.62 g
mol
× 0.025 L = 0.0275 mol CuSO 4
b. 1.099
L
a. 1.1169
c. i.
mol
× 0.03063 L = 0.003173 mol HCl = 0.003173 mol NH 3
L
g
0.003173 mol NH 3 × 17.034
= 0.05405 g NH 3
mol
0.1036
% NH 3 =
ii.
0.05405 g NH 3
× 100 = 27.7% NH 3
195 g sample
mol
= 6.11 × 10−4 mol EDTA ≡ 6.11 × 10−4 mol Cu 2+
L
g
6.11 × 10−4 mol Cu 2+ × 63.55
= 0.03881 g EDTA
mol
0.03881 g EDTA
% EDTA =
× 100 = 25.9% EDTA
150 g sample
0.02443 L EDTA × 0.0250
0.200 g compound - 0.185 g anhydrous compound = 0.015 g H 2O
0.0150 g H 2O
% H 2O =
× 100 = 7.5% H 2O
0.200 g sample
% SO 4 = 100 - [ 27.7 + 25.88 + 7.5] = 38.92 %
iv.
d. Assume 100 g; calculate moles using molar mass; divide all results by the smallest number
1 mol
27.7 g NH 3 ×
= 1.627 / 0.405 = 4.02
17.03 g
iii.
1 mol
= 0.407 / 0.405 = 1.00
63.55 g
1 mol
7.5 g H 2O ×
= 0.416 / 0.405 = 1.03
18.02 g
1 mol
38.9 g SO 4 ×
= 0.405 / 0.405 = 1
96.1 g
25.88 g Cu ×
Based on these results, we can see the molecular formula is, Cu(NH3)4 SO4.H2 O
whose molar mass is 245.28 g/mol
e. 245.28
g
× 0.275 mol = 6.76 g is the theoretical yield
mol
% yield =
6.127 g
× 100 = 90.7%
6.76 g
2.
o
o
o
o
o
a. ∆H rxn = ∆H f (N 2O 5 ) + ∆H f (O 2 ) - [2∆H f (NO 2 ) + ∆H f (O 3 )]
−198 kJ = 11 kJ + 0 - [2∆H of (NO 2 ) + 143 kJ so, ∆H of (NO 2 ) = 33 kJ
b. ∆So < 0 because 3 moles of gas are converted into 2 moles of gas.
c. ∆Go = ∆Ho – T∆So = -198000 J – 298 K(-168 J/K) = -148 kJ
Not valid for use as an USNCO National Examination after April 19, 2004.
Page 5
d. Spontaneity will decrease as temperature increases because the entropy change of the process is
negative. This makes the contribution of the T∆So term to the free energy positive. As temperature
rises, this term contributes more to the free energy and because it is positive, the spontaneity must
decrease.
e. Looking at the table we see from comparing experiments 1&2 that when [NO 2 ] increases by a factor
of 1.5 the rate of reaction also increases by 1.5, so it is first order in [NO2]. From experiments 2&3
we see that an increase in the [O 3 ] by a factor of 2.0 increases the rate of reaction by 2.0 so the
reaction is also first order in [O 3 ].
f. The rate law is, rate = k[NO 2 ][O 3 ]
4.8 × 10−8 M −1s−1 = k (0.0015 M )(0.0025 M ) so k = 0.0128 L mol-1 s-1
g. Mechanism #1 would suggest a rate law of rate = k[NO 2 ]2 and mechanism #2 would suggest a rate law
of, rate = k[NO 2 ] [O 3 ] . Neither of these mechanisms is consistent with the observed rate law.
h. The rate of the reaction would be expected to increase with an increase in temperature. Collision rate
increases as temperature increases and a larger fraction of collisions have the energy necessary for
reaction (Ea).
2
3.
[Ag + ]2 1.4 × 10−5
2=
= 0.56 M
a. K sp = [Ag + ]2[SO 24 ] so [SO 4 ] =
2
K sp
(0.005)
b. Cl– then Br– then I– .
For Cl, 1.8 × 10−10 = [Ag + ](0.10), so [Ag + ] = 1.8 × 10−9 ,
for 5 × 10−13 = [Ag + ](0.10), so [Ag + ] = 5 × 10−13
and for 8.3 × 10−17 = [Ag + ](0.10), so [Ag + ] = 8.3 × 10−17
The [Ag+] in equilibrium with AgCl in 0.1 M Cl– is sufficient to cause precipitation of AgBr in
0.1 M Br– . In turn, the [Ag+] in equilibrium with AgBr in 0.1 M Br– will cause precipitation of AgI
in 0.1 M I– . However, the reverse order of addition of the anions would not lead to this behavior.
c. We need the net equilibrium constant for the combined reaction in each case.
AgCl(s) → Ag + (aq) + Cl− (aq) K sp = 1.8 × 10−10
Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7
So K = K sp × K f = 3.06 × 10−2
and
AgBr(s) → Ag + (aq) + Br − (aq) K sp = 5 × 10−12
Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7
So K = K sp × K f = 8.5 × 10−7
and
AgI(s) → Ag + (aq) + I− (aq) K sp = 8.3 × 10−17
Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7
So K = K sp × K f = 1.41 × 10−9
Now calculate the reaction quotient, Q, for the given case (using chloride as the example),
AgCl(s) + 2NH 3 (aq) → Ag(NH 3 ) +2 + Cl− (aq)
So, Q =
[
[Cl– ] Ag(NH 3 ) 2+
[ NH 3 ]
2
] = (0.10)(0.005) = 1.39 × 10-3
(0.60) 2
This value is smaller than only one of the calculated values for the net equilibrium constant (the case of
chloride) so the silver chloride is the only one that will dissolve in 0.60 M NH3.
Page 6
Not valid for use as an USNCO National Examination after April 19, 2004.
4.
a. Al-O bonds are too strong to be broken by simple heating of the oxide, even in the presence of
carbon. The heat of formation of Al 2 O3 is much more negative than that of CO2, so the reaction:
2Al2O 3 + 3C → 4 Al + 3CO 2 is endothermic.
b. Na3 AlF6 is added to lower the melting point of Al2O3 . Lowering the melting point also lowers the
amount of energy needed to carry out the process.
2–
–
3+
–
c. Al + 3e → Al (cathode) and 2O → O 2 + 4e
3
d. 5.00 × 10 g Al ×
1 mol
3 mol e –
= 185.3 mol Al and 185.3 mol Al ×
= 556 mol e –
26.98 g
1 mol Al
e. We need to determine the charge (in coulomb) and time (in seconds) to calculate current.
96500 C
60 min 60 sec
= 5.365 × 10 7 C and 10.0 hours ×
×
= 3.6 × 10 4 sec
–
1 hour 1 min
1 mol e
(5.365 × 10 7 C)
= 1490A
so,
(3.6 × 10 4 sec)
556 mol e – ×
f.
556 mol e – ×
1 mol O 2
= 139 mol O 2
4 mol e –
nRT (139 mol O 2 )(0.0821 L ⋅ atm ⋅ mol-1 ⋅ K -1)(298 K)
V
=
=
= 3590 L
Now use Ideal Gas Law,
 1 atm 
P
720 mmHg × 

 760 mmHg 
5.
a.
b.
c.
d.
H + + NO -3 + FeS → NO 2 + Fe 3+ + S + H 2O
HC 2 H 3O 2 + Ca 3 (PO 4 ) 2 → Ca 2+ + C 2 H 3O –2 + HPO 2–
4
–
–
Al + H 2O + OH → Al(OH) 4 + H 2
2+
Cr 3+ + SO 2–
+ OH – → Cr(OH) 3 + BaSO 4
4 + Ba
2+
–
2–
e. Co + Cl → CoCl4
f. Br2 + CH 3CHCH 2 → CH 3CHBrCH 2 Br
6.
a. The Mg2+ ion is smaller and has a higher charge than the Mg+ ion, so the lattice energy that arises when
Mg2+ ions form compounds is much greater than what would be observed if Mg+ ions formed
compounds. The increase in lattice energy more than offsets the larger ionization energy of the Mg2+ ion.
b. Ti (atomic number 24) has a valence electron configuration of 4s2 3d2 and can form +2 ions by losing it’s
two 4s electrons, +3 ions by losing the two 4s electrons and one 3d electron and +4 by losing all four of
the valence electrons. The +2 and +3 ions are colored because of electronic transitions between d
orbitals. The +4 ion does not exhibit color because it has no valence d-electrons to undergo electronic
transitions.
c. Carbon dioxide (O=C=O) molecules are nonpolar and interact with each other only through weak
dispersion forces. These weak forces are easily overcome so CO2 is a gas at room temperature. SiO2
doesn’t have the same molecular formula, because Si does not form double bonds as readily as carbon
does. Si-O form single bonds that lead to a network solid held together with strong, covalent bonds, so it
is a solid that has a high melting point.
d. Nitrogen can form three bonds (NF 3 ) but not five (NF5 ) because it lacks d orbitals that are energetically
available for the formation of hybrid orbitals (or alternatively, because it is too small to accommodate
five atoms.) Both NF3 and PF3 are trigonal pyramidal because the central atom has three bonding pairs
and one lone pair of electrons (leading to sp 3 hybridization). PF5 is trigonal bipyramidal because it has
five bonding pairs (leading to dsp3 hybridization.)
Not valid for use as an USNCO National Examination after April 19, 2004.
Page 7
7.
a. Even numbers of protons can pair up, which makes nuclei more stable. These more stable nuclei can
accommodate a wider range of n/p ratios.
b. Light nuclei are more stable when the n/p is close to 1/1. C-14 loses a β– because it’s n/p ratio of 8/6
is too high, so a neutron is converted to a proton, giving a n/p ratio of 7/7. C-10 loses a β+, thereby
converting a proton to a neutron and changing its n/p ration from 4/6 to 5/5.
c. The difference in radioactivity between a fresh sample of C-14 containing material and a sample
between 500 years and 50,000 years old can be used to determine the age of the historical sample.
Samples less than 500 years old produce too little difference from new samples to provide reliable
dates. Samples older than 50,000 years contain too little C-14 to provide sufficient radioactivity
(“counts” measured) to give reliable dates.
8.
a. i. As the solution cools, the water freezes, leaving a more concentrated solution. The more
concentrated the solution, the lower the freezing temperature, so the line slopes downward once the
solution starts to freeze.
ii. The depression of the freezing point is due to supercooling of the solution before crystallization
begins.
b. The freezing point is approximately –0.026 o C. ∆T = k f m so m =
∆T −0.026
=
= 0.014 m
kf
1.86
c. The molality is 0.014 mol / kg water, so the molar mass can be calculated,

mol 
0.029 g
−4
= 207 g/mol
 0.014
 × (0.01 kg) = 1.4 × 10 mol and the molar mass is
kg 

1.4 × 10−4 mol
d. The molar mass determination by freezing point depression is roughly twice as great as that
determined by titration. Carboxylic acids tend to dimerize in solution so the freezing point depression
experiment, which observes the number of particles rather than their chemical properties,
overestimates the molar mass by almost a factor of two (dimers take two particles and combine them
into one.)
Page 8
Not valid for use as an USNCO National Examination after April 19, 2004.
2004 U.S. NATIONAL
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART III
Problem 1. You have been given a vial containing either maleic acid, C4 H4 O4 , fumaric acid, C4H4 O4 , or
tartaric acid, C4 H6 O6 . Your lab instructor will identify the acid you have been given.
Design and carry out an experiment with the materials provided to determine the number of ionizable H+
ions possible for each molecule of the acid given.
This is a solid acid titration experiment. Students had to determine the number of moles of NaOH base present
given a 0.25M solution, then perform a titration against a weighed sample of the solid maleic, fumaric, or
tartaric acid. The endpoint is detected using phenolphthalein as the indicator. All of these solid organic acids
have two ionizable protons.
Experimental Plan:
An experiment plan had to describe the steps needed to carry out a successful, microscale titration.
Data and Observations:
The recording of data and observations needed to include information about the mass of the acid used and the
volume of the base used. Where the quantity was obtained by difference, both observations used to determine
the amount should be shown. Significant figures had to be used appropriately.
Sample Calculations and Conclusions:
Students needed to provide calculations that supported their conclusions.
An example of excellent student work:
Plan:
Dissolve a pre-weighed acid sample in distilled water adding two drops phenolphthalein. Fill the 10.0 ml
graduated cylinder with .25M NaOH. Add base gradually until a permanent pink color. Record total volume of
base used. Repeat.
Data:
Acid used: Maleic
MM=116.0 g/mol
Mass of vial + acid
Mass of vial less acid
Mass of acid
Volume base
Key for 2004 USNCO National Exam, Part III
Trial One
4.50g
4.00 g
.50g
34.4 ml
Trial Two
4.00 g
3.48g
.52g
35.0 ml
Calculations:
Moles acid = mass acid/MM
Moles base= (liters base x Mbase)
Mole base/moles acid
.0043
.0086
2.0
.0045
.0088
2.0
Conclusion:
For maleic acid, there are two moles of ionizable protons/mol acid. It is diprotic.
An example of good student work:
Plan:
Dissolve maleic acid in water. Add phenolphthalein and titrate with .25M NaOH. Calculate.
Data:
1.462g maleic in 125 ml solution. Molarity solution = .101
Use 25 ml of solution and titrate. 19.5 ml NaOH used.
Calculations:
Mol Maleic = .0025
Mol NaOH = 19.5ml/1000ml/L(.25mol/L) = .004875
Mol NaOH/Mol Maleic = 2
Conclusion:
Maleic acid is diprotic.
Key for 2004 USNCO National Exam, Part III
(Work not shown)
(Poor use sig. fig)
(Only one trial)
Problem 2: You have been given 4 (four) black pens. Design and carry out an experiment to determine
whether the dye used in each pen is a compound or a mixture. You will need to provide evidence to justify
your conclusions.
This is a paper-chromatography identification experiment. Ammonia and water are provided as the carrying
liquids. Students had to create an experiment using pieces of filter paper to observe a possible separation of the
black ink dyes from each of the four pens provided. Situating the filter paper on or in provided beakers
allowed the solvent front to rise, showing a possible separation.
Experimental Plan
The experimental plan needed to identify a way to use paper chromatography to investigate the inks in the
various pens. Some detail about how to carry out the experiment – marking the filter paper above the liquid
level, for example, was useful in this component of the exercise.
Observations and results:
Students needed to summarize observations about the chromatography experiment. Students who did an
excellent job carried out more than one trial to verify results and included detail about separations. A table of
probable observations includes,
Pens
#1 Crayola ®
#2 Gel-Pen ®
#3 Papermate ®
#4 Sharpie ®
Water
Separation
No separation
or movement
Separation
No separation
Ammonia
Separation
No separation
Conclusions
= a mixture
= a compound
Separation
No separation
= a mixture
= a compound
An example of excellent student work:
Using Ammonia as solvent:
- Pen 1: Different colors appear: orange, blue, green and purple
- Pen 2: Solid black color only appears.
- Pen 2 (again): Still shows only solid black color.
- Pen 3: Different colors appear, but they are light: green, blue and purple
- Pen 3 (again): Different colors appear again, darker than previous run.
- Pen 4: Mark does not move, only black color appears.
- Pen 4 (again): Mark does not move, only black color.
Using water as solvent – retry #2 and #4.
- Pen 2: Solid black color only appears
- Pen 4: No movement observed.
- Pen 4 (again): No movement observed.
Conclusions:
Pen 1 is a mixture because different colored pigments are observed.
Pen 2 is a compound because only one color is observed even though the ink moved along the filter paper.
Pen 3 is a mixture because different colored pigments are observed. The pigments used are different than
those used in Pen 1.
Key for 2004 USNCO National Exam, Part III
Pen 4 can not be determined in this experiment because no solvent was found that dissolved the ink. Based on
the fact that separation of color is not observed, the ink might be a single compound.
An example of good student work:
Pen 1: Mulitple layers form. Fastest moving layer is light blue green. A darker blue-green is next and a
purple color is the slowest.
Pen 2: One dark black layer that doesn’t seem to move at all. One light black layer that moves slowly. One
gray layer that moves faster.
Pen 3: One black layer doesn’t move much. One light black layer that moves slowly. One blue-purple layer
that doesn’t move at all.
Pen 4: Black mark doesn’t move at all and remains dark black.
Conclusion:
The pens with multiple layers in the paper chromatography were mixtures – in this case pens 1, 2 and 3. The
pen with a single layer in the paper chromatography, pen 4, was a single compound.
Key for 2004 USNCO National Exam, Part III
2005 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM PART 1
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
Alice Johnsen, Bellaire High School, TX
William Bond, Snohomish High School, WA
Adele Mouakad, St. John’s School, PR
Peter E. Demmin (retired), Amherst Central High School, NY
Kimberley Gardner, United States Air Force Academy, CO,
David W. Hostage, Taft School, CT
Jane Nagurney, Scranton Preparatory School, PA
Ronald O. Ragsdale, University of Utah, UT
Jacqueline Simms, Sandalwood Sr. High School, FL
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 27, 2005, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
ABBREVIATIONS AND SYMBOLS
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G molar mass
A frequency
ν mole
N A gas constant
R Planck’s constant
°C gram
g pressure
c heat capacity
C p rate constant
C hour
h retention factor
E joule
J second
Ea kelvin
K temperature, K
H kilo– prefix
k time
S
liter
L volt
K
milli– prefix
m
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
T
t
V
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
0 °C = 273.15 K
1 atm = 760 mmHg
EQUATIONS
E = Eo −
1
1A
1
H
1.008
3
Li
RT
ln Q
nF
k  E  1 1 
ln 2  = a  − 
 k1  R  T1 T2 
 −ΔH  1 
ln K = 
  + constant
 R  T 
PERIODIC TABLE OF THE ELEMENTS
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
18
8A
2
He
4.003
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
19
K
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
26.98
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
(98)
44
Ru
101.1
45
Rh
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
85.47
87.62
88.91
91.22
92.91
95.94
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(269)
(272)
(277)
(2??)
(223)
(226)
(227)
€
58
Ce
59
Pr
(262)
60
Nd
(263)
61
Pm
(262)
62
Sm
(265)
63
Eu
(266)
64
Gd
65
Tb
66
Dy
67
Ho
173.0
175.0
101
Md
102
No
103
Lr
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
70
Yb
(209)
168.9
140.9
238.0
209.0
69
Tm
90
Th
231.0
207.2
68
Er
140.1
232.0
Page 2
(261)
204.4
(259)
(210)
(222)
71
Lu
(262)
Not valid as a USNCO National Exam after April 26, 2005
DIRECTIONS
 When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
 There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
 Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which solution produces a black precipitate when added
to an aqueous copper(II) solution?
(A) NH3
(B) (NH4)2S
(C) K2SO4
(D) NaOH
6. Which diagram best represents the change in electrical
conductivity of a solution of acetic acid as a solution of
sodium hydroxide is added?
(A)
(B)
(C)
(D)
2. Which oxide is the best reducing agent?
(A) CO2
(B) NO2
(C) SiO 2
(D) SO2
3. Solutions of which ion produce a red color when
vaporized in a Bunsen burner flame?
(A) calcium
(B) potassium
(C) sodium
(D) zinc
4. Which procedure for dispensing a liquid with a
volumetric pipet is correct?
(A) Draw the liquid up to the line on the pipet using a
pipet bulb. Squeeze the bulb to force all the liquid in
the pipet into the receiving container.
(B) Introduce the liquid into the top end of the pipet
until it is filled to the line. Allow the liquid to drain
into the desired container. Blow on the pipet to
release the last drop.
7. Methylamine, CH3NH2, reacts with O2 to form CO 2, N2,
and H2O. What amount of O2 (in moles) is required to
react completely with 1.00 mol of CH3NH2?
(A) 2.25
(B) 2.50
(C) 3.00
(D) 4.50
(C) Draw the liquid above the line on the pipet using a
pipet bulb. With a finger on the top of the pipet
allow the curve of the meniscus to drop to the line.
Place the tip of the pipet against the side of the
receiving container and allow the liquid to drain.
8. Iodine adds to the double bonds in fatty acids (one iodine
molecule per double bond). How many double bonds are
in a molecule of arachidonic acid (Molar mass = 304.5
g/mol) if 0.125 g of the acid require 0.417 g of iodine?
(D) Draw the liquid above the line on the pipet by
sucking on the open end of the pipet. Place a thumb
on the top of the pipet and allow the curve of the
meniscus to drop to the line. Allow the liquid to
drain into the receiving container pipet against its
side.
9. The solubility of a gas in a I. pressure of the gas
liquid increases when
II. temperature of the liquid
which of the following increases?
5. Which physical characteristic distinguishes copper from
brass (an alloy of copper and zinc)?
(A) Brass is a liquid at room temperature and copper is
not.
(A) 2
(B) 3
(C) 4
(D) 8
(A) I only
(B) II only
(C) both I and II
(D) neither I nor II
10. A mineral containing only manganese and oxygen
contains 69.6% Mn by mass. What is its empirical
formula?
(B) Brass is much less dense than copper.
(A) MnO
(B) Mn2O3
(C) Brass is attracted to a magnet but copper is not.
(C) Mn3O4
(D) MnO2
(D) Brass is a much poorer electrical conductor than
copper.
Not valid as a USNCO National Exam after April 26, 2005
Page 3
11. Toluene, C7H8, is added to gasoline to increase its octane
rating. What is the volume ratio of air to toluene vapor to
burn completely to form CO2 and H 2O? (Assume air is
20% O2 by volume.)
(A) 9/1
(B) 11/1
(C) 28/1
(D) 45/1
12. Acidified solutions of dichromate ion, Cr2O72-, oxidize
Fe2+ to Fe3+, forming Cr3+ in the process. What volume of
0.175 M K 2Cr2O7 in mL is required to oxidize 60.0 mL of
0.250 M FeSO4?
(A) 14.3
(B) 28.6
(C) 42.9
(D) 85.7
13. Which property is the same for 1.0 g samples of H2 and
CH4 in separate 1.0 L containers at 25 ˚C?
(A) pressure
(B) number of molecules
(C) average molecular velocity
(D) average molecular kinetic energy
14. When CsI, SiO2, CH3OH and C 3H8 are listed in order of
increasing melting point, which is the correct order?
(A) CsI, SiO2, CH3OH, C3H8
(B) CH3OH, C3H8, CsI, SiO2
17. When NaF, MgO, KCl and CaS are listed in order of
increasing lattice energy, which order is correct?
(A) MgO, NaF, KCl, CaS
(B) CaS, MgO, KCl, NaF
(C) KCl, CaS, NaF, MgO
(D) KCl, NaF, CaS, MgO
18. When compared to most
I.
boiling point
other substances of similar II. specific heat capacity
molar mass the values of
III. surface tension
which properties of liquid H2O are unusually large?
(A) I only
(B) I and II only
(C) II and III only
(D) I, II and III
19. Calculate ∆H˚ for the reaction;
∆Hf˚ kJ/mol
TiCl4(g) + 2H 2O(l)
TiCl4(g)
–763
r TiO 2(s) + 4HCl(g) H2O(l)
–286
TiO2(s)
–945
HCl(g)
–92
(A) –264 kJ
(B) 12 kJ
(C) 22 kJ
(D) 298 kJ
20. Use bond energies to estimate the
value of ∆H˚ for the reaction;
N2(g) + 3H 2(g) r 2NH 3(g)
(C) CH3OH, C3H8, SiO 2, CsI
(D) C 3H8, CH3OH, CsI, SiO2
15. According to the graph
(ln vapor pressure vs
1/T) what can be
concluded about the
enthalpies of
vaporization (∆Hvap ) of
liquids X and Y?
Bond Energies
kJ/mol
H-H
436
H-N
386
N-N
193
N=N
418
N≡N
941
(A) –995 kJ
(B) –590 kJ
(C) –67 kJ
(D) 815 kJ
Questions 21. and 22. should be answered using this
thermochemical equation;
N2(g) + 2O2(g) r 2NO2(g) ∆Hrxn > 0
21. Which relationship is correct for this reaction at a
pressure of 1 atm?
(A) ∆Hvap X > ∆Hvap Y
(B) ∆Hvap X = ∆Hvap Y
(C) ∆Hvap X < ∆Hvap Y
(D) No conclusions can be drawn about the relative
∆Hvap values from this diagram.
16. An unknown gas effuses through a pin-hole in a container
at a rate of 7.2 mmol/s. Under the same conditions
gaseous oxygen effuses at a rate of 5.1 mmol/s. What is
the molar mass (in g/mol) of the unknown gas?
(A) 16
Page 4
(B) 23
(C) 45
(D) 64
(A) ∆Erxn > ∆Hrxn
(B) ∆Erxn < ∆Hrxn
(C) ∆Erxn = ∆Hrxn + ∆S rxn
(D) ∆Erxn = ∆Hrxn – ∆S rxn
22. Under what temperature conditions is this reaction
spontaneous at standard pressure?
(A) at low temperatures only
(B) at high temperatures only
(C) at all temperatures
(D) at no temperature
Not valid as a USNCO National Exam after April 26, 2005
23. Diethyl ether has a normal boiling point of 35.0 ˚C and
has an entropy of vaporization of 84.4 J/mol. K. What is
its enthalpy of vaporization?
(A) 0.274 J/mol
(B) 2.41 J/mol
(C) 3.65 J/mol
(D) 26.0 kJ/mol
24. A 9.40 g sample of
Solution Properties
KBr is dissolved in 105 Molar mass KBr 119 g/mol
g of H 2O at 23.6 ˚C in a ∆Hsoln KBr
19.9 kJ/mol
coffee cup. Find the
C p solution
4.184 J/g˚C
final temperature of this system. Assume that no heat is
transferred to the cup or the surroundings.
(A) 20.0 ˚C
(B) 20.3 ˚C
(C) 26.9 ˚C
(D) 27.2 ˚C
25. For the reaction A r B which is
first order in A, which of the
following change as the
concentration of A changes?
I.
II.
III.
rate
rate constant
Half–life
(A) I only
(B) III only
(C) II and III only
(D) I, II and III
26. The equation and rate law for the gas phase reaction
between NO and H2 are;
2NO(g) + 2H2(g) r N 2(g) + 2H 2O(g)
Rate = k[NO]2[H 2]
What are the units of k if time is in seconds and the
concentration is in moles per liter?
(A) L. s. mol -1
(B) L2. mol -2. s-1
(C) mol . L-1. s-1
(D) mol 2. L-2. s-1
27. At a given temperature a first-order reaction has a rate
constant of 3.33×10-3 s-1. How much time is required for
the reaction to be 75% complete?
(A) 100 s
(B) 210 s
(C) 420 s
(D) 630 s
28. Most reactions occur more
I.
activation energy
rapidly at high temperatures
II.
collision energy
than at low temperatures. This III. rate constant
is consistent with an increase in which property at higher
temperatures?
(A) I only
(B) II only
(C) I and III only
(D) II and III only
29. Which graph is diagnostic of an irreversible second order
reaction A r B?
(A)
(B)
(C)
(D)
30. The reaction; 2NO(g) + 2H2(g) r 2H2O(g) + N2(g)
obeys the rate equation
Rate = k[NO]2[H 2]
This mechanism has been proposed:
1.
2NO(g) r N2O2(g)
2.
N2O2(g) + H 2(g) r 2HON(g)
3.
HON(g) + H2(g) r H 2O(g) + HN(g)
4.
HN(g) + HON(g) r N2(g) + H 2O(g)
Which step of the mechanism is the rate-determining
step?
(A) step 1
(B) step 2
(C) step 3
(D) step 4
31. For the hypothetical equilibrium reactions;
AsB
K = 2.0
BsC
K = 0.010
What is the value of K for the reaction;
2C s 2A?
(A) 2500
32. For which
reaction is
Kp = Kc ?
(C) 25
(D) 4.0×10-4
2N2(g) + O 2(g) s 2N2O(g)
I.
II. C(s) + O2(g) s CO2(g)
III. N2O4(g) s 2NO2(g)
(A) II only
(B) III only
(C) I and III only
(D) II and III only
33. What is the pH of a 0.010 M solution of a weak acid HA
that is 4.0% ionized?
(A) 0.60
Not valid as a USNCO National Exam after April 26, 2005
(B) 50
(B) 0.80
(C) 2.80
(D) 3.40
Page 5
34. Given the acid ionization
constants, when the
conjugate bases are
arranged in order of
increasing base strength,
which order is correct?
–
Acid Ionization Constant, Ka
HClO
3.5×10-8
HClO2
1.2×10-2
HCN
6.2×10-10
–
H2PO4
6.2×10-8
–
–
–
–
(C) CN , HPO42–, ClO , ClO 2
–
–
–
(D) CN , ClO , HPO42–, ClO 2
Base Ionization Constant, Kb
35. Calculate the
NH3
1.8×10-5
concentration of
hydrogen ion in mol/L of a 0.010 M solution of NH4Cl.
(A) 4.2×10-4
(B) 2.4×10-6
(C) 1.8×10-7
(D) 5.6×10-12
36. For the reaction;
–
PbI2(s) s Pb2+(aq) + 2I (aq)
Ksp = 8.4×10-9
2+
What is the concentration of– Pb in mol/L in a saturated
solution of PbI2 in which [I ] = 0.01 M?
(A) 8.4×10
-7
(B) 8.4×10
(C) 1.3×10
-3
(D) 2.0×10-3
-5
37. Which statement is correct about the
electrochemical cell
–
represented here? Ag | Ag + || NO3 , NO | Pt
(C) 1033
(D) 1078
(B) The major purpose of the Pt is to act as a catalyst.
(C) The Ag electrode decreases in mass as the cell
operates.
(A) I only
(B) II only
(C) both I and II
(D) neither I nor II
(A) 0.355 V
(B) 0.178 V
(C) –0.178 V
(D) –0.355 V
(C) II and III only
(D) II and IV only
(B) 0.22
(C) 0.33
(D) 0.66
43. How many orbitals are in an atomic sublevel with l = 3?
(A) 3
(B) 5
(C) 7
(D) 9
44. A ground state gaseous atom of which element has the
greatest number of unpaired electrons?
(A) As
(B) Br
(C) Ge
(D) Se
45. An atom of which element has the highest second
ionization energy?
(A) Na
(B) Mg
46. Which of these properties
increase across the period
from Na to Cl?
(C) Al
I.
II.
III.
(D) K
atomic radius
density
electronegativity
(A) I only
(B) III only
(C) I and II only
(D) II and III only
47. For the elements in group 14 (C to Pb), which property
increases with increasing atomic number?
(A) melting points
(B) covalent radius
(C) magnitude of stable oxidation state
(D) ability to form chains of atoms with themselves
48. What mode of radioactive decay is most likely for the
isotope 20
11 Na ?
+
39. The standard reduction potential for H (aq) is 0.00 V.
What is the reduction potential for a 1×10-3 M HCl
solution?
(B) I and IV only
(A) 0.16
(D) The voltage of the cell can be increased by doubling
the size of the Ag electrode.
38. The overall reaction for the lead storage battery when it
discharges is;
+
2Pb(s) + PbO 2(s) + 4H (aq) + 2SO4 (aq)
r 2PbSO 4(s) + 2H2O(l)
I.
PbSO4 is formed only at the cathode.
II. The density of the solution decreases.
Which statement(s) correctly describe(s) the battery as it
discharges?
(A) I and III only
42. A current of 0.20 amps is passed through an aqueous
solution of nickel(II) nitrate for 45.0 minutes. What mass
of Ni metal (in grams) will be deposited?
(A) NO undergoes oxidation at the anode.
Page 6
(B) 1026
41. Which products are formed by the electrolysis of an
aqueous solution of AlCl3?
I. Al(s)
II. Cl2(g) III. H2(g) IV. O2(g)
–
(B) ClO 2 , HPO42–, ClO , CN
–
(A) 1022
–
(A) ClO 2 , ClO , HPO42–, CN
–
Standard Reduction Potential, V
40. What is the
–
approximate value
Ag+(aq) + e r Ag(s) +0.80
–
of the equilibrium
Cr3+(aq) + 3e r Cr(s) –0.74
constant, Keq , at 25 oC for the reaction;
3Ag+(aq) + Cr(s) r Cr3+(aq) + 3Ag(s)
€
(A) alpha
(B) beta
(C) gamma
(D) electron capture
Not valid as a USNCO National Exam after April 26, 2005
49. Oxygen gas is paramagnetic. This observation is best
explained by
56. What is the most characteristic reaction of benzene?
(A) resonance.
(B) the Lewis structure of O2.
(D) the hybridization of atomic orbitals in O2.
50. What is the geometry of the iodine atoms in the I3- ion?
(A) bent
(B) linear
(C) T-shaped
(D) triangular
O
S
O
O
(A) 0, 0
(B) –2, 0
(C) +2, –1
(B) 2
(B) CH3COOH
(C) ClCH2COOH
(D) ClCH2CH2COOH
(B) two
(C) three
(D) four
(A) ClHC = CHCl
(B) meta-C6H4Cl2
(C) CH2ClBr
(D) CH3CH(Cl)CH2CH3
60. Which type of dietary fat is currently considered the least
harmful?
(A) monounsaturated fat
(B) polyunsaturated fat
(C) saturated fat
(D) trans fat
(D) +6, –2
53. How many different isomers
exist for the octahedral
+
complex [Co(NH3)4Cl2] ?
(A) 1
(A) HCOOH
59. Which compound can exist as two optical isomers?
(D) SF6
2–
O
(D) substitution
(A) one
(C) SbF5
52. In the Lewis structure what are
the formal charges on the sulfur
and oxygen atoms, respectively?
(C) reduction
58. How many structurally isomeric alcohols have the
formula C4H9OH?
51. Which species has a dipole moment other than zero?
(B) CF4
(B) polymerization
57. Which organic acid is the strongest?
(C) the molecular orbital description of O2.
(A) BrF 3
(A) addition
(C) 3
END OF TEST
(D) 4
54. Which order is correct when the species are arranged in
order of increasing average N-O bond length?
–
–
+
(B) NO , NO3 , NO2
–
–
+
(D) NO , NO2 , NO3
(A) NO3 , NO2 , NO
(C) NO2 , NO3 , NO
+
–
–
+
–
–
55. All of the classes of compounds contain at least one
oxygen atom EXCEPT
(A) esters
(B) aldehydes
(C) ethers
(D) alkynes
Not valid as a USNCO National Exam after April 26, 2005
Page 7
NATIONAL OLYMPIAD PART I
2005
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
B
D
A
C
D
D
A
C
A
B
D
A
D
D
A
A
D
D
C
C
A
D
D
B
A
B
C
D
C
B
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Property of the ACS DivCHED Examinations Institute
Answer
A
A
D
B
B
B
C
B
C
D
C
A
C
A
A
B
B
D
C
B
A
C
B
D
D
D
C
D
D
B
2005 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
Alice Johnsen, Bellaire High School, TX
William Bond, Snohomish High School, WA
Adele Mouakad, St. John’s School, PR
Peter E. Demmin (retired), Amherst Central High School, NY
Jane Nagurney, Scranton Preparatory School, PA
Kimberley Gardner, United States Air Force Academy, CO,
David W. Hostage, Taft School, CT
Ronald O. Ragsdale, University of Utah, UT
Jacqueline Simms, Sandalwood Sr. High School, FL
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for
a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2005, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the
front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to
use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper,
and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
Page 1
1. (15%) This question involves several reactions of copper and its compounds.
a. A sample of copper metal is dissolved in 6 M nitric acid contained in a round bottom flask. This reaction yields a blue
solution and emits a colorless gas which is found to be nitric oxide. Write a balanced equation for this reaction.
b. The water is evaporated from the blue solution to leave a blue solid. When the blue solid is heated further, a second reaction
occurs. This reaction produces a mixture of nitrogen dioxide gas, oxygen gas and a black oxide of copper.
i.
A sample of the dried gas, collected in a 125 mL flask at 35 ˚C and 725 mm Hg, weighed 0.205 g. Find the
average molar mass of the gas and the molar NO2/O2 ratio in it.
ii.
These data were obtained for the black solid;
Mass of empty flask:
39.49 g
Mass of flask + copper metal:
40.86 g
Mass of flask + oxide of copper:
41.21 g
Determine the formula of the oxide of copper.
c. If some of the blue solution were lost due to splattering during the evaporation, what would be the effect on the calculated
percentage of copper in the black oxide? Explain.
d. If all of the blue solid were not decomposed into the black oxide during the final heating, what would be the effect on the
calculated percentage of copper in the oxide? Explain.
2.
3.
(16%) Liquid hydrazine, N2H4, is sometimes used as a rocket propellant.
a. Write an equation for the formation of hydrazine from its elements and use the combustion equations below to derive an
equation in which ∆Hf˚ for liquid hydrazine is expressed in terms of ∆H1, ∆H2 and ∆H3.
1 N (g) + O (g) r NO (g)
∆H1
2
2
2 2
H2(g) + 1 2 O2(g) r H 2O(g)
∆H2
N2H4(l) + 3O2(g) r 2NO2(g) + 2H 2O(g) ∆H3
b. In€a rocket, liquid hydrazine is reacted with liquid hydrogen peroxide to produce nitrogen and water vapor. Write a balanced
equation for this reaction.
€
c. Calculate ∆Hrxn˚ for the reaction represented in 2b.
∆Hf˚ kJ/mol
N2H4(l)
50.6
H2O2(l)
–187.8
H2O(g)
–285.8
d. Calculate ∆Hrxn˚ for the reaction in 2b using bond dissociation energies.
Bond Dissociation Energy kJ/mol
N–N
167
O–O
142
N=N
418
O=O
494
N≡N
942
O–H
459
N-H
386
e. Which value of ∆Hrxn˚ (that calculated in part c or part d) is likely to be more accurate? Justify your answer.
f. Calculate the maximum temperature of the combustion gases if all the energy generated in the reaction goes into raising the
temperature of those gases. [The heat capacities of N2(g) and H2O(g) are 29.1 J/(mol . ˚C) and 33.6 J/(mol. ˚C), respectively.]
(13%) A solution of alanine hydrochloride, [H3NCH(CH3)COOH]+Cl-,
is titrated with a solution of sodium hydroxide to produce a
curve similar to the one shown.
a.
Give the formula(s) for the major species present at the points on the titration curve.
i. 1 ii. 2 iii. 3 iv. 4
b. If K1 and K 2 of alanine hydrochloride are 4.6×10–3 and 2.0×10–10 respectively,
i. write equations to represent the reactions responsible for K1 and K 2.
ii. determine the pH at points 1, 2 and 3.
c. Describe quantitatively how you could prepare a buffer solution with a pH = 10.0. Solutions of 0.10 M alanine
hydrochloride and 0.10 M NaOH are available.
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
4.
(10%) An electrochemical cell based on the reaction;
M(s) + Cu2+(aq) r M2+(aq) + Cu(s)
E˚ = 1.52 V
is constructed using equal volumes of solutions with all substances in their standard states.
a. Use the value of the reduction potential of Cu2+(aq) (E˚ = 0.34 V) to determine the standard reduction potential for the
–
reaction; M2+(aq) + 2e r M(s)
b. The cell is allowed to discharge until the [Cu2+] = 0.10 M. Find
i. the M 2+ concentration in moles per liter,
ii. the cell potential, E.
c. 50 mL of distilled water is added to each cell compartment of the original cell. Compare the potential of the cell after the
addition of water with the potential of the original cell. Explain your answer.
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Solid sodium sulfide is added to water.
b. An aqueous solution of potassium triiodide is added to a solution of sodium thiosulfate.
c. Excess aqueous sodium fluoride is added to aqueous iron(III) nitrate.
d. Strontium-90 undergoes beta decay.
e. Excess carbon dioxide is bubbled through a solution of calcium hydroxide.
f. Ethanol is warmed gently with acidified potassium dichromate.
6.
(10%) This mechanism has been proposed for the reaction between chloroform and chlorine.
Step 1: Cl2(g) s 2Cl(g)
fast
Step 2: CHCl3(g) + Cl(g) s CCl 3(g) + HCl(g)
slow
Step 3: CCl3(g) + Cl(g) s CCl 4(g)
fast
a. Write the stoichiometric equation for the overall reaction.
b. Identify any reaction intermediates in this mechanism.
c. Write the rate equation for the rate determining step.
d. Show how the rate equation in c. can be used to obtain the rate law for the overall reaction.
e. If the concentrations of the reactants are doubled, by what ratio does the reaction rate change? Explain.
7.
(14%) Xenon forms several compounds including XeF2, XeF4 and XeO3.
a. Draw a Lewis structure for each of these molecules.
b. Describe the geometry of each compound including bond angles.
c. State and explain whether each molecule is polar or nonpolar.
d. Account for the observation that these compounds are highly reactive.
8.
(10%) A section of the polymer polypropylene is represented here.
a. Sketch a structural formula for the monomer used to make this polymer.
b. State and explain whether this polymer is an addition or a condensation polymer.
c. Compare the melting points for the following polymers. Give reasons for your answers.
i. polypropylene containing 1000 monomer units vs polypropylene containing 10,000 monomer units
ii. polypropylene vs a polymer in which the CH3 group is replaced with a CH2CH2CH2CH3 group
iii. isotactic polypropylene (all the CH 3 groups on the same side of the carbon backbone) vs atactic polypropylene (CH 3
groups arranged at random)
END OF PART II
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
Page 3
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
N A gas constant
R Planck’s constant
h
°C gram
g pressure
P
c heat capacity
C p rate constant
k
C hour
h retention factor
Rf
E joule
J second
s
Ea kelvin
K speed of light
c
H kilo- prefix
k temperature, K
T
S liter
L time
t
volt
V
E = Eο –
USEFUL EQUATIONS
 –ΔH   1 
ln K = 
   +c
 R   T
RT
ln Q
nF
CONSTANTS
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
 k2  Ea  1 1 
=
−
 k1  R  T1 T2 
ln 
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
3
Li
4.003
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
26.98
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
112
116
118
(269)
(272)
(277)
(277)
(289)
(293)
(223)
226.0
227.0
58
Ce
59
Pr
(262)
(262)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
144.2
91
Pa
92
U
238.0
237.0
(244)
63
Eu
(266)
(243)
64
Gd
(247)
65
Tb
(247)
66
Dy
(251)
67
Ho
(252)
68
Er
(257)
69
Tm
(258)
70
Yb
(209)
(145)
140.9
62
Sm
(265)
209.0
61
Pm
90
Th
231.0
(263)
207.2
60
Nd
140.1
232.0
Page 4
(261)
204.4
(259)
(210)
(222)
71
Lu
(260)
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART II
1.
a. 3Cu + 8H+ + 2NO3 – → 3Cu2+ + 2NO + 4H 2 O
b. i.
mRT (0.205 g)(0.0821 L ⋅ atm⋅ mol-1 ⋅ K -1 )(308 K)
MM ave =
=
= 43.5 g ⋅ mol-1
1
atm
PV
(0.125 L)(725 mmHg ×
760 mm Hg)
€
€
For the ratio, let x be the fraction of NO2 ,
xMM NO 2 + (1- x)MM O 2 = 43.5
46x + 32(1- x) = 43.5
solving for x : x = 0.821
NO 2 0.821 4.59
=
=
O2
0.189
1
ii. 40.86 g - 39.49 g = 1.37 g Cu
41.21 g - 40.86 g = 0.35 g O
1 mol Cu
1.36 g Cu ×
= 0.0216 mole Cu
63.55 g Cu
1 mol O
0.35 g O ×
= 0.0219 mole O
16.0 g O
Therefore the ratio is 1 : 1 and the formula must be CuO.
c. The lost solution will cause the mass of CuO to be too low relative to the mass of Cu. Therefore the
percentage determination for copper will be too high
d. The mass of CuO will be too high. Therefore the percentage determination for copper will be too
low.
2.
a. looking at N2 + 2H2 → N 2 H4
N 2 + O 2 → NO 2
we must use the equations,
2H 2 + O 2 → 2H 2O
2NO 2 + 2H 2O → 3O 2 + N 2 H 4
N 2 + 2H 2 → N 2 H 4
ΔH f = 2ΔH1 + 2ΔH 2 – ΔH 3
2ΔH1
2ΔH 2
– ΔH 3
b. N 2 H4 (l) + 2H2 O2 (l) → N2 (g) + 4H2 O(g)
c. The structures of hydrazine
and peroxide are,
€
and
Page 5
Thus,
ΔH rxn = 4 mol × (–285.8 kJ ⋅ mol-1 ) – [2 mol × (–187.8 kJ ⋅ mol-1 ) + 1 mol × (50.6 kJ ⋅ mol-1 )]
= – 1143.2 + 325.0
= – 818.2 kJ
d.
ΔH rxn = N – N + 4 N – H + 2 O – O + 4 O – H – [N ≡ N + 8 O – H ]
€
(note : 4O – H cancel from each side)
= 167 + 4(386) + 2(142) – [942 + 4(459)]
= – 783 kJ
e. ΔH from part c should be more accurate. ΔHf values are determined for each compound individually,
whereas bond energies are average values. We should expect the actual values for the compounds in
this problem to vary from these averages.
€
f. C = 1 mol × (29.1 J⋅mol -1⋅o C-1) + 4 mol × (33.6 J⋅mol -1⋅o C-1) so C=163.5 J⋅o C-1
q = CΔT 818200 J = 163.5 J⋅o C-1 × ΔT so ΔT = 5004 o C.
3.
a.
i.
ii.
iii.
iv.
[H3 NCH(CH3 )COOH]+ and +H3 NCH(CH3 )COO–
+
H3 NCH(CH3 )COO–
+
H3 NCH(CH3 )COO– and H2 NCH(CH3 )COO–
H2 NCH(CH3 )COO–
b. i. K1 [H3 NCH(CH3 )COOH]+ → H + + +H3 NCH(CH3 )COO–
K2 +H3 NCH(CH3 )COO– → H + + H2 NCH(CH3 )COO–
ii. [H +] at point 1 = K1 = 4.6×10-3 so pH = 2.34
[H+] at point 2 = K1 × K 2 = 9.2 ×10−13 = 9.6 ×10−7 so pH = 6.02
[H+] at point 3 = K2 = 2×10-10 so pH = 9.07
[1.0 ×10−10 ][B– ]
[B– ] 2
so
=
[HB]
[HB] 1
Take a specified volume of 0.10 M alanine hydrochloride(such as 300 mL) in which the predominant
species is [H3 NCH(CH3 )COOH]+. Add 0.10 M NaOH so there is 5/3 as much as there is 0.10 M alanine
[H NCH(CH 3 )COO – ] 2
(such as 500 mL). This will€give 800 mL of solution with a ratio of + 2
= .
[ H 3NCH(CH 3 )COO – ] 1
c. 2.0 ×10−10€=
€
a. Ε ocell = Ε oox + Ε ored and Ε oox = 1.52 V – 0.34 V = 1.18 V
so Ε ored (M 2+ + 2e – → M) = –1.18 V
€ 1.90 M
b. i. If Cu2+ decreases to 0.10 M then M2+ must increase to
0.0257 1.90
−
ln
ii. Ε = 1.52 €
= 1.52 – 0.01285 ln(19) = 1.52 – 0.0378 = 1.48 V
€
2
0.10
€ c. The E of the cell with dilute solutions will be the same as the original Eo . Because the solutions are
diluted by the same amount and the ions have the same coefficients (from the balanced chemical
equation), Q in the Nernst equation is 1, and lnQ = 0.
€
4.
Page 6
Not valid for use as an USNCO Olympiad National Exam after April 26, 2005.
5.
a. Na2 S + H2 O → 2Na + + HS– + OH–
b. I3 – + S2 O3 2– → 3I– + S4 O6 2–
c. F– + Fe3+ → FeF 6 3–
90
0
d. 90
38 Sr → 39Y + -1β
e. CO2 + Ca(OH) 2 → Ca2+ + HCO3 –
f. CH3 CH2 OH + Cr2 O7 2– + H+ → CH 3 CH=O (or CH3 COOH) + Cr3+ + H2 O
6.€
a.
b.
c.
d.
7.
a.
Cl2 + CHCl3 → CCl4 + HCl
Cl and CCl3
Rate = k[CHCl3 ][Cl]
Because this step is at equilibrium, we can express the [Cl] in terms of [Cl2 ] by looking at the
[Cl]2
1
1
equilibrium constant expression. K =
so [Cl]2 = K[Cl2 ] and [Cl] = K 2 [Cl2 ] 2 . Thus by
[Cl2 ]
substituting we get the overall expression to be: Rate = k[CHCl3 ][Cl2 ]1/2
e. If [CHCl3 ] and [Cl2 ] are doubled, rate will increase by (2)•(2)1/2 = 2.83 times.
€
and
and
b. XeF 2 is linear, 180o .
XeF4 is square planar, 90o
XeO3 trigonal pyramid, ~107o
c. XeF2 is nonpolar. Both Xe–F bond dipoles are the same size, but due to the linear geometry they offset
each other.
XeF4 is nonpolar. All Xe–F bond dipoles are the same size, but due to the square planar geometry they
offset each other.
XeO3 is polar. The Xe–O bond dipoles are the same size, and the non planar geometry leads to a net
dipole.
d. Xe has a formal positive charge in all of these compounds. This makes them good oxidizing agents.
8.
a.
b. It is an addition polymer. To form, the double bond in a monomer breaks to give a lone electron that
forms bonds to other monomers. No other product(s) are formed, so it cannot be condensation.
c. i. Polypropylene with 10,000 units melts at a higher temperature than one with 1000 units. The larger
molecule, with higher molar mass has stronger dispersion forces.
ii. Replacing CH3 with CH2 CH2 CH2 CH3 will lower the melting temperature. The bulk of this larger
group will impede the packing of the polymer chains and decrease the strength of the intermolecular
forces.
iii. Isotactic polypropylene will melt at a higher temperature than atactic polypropylene. The more
regular structure of the isotactic form allows better packing and stronger intermolecular forces.
Page 7
2005 U.S. NATIONAL
CHEMISTRY OLYMPIAD
KEY FOR NATIONAL EXAM – PART III
Lab Problem 1
You have been given three beakers containing NaHCO3 (sodium hydrogen carbonate), CaCl2
(calcium chloride), and tap water. These two compounds react in the presence of water.
Propose a balanced chemical equation to account for this reaction, and support your proposal
with all possible qualitative and quantitative observations and measurements.
This problem requires the students to carry out a reaction and make qualitative and quantitative
observations.
1. Sample answer:
CaCl2
+
2NaHCO3 → Ca(OH) 2 (s) + 2 CO2 (g) + 2 Na+(aq) + 2 Cl– (aq)
+
2NaHCO3 → CaCO3 (s) + CO2 (g) + H2 O
or
CaCl2
+
2 Na+(aq) + 2 Cl– (aq)
Weigh 1 mmol CaCl2 and add to one corner of the plastic bag, then weigh 2 mmol NaHCO3 and
add to the other corner of the plastic bag. Add water to the NaHCO 3 corner without mixing, seal
the bag, and weigh. Allow the reagents to mix, observing for bubbling and inflation of the bag
(to verify CO 2 production) and formation of a white insoluble solid (to verify either Ca(OH)2 or
CaCO3 ). Carefully expel the excess gas and reweigh the bag to measure the amount of CO2 lost,
and determine whether that mass corresponds to one or two moles per mole of calcium. The pH
of the final mixture will be measured as well; the first reaction should be close to neutral (both a
weak base and a weak acid are formed), while the second should be noticeably acidic (from the
CO2 ).
An excellent answer will clearly note both qualitative and quantitative observations that will be
made, and will indicate how they will be interpreted to provide evidence of the nature of the
reaction.
An average answer will note qualitative or quantitative observations that will be made, but will
not have clear indications of how those observations will be used.
A poor answer will show little awareness of the significance of mixing the reagents, or will not
describe observations that will be attempted.
Alternative approaches might include titrating the reagents against each other to determine their
mole ratio in the balanced reaction (using cessation of bubbling to mark the endpoint); weight
the precipitate formed; measuring the volume of CO 2 produced by inflating the bag and either
directly measuring its volume by water displacement or its dimensions, or by measuring the
drop in weight of the sealed bag after reaction and calculating the volume change from the mass
of air displaced; observing the pH of the precipitate after resuspension in pure water (Ca(OH)2 )
is noticeably basic, CaCO3 is not); observing the exothermicity of the reaction.
Key for 2005 USNCO National Exam, Part III
2.
Data and observations
Weighed 1.15 g CaCl 2 into one corner of the bag. Weighed 1.72 g NaHCO3 and carefully added
to the other corner of the bag. Added about 15 mL water and carefully added to the CaCl 2 side;
the compound dissolved and the water got hot. Let the solution cool, then sealed the bag and
weighed; total mass – 23.03 g. Allowed the CaCl2 solution to mix with the solid sodium
bicarbonate; the solution gets colder and bubbles vigorously, partially inflating the bag. The
solution turns milky white. After bubbling stops, open the bag and carefully expel the excess
gas; the mass is now 22.51 g. The pH of the suspension is about 6.
3.
CaCl2
A balanced chemical equation and evidence of reaction.
+
2NaHCO3 
CaCO3 (s) + CO2 (g) + H2 O
+
2 Na+(aq) + 2 Cl– (aq)
An excellent equation will be balanced and demonstrate reasonable chemical reactivity; an
average equation will be chemically reasonable, but may show inappropriate phases (e.g.
NaCl(s)); a poor equation will have chemically unreasonable products (H2 , Cl2 , HOCl; or the
production of Ca(OH)2 and H+ at the same time).
Evidence for reaction:
A colorless gas is evolved, confirming CO 2.
A white precipitate forms, consistent with CaCO 3.
The pH of the final solution is weakly acidic, consistent with formation of CO2 but not with OHof H 3O+.
The solution initially becomes hot, and then cools; CaCl2 dissolves exothermically, but release
of CO2 is endothermic.
An excellent answer will clearly relate all observations to the features of the reaction; an
average answer will relate only some observations to the reaction; a poor answer does not relate
the observations to the equation.
4.
Calculations
1.15 g CaCl2 /(110.9 g/mol) = 1.04 × 10–2 mol
1.72 g NaHCO3 /(84 g/mol) = 2.05 × 10–2 mol (limiting reagent)
Mass CO2 released = 23.03 g – 22.51 g = 0.52 g
0.52 g CO2 /(44 g/mol) = 0.012 mol
0.012 mol CO2 /0.0205 mol NaHCO3 = 0.59 mol CO2 per mol bicarbonate, close to the 0.5 mol
expected from the balanced reaction.
An excellent answer relates the mass of produced product to the predicted mass of CO2 based on
the balanced equation and the number of moles of the limiting reagent; an average answer will
note the amount of mass lost but will not relate it to the amounts of reagent; a poor answer will
note the number of moles of reagent but not of product.
Key for 2005 USNCO National Exam, Part III
Lab Problem 2
You have been given five vials, labeled #1-5. These vials contain methanol, 2-propanol,
acetone, hexane and water (though not necessarily in this order). You have also been given a
container of table sugar (C12H22O11 ). Design and carry out an experiment to determine which
liquid is in each labeled vial. You have access to a clock or timer.
This problem asks students to distinguish 5 different liquids.
Observations and data can be used to identify each liquid by observing:
1. Relative viscosity
2. Miscibility
3. Evaporation rates
4. Solubility with the provided sucrose sample
5. Rates of solubility with sucrose
6. Odor
7. Densities
Excellent answers for this problem included at least three of these observations in chart or table
form as part of their plan. In the data/observations, we looked for quantitative data (time for
evaporation, time for completely dissolved sugar samples using set volumes of each liquid and
set masses of sugar samples, etc.). Credit was given for multiple trials with quantitative data and
for complete observations with all five samples. Conclusions included correct identification of
each of the five liquids using more than one observation for each as evidence for student results.
The best student answers were organized, provided observations in chart or tabulated form, and
had a high degree of specificity with units were included for quantitative data.
Sample data:
A single drop of each liquid was placed on a watch glass and the time was recorded it took for
that drop to completely evaporate:
Vial
#1
#2
#3
#4
#5
liquid
2-propanol
Water
Methanol
Hexane
acetone
evaporation time (one drop) in min. (at room temp.)
2:15
still there
1:05
0:30
0:20
These data correspond nicely to predictions based on the degree of hydrogen bonding and
relative intermolecular forces of attractions between molecules of each of the samples present.
Solubility with sucrose:
Vial #2 dissolved the sugar easiest and quickest and in the greatest amount, evidence that #2 was
water; #1 and #3 slightly; #4 and #5 less so. Excellent student answers showed experiments with
consistent amounts of sugar and volumes of each vial used with recorded times for completely
dissolved sugar observed.
Miscibility:
Vial #4 was immiscible with #2, forming two distinct layers, conclusive evidence that #4 was
hexane on the top layer, water on the bottom layer. Vials #1, 2, 3 were miscible with one
another; #4 slightly with #5.
Key for 2005 USNCO National Exam, Part III
Odors:
For excellent data, students might have noted that #2 had no apparent odor, #1, 3 had an
‘alcohol-like’ smell, or even ‘smells like rubbing alcohol (#1), that #4 smelled like ‘gasoline’, or
that #5 had a ‘nail polish-like’ smell.
Key for 2005 USNCO National Exam, Part III
2006 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM PART 1
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
David W. Hostage, Taft School, CT
William Bond, Snohomish High School, WA
Adele Mouakad, St. John’s School, PR
Peter E. Demmin (retired), Amherst Central High School, NY
Jane Nagurney, Scranton Preparatory School,
PA
Marian Dewane, Centennial High School, ID
Ronald O. Ragsdale, University of Utah, UT
Kimberly Gardner, United States Air Force Academy, CO,
Todd Trout, Lancaster Country Day School, PA
Preston Hayes, Glenbrook South High School, IL
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April
25, 2006, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you
are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer
sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID
number because you will use the same number on both Parts II and III. Each item in Part I consists o f a question or an incomplete
statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement.
Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination,
but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at
the end of one hour and 30 minutes), you must turn in all testing materials , scratch paper, and your Scantron answer sheet. Do not
forget to turn in your U.S. citizenship statement before leaving the testing site today.
Page 1
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
ABBREVIATIONS AND SYMBOLS
A Faraday constant
F molal
atm formula molar mass
M
molar
u free energy
G molar mass
A frequency
ν mole
NA gas constant
R Planck’s constant
°C gram
pressure
g
c heat capacity
rate constant
Cp
C hour
retention factor
h
E joule
J second
Ea kelvin
K temperature, K
H kilo– prefix
k time
S
liter
L volt
K
milli– prefix
m
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
T
t
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s –1
0 °C = 273.15 K
1 atm = 760 mmHg
EQUATIONS
E= E −
o
1
1A
1
H
 −∆H  1 
ln K = 
  + constant
 R  T 
RT
ln Q
nF
k  E  1 1 
ln 2  = a  − 
 k1  R  T1 T2 
PERIODIC TABLE OF THE ELEMENTS
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
6.941
9.012
10.81
12.01
11
Na
12
Mg
13
Al
14
Si
22.99
24.31
26.98
19
K
1.008
18
8A
2
He
16
6A
8
O
17
7A
9
F
14.01
16.00
19.00
20.18
15
P
16
S
17
Cl
18
Ar
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
114
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
Page 2
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
Property of the ACS DivCHED Examinations Institute
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 3
DIRECTIONS
§ When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
§ This is a single use exam, so you may make marks in the test booklet.
§ There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
§ Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which substance is NOT paired correctly with its name?
(A) baking soda - potassium hydrogen tartrate
(B) chalk - calcium carbonate
(C) Epsom salt - magnesium sulfate heptahydrate
(D) Plaster of Paris - calcium sulfate hemihydrate
2. Which acid should be stored in plastic containers rather
than in glass ones?
(A) hydrofluoric acid
(B) nitric acid
(C) phosphoric acid
(D) sulfuric acid
(B) silicon
(C) sulfur
(D) tin
(B) HCl
(C) NH3
(D) O3
5. Which technique is preferred for delivering a solid into a
pre-weighed beaker for weighing?
(A) Transfer more of the reagent than is needed to the
beaker. Return the excess to the bottle with a
spatula.
(B) Transfer the desired amount of solid from the
reagent bottle by holding the neck of the open bottle
over the beaker and tapping the bottle. Then weigh
the beaker and solid.
(C) Weigh a spatula, scoop the desired amount of solid
from the bottle, transfer it to the beaker and reweigh
the spatula.
(D) Weigh a piece of filter paper, tap the neck of the
bottle to transfer solid to the filter paper, weigh the
filter paper and transfer the solid to the beaker.
6. Bronze is an alloy of
(A) copper and tin
(B) copper and zinc
(C) nickel and tin
(D) nickel and zinc
Page 4
(B) 0.313 M
(C) 0.500 M
(D) 0.625 M
8. What is the concentration of the solution that results
from mixing 40.0 mL of 0.200 M HCl with 60.0 mL of 0.100
M NaOH? (You may assume the volumes are additive.)
(A) 0.150 M NaCl
(C) 0.0200 M NaCl and 0.0600 M HCl
(D) 0.0600 M NaCl and 0.0200 M HCl
9. Mole fractions are
I. freezing point depression
typically used to
II. osmotic pressure
calculate which
III. vapor pressure
properties for solutions containing nonvolatile solutes?
4. Which gas is odorless?
(A) CH4
(A) 0.0301 M
(B) 0.0200 M NaCl and 0.0200 M HCl
3. Which element does NOT occur as distinct allotropes at
temperatures between 0°C and 150°C?
(A) phosphorus
7. What is the molarity of KI in a solution that is 5.00% KI
by mass and has a density of 1.038 g·cm-3?
(A) I only
(B) III only
(C) I and II only
(D) II and III only
10. An unknown anion in solution is to be identified by
adding Ag+ and Ba 2+ ions to separate portions of it.
Which anion would produce the results listed for it?
(+ indicates the presence of a precipitate)
Ag +
Ba2+
(A) carbonate
+
–
(B) hydroxide
–
+
(C) iodide
+
–
(D) sulfide
–
–
11. A 1.0 L portion of a 0.30m solution of which of the
following would be most effective at removing ice from a
sidewalk?
(A) C6H12O6
(B) NaBr
(C) KNO3
(D) CaCl2
Property of the ACS DivCHED Examinations Institute
12.
C6H6 + Br2 r C6H5Br + HBr
In an experiment to prepare bromobenzene according to
the equation, a student reacted 20.0 g of C6H6 with 0.310
mol of bromine. If 28.0 g of C6H5Br was obtained, what
was the percentage yield?
(A) 31.5
(B) 40.3
(C) 57.6
13. When the substances are
arranged in order of increasing
boiling points, which
order is correct?
(B) II < III < I
(C) III < II < I
(D) III < I < II
14. A 225 mL sample of H2 is
Vapor Pressure at 25 °C
collected over water at
H2O
24 mmHg
25 °C and 735 mmHg pressure. Which expression
represents the set-up to find the volume of dry H2 at 0 °C
and 1 atmosphere?
225 × (735 − 24 )× 273
760 × 298
(A)
V=
(B)
V=
(C)
V=
(D)
225 × (735 + 24 )× 298
V=
760 × 273
Pressure
(D) 69.7
I. CH3(CH2)3CH3
II. CH3CH2CH(CH3)2
III. C(CH3)4
(A) I < II < III
17. In the van der Waals equation for real gases, corrections
are introduced for both the pressure and the volume terms
of the Ideal Gas Equation. Identify the origin of both
correction factors and specify whether each is added to or
subtracted from the corresponding term.
Volume
(A) attractive forces /
subtracted
molecular size / added
(B) attractive forces / added
molecular size / subtracted
(C) molecular size / subtracted
attractive forces / added
(D) molecular size / added
attractive forces /
subtracted
18. The structure of a unit
cell of an oxide of
niobium is depicted
here. Niobiums are dark
and oxygens are light.
What is the empirical
formula of this
compound?
225 × 760 × 298
(735 − 24) × 273
225 × 273 × 760
(735 + 24 ) × 298
(A) NbO
15. What is ? Hvap for the
substance whose vapor
pressure is represented by
the diagram?
(C) NbO3
(D) Nb 2O3
19. For a reaction that is exothermic and non-spontaneous at
25 °C, which quantity must be positive?
(A) ?E°
(B) ?G°
(C) ?H°
(D) ?S°
20. Use the thermochemical data given to calculate ? Hf° for
N2O5(g) in kJ·mol-1.
N2(g) + O2(g) r 2NO(g)
?H° = +180.5 kJ
2NO(g) + O2(g) r 2NO2(g)
?H° = –114.1 kJ
4NO2(g) + O2(g) r 2N2O5(g)
?H° = –110.2 kJ
(A) –332.8
(A) 4.8 kJ·mol-1
(B) 33 kJ·mol-1
(C) 44 kJ·mol-1
(D) 50 kJ·mol-1
16. What occurs
when liquid
CH2F2
evaporates?
(B) NbO2
I. Dispersion forces are overcome.
II. Dipole-dipole forces are overcome.
III. Covalent bonds are broken.
(A) II only
(B) III only
(C) I and II only
(D) I, II and III
(B) –43.8
(C) 11.3
(D) 22.6
21. Bromine boils at 59°C with ?H°vap = 29.6 kJ·mol-1. What is
the value of ?S°vap in J·mol-1·K-1?
(A) 11.2
(B) 89.2
(C) 501
(D) 1750
22. The Ksp of calcium fluoride is 3.2×10-11. Calculate the ?G°
(in kJ·mol-1) for the dissolving of solid calcium fluoride at
25°C.
(A) 2.18
(B) 5.02
(C) 26.0
(D) 59.9
23. For which exothermic reaction is ?E more negative than
?H?
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 5
29. According to the reaction
profile given, which
reaction step is ratedetermining in the forward
direction?
(A) Br2(l) r Br2(g)
(B) 2C(s) + O2(g) r 2CO(g)
(C) H2(g) + F2(g) r 2HF(g)
(D) 2SO2(g) + O2(g) r 2SO2(g)
24. For a reaction at 25°C, ?G = 12.7 kJ when the reaction
quotient, Q = 10.0. What is the value of ?G° for this
reaction?
(A) –12.1 kJ
(B) 7.0 kJ
(A) I r II
(B) II r III
(C) 18.4 kJ
(D) 37.5 kJ
(C) III r II
(D) III r IV
25. How can the rate of reaction at a specific time be
determined from a graph of concentration against time?
(A) concentration at that time divided by the time
(B) logarithm of the concentration divided by the time
(C) absolute value of the slope of the graph at that time
(D) logarithm of the slope divided by the time
26. The rate constant for the radioactive decay of C-11 is
0.0341 min -1. How long will it take for a sample of C-11 to
decrease to 1/4 of its original activity?
(A) 20.3 min
(B) 29.3 min
(C) 40.6 min
(D) 58.6 min
(B) ln[A] vs time
(C) [A]2 vs time
(D) 1/ln[A] vs time
(A) This mechanism is consistent with the rate law IF
step 1 is the rate determining step.
(B) This mechanism is consistent with the rate law IF
step 2 is the rate determining step.
27. If a reaction A r B has the rate law k[A]2, which graph
produces a straight line?
(A) 1/[A] vs time
30. For the reaction;
2H2(g) + 2NO(g) r N2(g) + 2H2O(g),
rate = k[H2][NO]2.
This mechanism has been proposed:
step 1 H2 + NO r H2O + N
step 2 N + NO r N2 + O
step 3 O + H2 r H2O
Which statement about this rate law and mechanism is
correct?
28. Two unimolecular reactions, I and II, have the same rate
constant at 25 °C but Ea for reaction I is larger than Ea for
reaction II. Which statement about these two reactions is
correct?
(A) k reaction I is the same as k reaction II at all temperatures.
(B) k reaction I is larger than k reaction II at lower temperatures
but smaller at higher temperatures.
(C) k reaction I is smaller than k reaction II at lower temperatures
but larger at higher temperatures.
(D) k reaction I is larger than k reaction II at temperatures both
lower and higher than 25 °C.
(C) This mechanism is consistent with the rate law IF
step 3 is the rate determining step.
(D) This mechanism can not be consistent with the rate
law, regardless of which step is rate-determining.
31. C(s) + CO2(g) s 2CO(g)
I. raising the temperature
II. adding solid C
If this system is at
III. decreasing the pressure
equilibrium, which
change(s) will alter the value of Kp?
(A) I only
(B) II only
(C) I and III only
(D) II and III only
32. A 0.10 M solution of a weak acid is 5.75% ionized. What
is the Ka value for this acid?
(A) 3.3×10-3
(B) 3.5×10-4
(C) 4.2×10-5
(D) 3.3×10-5
33. Which base is most suitable to prepare a buffer solution
with a pH = 11.00?
(A) ammonia (Kb = 1.8×10-5)
(B) aniline (Kb = 4.0×10-10)
(C) methylamine (Kb = 4.4×10-4)
(D) pyridine (Kb = 1.7×10-9)
Page 6
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
34. Calculate the pH of H2CO3 Acid Ionization Constants
a 0.10 M solution
K1a
4.4×10-7
of H2CO3.
K2a
4.7×10-11
(A) 3.68
(B) 5.76
(C) 7.36
(D) 9.34
35. Which saturated solution has the highest [OH-]?
(A) aluminum hydroxide (Ksp = 1.8×10-32)
(B) calcium hydroxide (Ksp = 8.0×10-6)
(C) iron(II) hydroxide (Ksp = 1.6×10-14)
(D) magnesium hydroxide (Ksp = 1.2×10-11)
(A) Increasing the [Cu 2+] two-fold has the same effect on
the cell voltage as increasing the [Ag+ ] four-fold.
(B) Decreasing the [Cu 2+] ten-fold has the same effect on
the cell voltage as decreasing the [Ag+ ] by the same
ratio.
(C) Decreasing the [Cu 2+] ten-fold has less effect on the
cell voltage than decreasing the [Ag+ ] by the same
amount.
(D) Doubling the sizes of the cathode has exactly the
same effect on the cell voltage as decreasing the
[Cu 2+] by a factor of two.
39. 3Ni2+ + 2Al r 2Al3+ + 3Ni
E° = 1.41 V
For the reaction given, which expression gives the value
of ?G° in kJ·mol-1?
36. Consider these mixtures:
Mixture I. 100 mL of 0.006 M Pb(NO3)2 plus 50 mL
of 0.003 M NaBr
Mixture II. 100 mL of 0.008 M Pb(NO3)2 plus 100 mL of
0.006 M NaBr
Which statement is correct?
Ksp
PbBr2
6.6×10-6
(A) −3× 96.5
1.41
(B) −6× 96.5
1.41
(C) −3× 96.5× 1.41
(D) −6 × 96.5× 1.41
(A) A precipitate will not form in either mixture.
(B) A precipitate will form only in mixture I.
40. In which species is the oxidation number for hydrogen
different from those in the other three?
(C) A precipitate will form only in mixture II.
(A) AlH3
(B) H3AsO4
(D) A precipitate will form in both mixtures.
(C) H3PO3
(D) NH3
37. The equation for one of the half-reactions in a lead
storage battery is:
PbO2 + 4H+ + SO42- + 2e - r PbSO4 + 2H2O
What happens to the properties of the electrolyte as this
cell discharges?
41. A solution
Standard Reduction Potential (V)
containing
Ni2+(aq) + 2e – r Ni(s)
–0.236
equimolar
Sn 2+(aq) + 2e – r Sn(s)
–0.141
amounts of NiCl2 Br2(aq) + 2e – r 2Br –(aq)
1.077
and SnBr2 is
Cl2(aq) + 2e – r 2Cl –(aq)
1.360
electrolyzed using
a 9V battery and graphite electrodes. What are the first
products formed?
Density
pH
(A)
increases
increases
(B)
increases
decreases
(A) Ni(s) at cathode, Cl2(aq) at anode
(C)
decreases
decreases
(B) Ni(s) at cathode, Br2(aq) at anode
(D)
decreases
increases
(C) Sn(s) at cathode, Br2(aq) at anode
38. For the voltaic cell based on this reaction:
2Ag+ (aq) + Cu r Cu 2+(aq) + 2Ag
the concentrations of the aqueous ions and sizes of the
electrodes can be changed independently. Which
statement is correct?
(D) Sn(s) at cathode, Cl2(aq) at anode
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 7
42. True statements about
the system shown
after the passage of
one Faraday of
electricity include
which of those given?
(A) K-38
(B) K-39
I. The number of moles of Al formed is greater than
the number of moles of silver formed.
II. The final [Al3+] is greater than the final [Ag+ ].
III. The number of electrons reacting with Al3+ ions is
the same as the number reacting with Ag+ ions.
(A) I only
(B) I and III only
(C) II and III only
(D) I, II and III
43. When l = 3, what are the possible values for the quantum
number ml?
(A) 2, 1, 0
(B) 3, 2, 1, 0
(C) 2, 1, 0, -1, -2
(D) 3, 2, 1, 0, -1, -2, -3
44. The first ionization energy of cesium is 6.24×10-19 J/atom.
What is the minimum frequency of light that is required to
ionize a cesium atom?
(A) 1.06×10-15 s -1
(B) 4.13×1014 s -1
(C) 9.42×1014 s -1
(D) 1.60×1018 s -1
(A) H2Se, H2S, H2O
(B) H2S, H2Se, H2O
(C) H2S, H2O, H2Se
(D) H2O, H2S, H2Se
(B) II only
(C) I and III only
(D) II and III only
51. What is the shape of the TeF5– anion?
(A) see-saw
(B) square pyramidal
(C) trigonal pyramidal
(D) trigonal bipyramidal
52. How many sigma and pi bonds are in maleic acid,
HO2CCHCHCO2H?
(A) 7 sigma, 2 pi
(B) 8 sigma, 3 pi
(C) 9 sigma, 2 pi
(D) 11 sigma, 3 pi
53. How many isomers exist for the octahedral compound,
Pt(NH3)2Cl4?
(A) 1
–
2–
+
2–
–
(B) K , S , Cl
2–
–
+
–
2–
+
(B) 2
(C) 3
(D) 4
54. What is the formal charge on the sulfur atom in SO2?
(Assume a Lewis dot structure in which all atoms obey
the octet rule.)
45. When the isoelectronic ions, Cl–, S2– and K+ are arranged
in order of increasing size, which order is correct?
+
I. Na(g) r Na + (g) + e–
II. F(g) + e– r F–(g)
III. Na + (g) + F–(g) r NaF(s)
(A) I only
(A) +1
(C) S , Cl , K
(D) K-43
49. When the species listed are arranged in order of
increasing bond angle, which order is correct?
50. Which terms are
exothermic for the
formation of NaF(s)?
(A) K , Cl , S
(C) K-42
(C) –1
(D) –2
55. How many structural isomers are possible for C6H14?
(A) 2
(D) Cl , S , K
(B) +2
(B) 3
(C) 4
(D) 5
56. Which is an ester?
46. Which is most similar for the elements in a group in the
periodic table?
(A) CH3COOCH2CH3
(B) (CH3)3COOC(CH3)3
(C) CH3OCH3
(D) (CH3)3CCOOH
(A) physical state
57. Which type of reaction is typical of aromatic compounds?
(B) melting point
(A) addition
(C) first ionization energy
(B) free-radical substitution
(D) ground state electron configuration
(C) substitution by positively-charged reagents
47. How many unpaired electrons are present in a gaseous
Co 3+ ion in its ground state?
(A) 1
(B) 3
(C) 4
(D) 5
(D) substitution by negatively-charged reagents
58. What is the IUPAC name of (CH3)2CHCH=CHCH3?
48. Which nucleus is not radioactive?
Page 8
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
(A) 1,2-methyl-isopropylethene
(B) 1,1-dimethyl-2-butene
(C) 1-isopropylpropene
(D) 4-methyl-2-pentene
59. Which compound can exist in optically active forms?
(A) CH3CH2CH2CH2OH
(B) CH3CH2CH(OH)CH3
(C) (CH3)2CHCH2OH
(D) (CH3)3COH
60. How many different tripeptides can be formed from the
amino acids glycine, alanine and valine if each is used
only once in each tripeptide?
(A) 3
(B) 4
(C) 5
(D) 6
END OF PART I
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 9
CHEMISTRY OLYMPIAD 2006
National Test, PART I
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
A
A
B
A
B
A
B
D
B
C
D
D
C
A
C
C
B
A
B
C
B
D
D
B
C
C
A
C
D
B
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Property of the ACS DivCHED Examinations Institute
Answer
A
B
C
A
B
A
D
C
D
A
C
C
D
C
A
D
C
B
A
D
B
D
B
A
D
A
C
D
B
D
2006 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II – ANSWER KEY
Prepared by the American Chemical Society Olympiad Examinations Task Force
1.
(11%) A(g) + 3B(g) r 2C(g)
Use the tabulated data to answer the questions about this reaction, which is carried out in a 1.0 L container at 25°C.
Experiment
Ao, mol
Bo, mol
Initial rate of formation of C, mol .L-1.min-1
1
2
3
a.
b.
c.
d.
0.10
0.20
0.10
0.10
0.20
0.20
0.25
2.0
2.0
For experiment 1, give the initial rate of disappearance of
i. A ii. B
Determine the orders of A and B and write the rate law for the reaction.
Calculate the value of the rate constant and give its units.
For the initial amounts of A and B in experiment 1, state the initial rate of formation of C under the following conditions.
Justify your answer in each case.
i.
0.50 mol of neon gas is added to the 1.0 L container.
ii.
the volume of the container is increased to 2.0 L.
a. rate of formation C = 0.25 mol·L-1·min-1 so
1
= 0.13 mol ⋅ L-1 ⋅min −1
2
i. rate of disappearance of A =
3
0.25 × = 0.38 mol ⋅ L-1 ⋅ min −1
2
ii rate of disappearance of B =
0.25 ×
b. for experiments 2 and 3: [B] is constant while [A] doubles and the rate of the reaction is unchanged.
The reaction order with respect to A is zero.
for experiments 3 and 1: [A] is constant while [B] doubles and the rate of the reaction increases by a factor of 8.
The reaction order with respect to B is 3.
Rate = k[B]
3
k=
c.
0.25
(0.10)
3
= 250 L2 ⋅ mol -2 ⋅ min
d. i. The rate of formation of C is 0.25 mol·L-1·min-1. Adding an inert gas does not change the rate law.
ii. The rate of formation of C is 0.031 mol·L-1·min-1. Doubling the volume of the container changes the concentration of B.
[B] = 0.050 mol·L-1·
Rate = k [B] = 250 L2 ⋅ mol -2 ⋅ min [0.050 mol ⋅ L-1] = 0.031
3
3
2. (13%) A 0.472 g sample of an alloy of tin and bismuth is dissolved in sulfuric acid to produce tin(II) and bismuth(III) ions. This
solution is diluted to the mark in a 100 mL volumetric flask and 25.00 mL aliquots are titrated with a 0.0107 M solution of KMnO4,
forming tin(IV) and manganese(II) ions. (The bismuth ions are unaffected during this titration.)
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 1
a.
b.
c.
d.
e.
Write a balanced equation for the reaction of the MnO4- ion with Sn(II) in acid solution.
If an average titration requires 15.61 mL of the MnO4- solution, calculate the number of moles of MnO4- used in an average
titration.
Determine the percentage of tin in the alloy.
State how the end point of the titration is detected.
Describe and explain the effect on the calculated percentage of tin in the alloy if the same volume of MnO4- solution is used
with the following differences:
i.
During the titration the solution pH increases so that MnO2 is formed rather than Mn(II).
ii.
The solution volume in the volumetric flask was above the mark on the flask but a volume of 100. mL was
assumed.
iii.
The sample of the original alloy had an oxide coating on it.
–
2+
+ → 2Mn 2+ + 5Sn 4 + + 8H O
2
a. 2MnO 4 + 5Sn + 16H 
b.
0.01561 L × 0.0107 mol ⋅ L-1 = 1.669 ×10 -4 mol
1.669 × 10 -4 mol ×
c.
%Sn =
5 mol Sn 2+ 118.7 g Sn 2+
×
= 0.04953 g of Sn in 25 mL
2 mol MnO –4 1 mol Sn 2+
4 × 0.04953 g
×100 = 41.97%
0.472 g sample
d. The endpoint is shown by the persistence of a faint purple color. (Indicating that there is an excess of MnO4- .)
e. i, The %Sn that is determined is too high. The MnO4-/Sn2+ ratio is 2:3 when MnO2 is formed. Thus, for a given number of moles of
2
2
Sn 3 as many moles of MnO4- will be used rather than 5 . Since the calculations assume MnO4- , the latter ratio is used. So
5 ×2 = 5
( 2 3 3 ) of the correct moles of Sn would be calculated.
ii. The %Sn that is determined would be too low. Each 25 mL aliquot of contain fewer Sn2+ ions, requiring less MnO4- in the
titration.
iii. The %Sn that is determined would be too low. The oxide coating leads to fewer Sn2+ ions released into solution per g of sample
weighed.
3.
(15%) In water, HCN is a weak acid with pKa = 9.6.
a. Calculate the the Ka and the [H+ ] in a 0.15 M solution of HCN.
b. In a closed system, these equilibria are established:
HCN(g) s HCN(aq)
HCN(aq) s H+ (aq) + CN-(aq)
i.
Calculate Kp for the equilibrium between HCN(g) and HCN(aq) at 298 K.
? Gf° kJ.mol-1
HCN(g)
124.7
HCN(aq) 119.7
ii.
If the total cyanide concentration in solution (i.e. [CN-] + [HCN]) is 0.10M, calculate the partial pressure of HCN(g)
in this system at pH = 7.
iii.
A concentration of 300 ppm of HCN in air is reported to be toxic to humans after a few minutes exposure.
Determine the ratio of the pressure calculated in 3.b.ii. to this value.
c. Gold can be extracted from its ores by reacting the ore with O2 gas in the presence of aqueous CN- ions according to this
equation.
4Au(s) + 8CN-(aq) + O2(g) + 2H2O(l) s 4Au(CN)2-(aq) + 4OH-(aq)
i.
Write the equilibrium expression for this reaction.
ii.
At fixed [CN–] and O2 pressure will the amount of Au(CN)2– be greatest at high or low pH? Justify your
answer.
iii.
What purpose does O 2 serve in the extraction process?
−9.6 = 2.5× 10 −10
a. pKa = 9.6, so K a = 10
2
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Ka =
[H + ][CN– ] = 2.5 ×10−10 = x 2
0.15 so x = [H+] = 6.1×10-6
[HCN]
b. i. HCN(g) s HCN(aq) so, ? Go = 119.7 – 124.7 = –5.0 kJ·mol-1
? Go = –RTlnK, so –5000 J·mol-1 = –(8.314 J·mol-1·K-1)(298 K)lnK
K = e 2.018 = 7.52
ii. [CN–]+[HCN]=0.10, pH=7.0 so [H+]=1.0×10-7
2.5 ×10−10 =
(1.0 × 10−7 )(0.10 − x)
x
2.5 ×10−10 x = 1.0 × 10−8 − 1.0 × 10−7 x
1.0025 ×10−7 x = 1.0 ×10−8 so x = [HCN]eq = 0.0998
Kp =
[HCN]eq
pHCN
pHCN =
so
0.0998
= 0.0133 atm
7.52
iii. This problem is ambiguous in terms of how it might be solved – by mass or by volume
Perhaps, the more obvious method is to solve by volume:
0.0133 atm
= 0.0133 = 13300 ppm (vol)
1 atm
so the ratio is 13300 ppm / 300 ppm = 44.3
We can also use mass…
0.0133 atm HCN 27 g mol -1 HCN
×
= 0.0125 = 12500 ppm (mass)
-1
1 atm air
28.8 g mol air
so the ratio is 12500 ppm / 300 ppm = 41.6
[Au(CN)–2 ] [OH – ]
K=
[CN– ]8 [O 2]
4
c. i.
4
ii. Low pH should be used because it will lower the [OH–], shifting the equilibrium towards production of products.
iii. O2 is an oxidizing agent in this reaction.
4.
(13%) The combustion of ethane, C2H6, produces carbon dioxide and liquid water at 25°C.
a.
b.
c.
d.
Write an equation for this reaction.
Given that ? H°comb for ethane under these conditions is -1560.5 kJ/mol ethane, calculate
i. ? Hf° for ethane.
? Hf° kJ.mol-1
Bond Energies, kJ.mol-1
CO2(g)
–393.5
ii. the bond energy of the C=O bond.
C–C
347
H2O(l)
–285.8
H–C
413
Given ? G°= –1467.5 kJ/mol, Calculate ? S° for this reaction in J.mol-1.K-1.
H–O
464
Compared with combustion to form liquid water at 25 °C, how would combustion
O=O
495
to form H2O(g) affect each of the following;
i. ? H°combustion ii. ? S°combustion iii. ? G°combustion
a. 2C2H6(g) + 7O2(g) r 4CO2(g) + 6H2O(l) (note: dividing all coefficients by 2 was given full credit)
b. i.
∆H rxn = 4∆H f (CO 2 ) + 6∆H f (H 2O) – 2∆H f (C 2H 6 )
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 3
so –3121 = 4(–393.5) + 6(–285.8) – 2∆Hf(C 2H6) and ∆Hf(C 2H6) = –83.9 kJ·mol-1
ii.
∆H rxn = 2BE C– C +12BE C– H + 7BE O= O – 8BE C= O – 12BE H– O
so –3121 = 2(347) + 12(413) + 7(495) – 12(464) – 8BEC=O and = BEC=O = 6668 / 8 = 833 kJ.mol-1
c.
∆G o = ∆H o – T∆S o so –1467.5 = –1560.5 – 298 ∆So solving for ∆So yields
∆So = (1560.5 – 1467.5) / –298 = –312 J·K-1
d. i.
. The measured ∆Hcombustion is less negative because the heat of vaporization was not released.
ii. The ∆Scombustion is more positive (less negative) because H2O(g) has greater entropy than H2O(l)
iii. The ∆Gcombustion is less negative due to the combination of these two effects.
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You
need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Excess carbon dioxide is bubbled through a suspension of calcium hydroxide.
b. Acidified solutions of cerium(IV) and iron(II) are mixed.
c. Solid calcium carbide is added to water.
d. Excess concentrated ammonia is added to aqueous nickel(II) nitrate.
e. Solutions of silver acetate and hydrobromic acid are mixed.
f. Gaseous hydrogen chloride is reacted with gaseous propene.
→ Ca 2+ + HCO –3
a. Ca(OH) 2 + CO 2 + H 2O 
→ Ce 3+ + Fe 3 +
b. Ce 4 + + Fe 2+ 
c.
CaC 2 + H 2 O 
→ Ca 2+ + OH – + C 2 H 2
d.
Ni 2+ + NH 3 
→ Ni(NH 3 ) 2+
6
e.
Ag + C 2 H 3O 2 + H + Br 
→ HC 2 H 3 O 2 + AgBr
f.
HCl + H 3C – CH = CH 2 
→ H 3C – CHCl – CH 3
+
–
+
(note: Ca(OH)2 was accepted as a product)
–
6. (13%) Answer the following questions.
a. For the molecule XeOF4.
i.
Write a Lewis structure.
ii.
Predict its geometry and specify the bond angles.
iii.
State whether it is polar or nonpolar. Explain your answer.
b. Nitric acid, HNO3, is a strong acid while phosphoric acid, H3PO4, is a weak acid .
i.
Draw Lewis structures for each acid .
ii.
Explain why H3PO4 is stable while H3NO4 is not.
iii.
Suggest and explain two reasons that nitric acid is stronger than phosphoric acid.
c. Ethane and diborane have similar formulas, C2H6 and B2H6, but B2H6 is more reactive. Sketch the structure of C2H6 and
explain why B2H6 does not adopt this structure.
4
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
O
O
F
F Xe
F
or
F
F
F Xe
F
F
a. i.
ii. square pyramidal with 90º bond angles.
O
F
F Xe
F
F
iii. The molecule is polar with charges…
O
O
H
O N
H
O
H O P O H
O
O
O N
O
b. i.
and
H
ii. H3NO4 would have more than 8 e– around N (OR) would put a positive (+) formal charge on N (OR) would be too sterically
hindered with 4 oxygen atoms around the nitrogen.
iii. (1) N is more electronegative that P, so electron density is shifted from H atoms towards the N, so the H+ can be more readily
removed. (2) NO3– is stabilized by resonance more than H2PO4–. (3) HNO3 has two free oxygen atoms that attract electron
density from the H atom, whereas H3PO4 has only one free oxygen atom.
H H
H
c.
7.
C C
H
H H
B2H6cannot adopt this structure because it has only 12 valence electrons where C2H6 has 14.
(11%) A common lecture demonstration involves electrolyzing a 1.0 M aqueous NaI solution containing phenolphthalein with a
9V battery.
a. Write a balanced equation for the half-reaction that occurs at the
i. anode. ii. cathode.
b. Describe what is observed in the solution at the
i. anode. ii. cathode.
c. If a current of 0.200 amperes is passed through a 25.0 mL solution for 90.0 minutes, calculate the;
i.
number of moles of electrons passed through the solution.
ii.
number of moles of each of the products formed.
a. i. 2I– r I2+ 2e–
ii. 2H2O + 2e– r H2+ 2OH–
b. i. At the anode, a yellow-brown color appears due to formation of iodine. If starch is added the color is blue
ii. At the cathode, bubbles form due to the formation of hydrogen gas. The solution turns pink if phenolphthalein is added.
c. i. 0.200 C ⋅s -1 × 90 min × 60 s ⋅ min -1 = 1080 C
1080 C ×
1 mol e –
= 1.12 ×10−2 mol e –
96500 C
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 5
1.12 ×10−2 mol e – ×
1 mol I 2
= 5.6 ×10−3 mol I 2
2 mol e –
1.12 ×10−2 mol e – ×
1 mol H 2
= 5.6 ×10 −3 mol H 2
2 mol e –
1.12 ×10−2 mol e – ×
2 mol OH –
= 1.12 ×10−2 mol OH –
2 mol e –
ii.
8.
(12%) There are four isomeric unsaturated compounds (alkenes) with the formula C4H8.
a. Draw and name each of these isomers.
b. These compounds all react with water in the presence of H2SO4 as a catalyst.
i. Name the type of compound formed in this reaction.
ii. Three of the four isomers form the same compound during this reaction. Identify these three isomers and outline your
reasoning.
c. Draw the structure of a saturated compound with the formula C4H8 and describe a chemical test that could be used to
distinguish between this compound and one of the alkenes above. (Describe the results obtained for the saturated and
unsaturated compound.)
a.
H
H
H
C C
H
C
H
H
H
C
H
1-butene
H
C
C
H
H
C H
H
H
H
C
H
H
2-methylpropene
H
H
C C H
H C
C
H
H
H
H
H
H
C H
C C
H C
H
H H
cis -2-butene
trans-2-butene
b. i. An alcohol.
ii. 1-butene, cis -2-butene and trans-2-butene all form 2-butanol. The OH of H2O will attack the 2° C of the double bond rather than
the 1o C.
H
H
H
C
H C
H
iii.
H
H
C
H
H
C
C H
C C H
C H
H
H
H
H
or
Add Br2 to each. Br2 will be decolored with an unsaturated compound because Br2 adds to the double bond. Br2 will not change in the
presence of saturated compounds.
6
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
END OF KEY PART II
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
mol
ν mole
NA gas constant
h
R Planck’s constant
°C gram
pressure
P
g
c heat capacity
k
Cp rate constant
C hour
retention factor
Rf
h
E joule
s
J second
Ea kelvin
c
K speed of light
H kilo- prefix
T
k temperature, K
S liter
t
L time
volt
V
CONSTANTS
R = 8.314 J·mol–1 ·K–1
R = 0.0821 L·atm·mol–1 ·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1 ·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s –1
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
112
116
118
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(277)
(289)
(293)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO Olympiad National Exam after April 25, 2006.
Page 7
2006 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Brookline High School, Brookline, MA
Chair
Linda Weber, Natick High School, Natick, MA
Sheldon L. Knoespel, Mott Community College, Flint, MI
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 26, 2006.
Page 1
2006 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
Turmeric, a natural compound, is added to mustard for flavor and color. It changes color from yellow to red at
a pH of 7.4. Mustard also contains acetic acid. Given a sample of 0.50 M NaOH and the packets of mustard,
create and perform an experiment to determine the mass percentage of acetic acid in mustard.
Lab Problem 2
Given a sample of 3.0 M hydrochloric acid, phenolphthalein, and some common laboratory equipment, devise
an experiment using both qualitative and quantitative evidence to determine the provided unknown metal
given these possible choices: Ag, Al, Ca, or Cr.
Page 2
Not valid for use as an USNCO National Examination after April 26, 2006.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
This is a titration experiment. Yellow mustard must be between 2.6 – 3.5% acetic acid by
law (See: Current CFR 21 for 2005 [ http://www.gpoaccess.gov/cfr/index.html
]http://www.gpoaccess.gov/cfr/index.html). Yellow mustard contains turmeric, here used
as an indicator for this experiment.
Before beginning your experiment, you must get
Examiner’s Initials:
Not valid for use as an USNCO National Examination after April 26, 2006.
Page 3
approval (for safety reasons) from the examiner.
2. Record your data and other observations.
3. Calculations.
The calculations would be:
1.
moles base used (V x M) = moles acid present
2.
moles acid present x molar mass acetic acid = mass acetic acid
3.
Percentage of acetic acid in mustard = mass acetic acid present / mass
mustard used
Sample Calculation:
0.50 g mustard weighed, titrated with a volume of 0.5 mL NaOH
moles OH- = 0.0005 L
x
0.5M = 0.00025 mol OH-
= 0.00025 mol H+ from HC2H3O2 in mustard
mass HC2H3O2 = 0.00025 mol
x
60 g/mol
=
0.015 g HC2H3O2
finally, % acetic acid in mustard = 0.015 g /.50 g x100 = approx. 3.0%
The percentage of acetic acid in your sample of mustard
=
___3.0%__________________
Excellent work:
Student was able to complete two or more trials and average their results, using a
minimum amount of both mustard and NaOH for each titration. Results were clearly shown
and observations, i.e. color changes and endpoint were clearly noted.
Student thought to make dilute aqueous solutions with each of the samples of mustard in
order to completely dissolve the mustard and be able to more clearly note a uniform and
lasting color change
Average work:
Student only completed one trial. Evidence of a titration was performed. Measurements
between trials were fairly consistent.
Below average work:
Student was not able to conclude that this was a titration experiment, or did so, but did
not perform the titration correctly to obtain a mass/volume of NaOH added. Only one trial
was performed. Measurements were inaccurate or inconsistent between trials.
Page 4
Not valid for use as an USNCO National Examination after April 26, 2006.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name:__________________________________________________________________________
ACS Section Name : ________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Students were to provide both qualitative and quantitative evidence to determine the
unknown metal. The metal provided was calcium. The results to this experiment should have
included both evidence form data obtained and exclusive information about what was not
observed from students’ previous chemical knowledge. Conclusions come from knowledge
about each metals’ reactivity to both water and HCl, with phenolphthalein, a possible
titration, and gas generation. Students might also have explored reactivity of the metal
with NaOH from Problem #1 (this is allowed, though not necessary to successfully complete
this problem).
Excellent work:
Student combined HCl with the unknown metal (Ca) to obtain hydrogen gas in the wellplate, clearly showing evidence of gas production and exothermic reaction. A titration
The student then performed this reaction with a measured amount of Ca and excess HCl
using the Luer-Lok syringe to quantify the hydrogen gas produced (since room temperature
and pressure were not given, student had to make some assumptions about the Kelvin
temperature and room pressure, perhaps estimating 298K and 1 atm) to determine the
expected volume of hydrogen and compare it to a theoretical volume produced from
Ca + HCl
CaCl2 + H2
Noting the color change when phenolphthalein is added to the metal reacted to either
water or HCl. It is possible that a student might have thought to combine mustard (from
Prob.#1) with the metal from this experiment. If so, mustard on the surface of Ca
produces over time a crusty white solid, Ca(C2H3O2)2 (there is no evidence of reaction
with mustard on the surface of Cr and with pure Ag, no visible reaction).
Concluding what DIDN’T occur:
If Cr + HCl
greenish color indicating CrCl3 (or green color with many
chromium salts
If Ag + HCl
no reaction
If Al + HCl
no visible reaction due to aluminum oxide layer (though
student might have attempted to dissolve the metal with the NaOH from Exp.
#1, if Al, would dissolve; Ca + NaOH
gives Ca(OH)2, a noticeable milky
white precipitate, with phenolphthalein produces a pink color.
a)
Sample titration experiment conclusions:
Reacting a 0.10 g metal turning with water completely, adding phenolphthalein, then
titrating with the 3M HCl to obtain a 2 : 1 ratio of OH- : H+ in solution, confirms
that OH- must be present in the metal hydroxide form, M(OH)2.
b)
Sample data for quantifying hydrogen gas generated using the Luer-Lok®
syringe:
One metal turning, approx. 0.07 g Ca in excess 3M HCl
Not valid for use as an USNCO National Examination after April 26, 2006.
Page 5
Begin at 12 mL mark on syringe
End at 40 mL mark on syringe
40 – 12 = 28 mL hydrogen gas generated, strongly exothermic reaction.
Assume room temp. 25oC (298K) and 1 atm:
using the ideal gas law, PV = nRT
(1 atm)(0.028L) = n (0.0821atm L/mol K) (293K) ; n = 0.00116 mol H2(g)
if given 0.07g of calcium, Ca + 2HCl
CaCl2(aq) + H2(g) , then 0.0035
g of hydrogen gas is produced, corresponds roughly to number of moles of
H2(g) made with these assumed conditions.
Average work:
Student reacted metal with HCl and concluded hydrogen gas was present but didn’t quantify
the gas produced, or did but incorrectly. Student wrote out possible reactions with the
other possibilities but did not do so correctly.
Below average work:
Student was unable to conclude that hydrogen gas was produced, did not use either a
titration or quantitative method of data collection, or unable to use the phenolphthalein
to qualitatively justify the metal.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
2. Record your data and other observations. (See comments above)
Page 6
Not valid for use as an USNCO National Examination after April 26, 2006.
3. Conclusions and Evidence.
The unknown metal is
= __Calcium__________
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
111
112
112
116
118
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(272)
(277)
(277)
(289)
(293)
(269)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 26, 2006.
Page 7
2007 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM PART 1
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
David W. Hostage, Taft School, CT
Peter E. Demmin (retired), Amherst Central High School, NY
Marian Dewane, Centennial High School, ID
Jane Nagurney, Scranton Preparatory School, PA
Kimberly Gardner, United States Air Force Academy, CO,
Preston Hayes, Glenbrook South High School, IL
Adele Mouakad, St. John’s School, PR
Ronald O. Ragsdale, University of Utah, UT
Todd Trout, Lancaster Country Day School, PA
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
May 1, 2007, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
ABBREVIATIONS AND SYMBOLS
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G molar mass
A frequency
ν mole
N A gas constant
R Planck’s constant
°C gram
g pressure
c heat capacity
C p rate constant
C hour
h retention factor
E joule
J second
Ea kelvin
K temperature, K
H kilo– prefix
k time
S
liter
L volt
K
milli– prefix
m
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
T
t
V
R = 8.314 J·mol –1·K–1
R = 0.0821 L·atm·mol –1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
N A = 6.022 × 10 23 mol–1
h = 6.626 × 10 –34 J·s
c = 2.998 × 10 8 m·s –1
0 °C = 273.15 K
1 atm = 760 mmHg
EQUATIONS
E = Eo −
1
1A
1
H
1.008
3
Li
RT
ln Q
nF
k  E  1 1 
ln 2  = a  − 
 k1  R  T1 T2 
 −ΔH  1 
ln K = 
  + constant
 R  T 
PERIODIC TABLE OF THE ELEMENTS
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
18
8A
2
He
4.003
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
19
K
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
26.98
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
(98)
44
Ru
101.1
45
Rh
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
85.47
87.62
88.91
91.22
92.91
95.94
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Uub
(223)
(226)
(227)
€
58
Ce
59
Pr
(262)
60
Nd
(263)
61
Pm
(262)
62
Sm
(265)
63
Eu
(266)
64
Gd
(269)
65
Tb
(272)
66
Dy
(277)
67
Ho
173.0
175.0
101
Md
102
No
103
Lr
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
(243)
(247)
(247)
(251)
(252)
(2??)
168.9
144.2
(244)
116
Uuh
69
Tm
91
Pa
(237)
(209)
68
Er
140.9
238.0
209.0
(2??)
90
Th
231.0
207.2
114
Uuq
140.1
232.0
Page 2
(261)
204.4
(257)
(258)
70
Yb
(259)
(210)
(222)
118
Uuo
(2??)
71
Lu
(262)
Not valid as a USNCO National Exam after May 1, 2007
DIRECTIONS
 When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
 There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
 Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which absorbs gaseous carbon dioxide most effectively?
(A) solid KOH
(B) solid SiO 2
(C) aqueous HCl
(D) aqueous NaF
2. Which laboratory results will tell whether an unknown
white solid is NaOH or NH4NO3?
6. When a liquid is delivered from a volumetric pipet a
small amount is typically retained in the tip. How should
a student proceed in order to deliver the volume of liquid
stated on the pipet?
(A) Leave the small amount in the tip.
(B) Use a pipet bulb to expel the remaining droplet.
(A) NaOH is soluble in H 2O but NH4NO3 is not.
(C) Shake the pipet to dispense the amount left in the
tip.
(B) Aqueous NaOH turns litmus blue but NH4NO3 does
not.
(D) Draw the liquid above the line initially to
compensate for the amount that remains in the tip.
(C) Aqueous NaOH reacts with copper metal but
NH4NO3 does not.
(D) NaOH gives a green flame test but NH4NO3 is
colorless in a flame.
3. Which sets of chemicals, when mixed, produce the
observation(s) listed?
Combination
Observation
I. NH4Cl(s) and H2O(l)
endothermic
II. 9 M H2SO4(aq) and H2O(l)
exothermic
III. 1M NaOH(aq) and 1 M HCl(aq)
exothermic
(A) III only
(B) I and II only
(C) II and III only
(D) I, II and III
4. What happens when 6 M nitric acid is added to an
aqueous solution that contains 0.1 M Cl– and
0.1 M Ag(NH3)2+?
(A) A deposit of silver metal forms.
(B) A precipitate of AgCl forms.
(C) Chlorine gas is released.
(D) Gaseous ammonia is released.
5. A mixture of which 0.2 M aqueous solutions will form a
precipitate that dissolves in 6 M nitric acid?
(A) Co(NO3)2 and NH4Cl
(B) Pb(NO 3)2 and NaBr
(C) Ba(NO3)2 and Na2CO3
(D) Al(NO3)3 and K 2SO4
Not valid as a USNCO National Exam after May 1, 2007
7. What is the molarity of a 0.500 molal aqueous solution of
calcium nitrate that has a density of 1.045 g·mL-1?
(A) 0.483 M
(B) 0.500 M
(C) 0.522 M
(D) 0.567 M
8. What volume of 0.150 M H2SO4 would be required to
completely neutralize a mixture of 20.0 mL of
0.200 M NaOH and 40.0 mL of 0.0500 M Ca(OH)2?
(A) 20.0 mL
(B) 26.7 mL
(C) 40.0 mL
(D) 53.3 mL
9. A compound with the formula X2O5 contains 34.8%
oxygen by mass. Identify element X.
(A) arsenic
(B) carbon
(C) phosphorous
(D) samarium
10. A solution of 0.0400 mol of C2H4Br2 and 0.0600 mol of
C 3H6Br2 exerts a vapor pressure of 145.4 mm Hg at a
certain temperature. Determine the vapor pressure of pure
C 3H6Br2 at this temperature. Assume the vapor pressure
of C 2H4Br2 at this temperature is 173 mm Hg and that the
solution obeys Raoult's Law.
(A) 76.2 mm Hg
(B) 118 mm Hg
(C) 127 mm Hg
(D) 138 mm Hg
Page 3
11. When 0.1 M aqueous solutions of aluminum nitrate,
magnesium nitrate, sodium nitrate and urea, (NH2)2CO,
are arranged in order of increasing boiling point, which
order is correct?
(A) Al(NO3)3 = Mg(NO3)2 = (NH2)2CO = NaNO 3
(B) Mg(NO 3)2 < (NH2)2CO < NaNO 3 < Al(NO3)3
(C) (NH2)2CO < NaNO 3 < Mg(NO3)2 < Al(NO3)3
(D) NaNO3 < Mg(NO 3)2 < Al(NO3)3 < (NH2)2CO
12. What is the maximum mass
Molar Mass / g·mol–1
of Ba3(PO4)2 that can be
Ba3(PO4)2
601.84
formed from
Na 3PO4
163.94
0.00240 mol of Ba(NO3)2 and 0.131 g of Na3PO4?
(A) 0.240 g
(B) 0.480 g
(C) 1.44 g
(D) 7.22 g
13. Which segment
of the heating
curve obtained at
constant
pressure
corresponds to
the transition
denoted by the
arrow in the
phase diagram?
16. The vapor pressure of phosphorus trichloride is
100 mm Hg at 21.0˚C and its normal boiling point is
74.2˚C. What is its enthalpy of vaporization in kJ. mol –1?
(A) 0.493
(B) 3.93
(C) 23.0
(D) 32.4
17. If the absolute temperature of a sample of gas is
increased by a factor of 1.5, by what ratio does the
average molecular speed of its molecules increase?
(A) 1.2
(B) 1.5
(C) 2.2
(D) 3.0
18. The curves in
the
accompanying
diagram
represent the
PV/RT
behavior of
the gases: He,
CH4 and C3H8.
Which
assignment of
behavior to gas is correct?
(A) 1 = He
(B) 1 = C3H8
2 = CH 4
2 = CH 4
3 = C3H8
3 = He
(C) 1 = CH 4
(D) 1 = C3H8
2 = C3H8
2 = He
3 = He
3 = CH 4
19. Calculate the standard enthalpy of formation of acetylene
(in kJ. mol –1).
2C 2H2(g) + 5O 2(g)
r 4CO2(g) + 2H 2O(l) ∆H˚ = –2243.6 kJ
C(s) + O2(g) r CO2(g)
∆H˚ = –393.5 kJ
H2(g) + 1/2 O2(g) r H 2O(l)
∆H˚ = –285.8 kJ
(A) a
(B) b
(C) c
(D) d
14. What is the molar mass of a gas that has a density of
5.66 g. L–1 at 35˚C and 745 mm Hg?
(A) 127
(B) 141
(C) 143
(D) 146
15. Consider the solids: body-centered cubic (bcc), facecentered cubic (fcc), simple cubic (sc) (or primitive),
constructed of spheres of the same size. When they are
arranged in increasing order of the percentage of free
space in a unit cell, which order is correct?
(A) fcc, bcc, sc
(B) bcc, sc, fcc
(C) sc, fcc, bcc
(D) bcc, fcc, sc
(A) 49.0
(C) 1121.8
(D) 1564.3
20. The boiling point of diethyl ether is 34.6˚C. Which is true
for the vaporization of diethyl ether at 25.0˚C?
(A) ∆G˚vap > 0
(B) ∆H˚vap < 0
(C) K vap = 1
(D) ∆S˚ vap < 0
21. Estimate the Bond Dissociation Enthalpies / kJ. mol–1
enthalpy of
C–C
350
C–O
350
combustion
C–H
410
C=O
732
of methane in
O–H
460
O–O
180
kJ. mol –1.
O=O
498
CH4(g) + 2O2(g) r CO2(g) + 2H 2O(g)
(A) 668
Page 4
(B) 98.0
(B) 540
(C) –540
(D) –668
Not valid as a USNCO National Exam after May 1, 2007
22. Which reaction has a positive ∆S˚reaction?
28. For the reaction A r B that is first-order in A, the rate
constant is 2.08×10–2 s–1. How long would it take for [A]
to change from 0.100 M to 0.0450 M?
(A) Ag+(aq) + Br–(aq) r AgBr(s)
(B) 2C 2H6(g) + 3O 2(g) r 4CO2(g) + 6H 2O(l)
(A) 0.0166 s
(C) N2(g) + 2H 2(g) r N 2H4(g)
(D) 2H2O2(l) r 2H 2O(l) + O2(g)
23. For reactions
I. constant number of moles
conducted at constant
II. constant temperature
pressure, under what
III. constant volume
conditions are ∆E and ∆H equal?
(A) I only
(B) II only
(C) III only
(D) I and II only
24. For the reaction,
H2(g) + I 2 (g) s 2HI(g)
K p = 50.0 at 721 K. What is the value of ∆G˚ for this
reaction (per mole of H2) at 721 K?
(A) –32.3 kJ
(B) –23.5 kJ
(C) –10.2 kJ
(D) –0.231 kJ
25. Which of these factors
affect the value of the
rate constant for a
reaction?
I. temperature
II. reactant concentration
III. use of a catalyst
(A) I only
(B) II only
(C) I and III only
(D) I, II and III
26. Which is the correct exponential form of the Arrhenius
equation?
(A) E = Ae
a
(C) k = Ae
–k
– RT
RT
(B)
Ea
(D) k = Ae
Ea = Ae
k
– Ea
RT
RT
€27. For the reaction A r B, €
€
what is the order with
respect to A that gives
this graph?
(A) zero
(C) 38.4 s
(D) 107 s
29. These data were obtained for the reaction: X + Y r Z.
X (M) Y (M) Rate: ∆Z/∆t / M·min–1
1.00
1.00
2.36×10-4
2.00
2.00
1.89×10-3
2.00
4.00
3.78×10-3
What is the rate law?
(A) Rate = k[X][Y]
(B) Rate = k[X]2[Y]
(C) Rate = k[X][Y]2
(D) Rate = k[X]2[Y]2
30. A possible mechanism for the conversion of ozone to
oxygen in the upper atmosphere is
O3(g) s O2(g) + O(g)
(fast equilibrium)
O(g) + O 3(g) s 2O2(g)
(slow)
Which rate law is consistent with this mechanism?
(A) Rate = k[O3]
(B) Rate = k[O3]2
(C) Rate = k[O3][O]
(D) Rate = k[O3]2[O 2]–1
31. A 0.050 M solution of an unknown acid is 1.0% ionized.
What is the value of its K a ?
(A) 2.5×10–7
(B) 5.0×10–6
(C) 5.0×10–4
(D) 5.0×10–2
32. Which mixture(s) form(s) buffer solutions?
I. 100 mL of 0.200 M HF
and 200 mL of 0.200 M NaF
II. 200 mL of 0.200 M HCl
and 200 mL of 0.400 M CH3CO2Na
III. 300 mL of 0.100 M CH3CO2H
and 100 mL of 0.300 M CH3CO2Na
(A) I only
(B) III only
(C) II and III only
(D) I, II and III
33. Determine the equilibrium constant for the reaction:
HF(aq) + NH3(aq) s NH4+(aq) + F–(aq)
given the equilibrium constants for the reactions.
K a = 6.9×10–4
HF(aq) + H2O(l) s H3O+(aq) + F–(aq)
+
–
NH3(aq) + H2O(l) s NH4 (aq) + OH (aq) K b = 1.8×10–5
K w = 1.0×10–14
2H2O(l) s H3O+(aq) + OH–(aq)
€
(B) first
(B) 16.7 s
(C) second
(D) third
Not valid as a USNCO National Exam after May 1, 2007
(A) 1.2×10–8
(B) 1.2×106
(C) 8.1×107
(D) 3.8×1015
Page 5
34. Calculate the pH of a
0.15 M solution of HOCl.
(A) 3.77
Ka
HOCl
(B) 4.18
(C) 6.71
2.9×10
–8
(D) 8.36
35. For which reaction does K p = K c ?
(A) 2C(s) + O2(g) s 2CO(g)
42. According to the tabulated standard reduction potentials
E˚ = –2.38 V
Mg2+(aq) + 2e– r Mg(s)
–
–
2H2O(l) + 2e r H2(g) + 2OH (aq) E˚ = –0.83 V
E˚ = 0.53 V
Br2(l) + 2e – r 2Br– (aq)
+
E˚ = 1.23 V
O2(g) + 4H (aq) r 2H 2O(l)
what products are formed during the electrolysis of an
aqueous MgBr2 solution?
(B) N2(g) + 3H 2(g) s 2NH3(g)
(A) Mg and H 2
(B) H2 and Br2
(C) 2H2(g) + O2(g) s 2H2O(g)
(C) H2 and O 2
(D) Mg and O 2
(D) H2(g) + I2(g) s 2HI(g)
36. CaF 2 has a Ksp = 3.9×10–11 at 25˚C. What is the [F–] in a
saturated solution of CaF2 at 25˚C?
(A) 2.1×10-4
(B) 3.4×10-4
(C) 4.3×10-4
(D) 6.8×10-4
(B) 2 / 1
(C) 3 / 1
(D) 5 / 1
38. Which change could occur at the anode of an
electrochemical cell?
(A) Cl– r Cl2
(B) H2O r H2
+
(C) Na r Na
(D) O2 r H2O
39. E˚ = 0.93 V
for the reaction:
Standard Reduction Potential / E˚
Fe2+(aq) + 2e– r Fe(s)
–0.41 V
Fe(s) + 2M+(aq) r Fe2+(aq) + 2M(s).
What is the standard potential for M+ + e– r M?
(A) 0.26 V
(B) 0.52 V
(C) 0.67 V
(D) 1.34 V
3+
(A) V (aq) r V (aq) + e
(B) VO3- + 2H+ r VO2+ + H2O
(C)
Mg
12
(D)
Ar
18
n
l
ml
ms
(A) 1
0
0
–_
(B) 2
2
1
_
(C) 3
1
1
_
(D) 4
3
–3
_
45. Which change(s) in electron structure occur when a gas
phase Mn atom is converted to a Mn2+ ion in the gas
phase?
I. The number of occupied energy levels decreases.
II. The number of half-filled orbitals decreases.
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
(A) F, Ne, Na
(B) Al, Mg, Na
(C) Sr, Ca, Mg
(D) Cl, Br, I
–
47. How many unpaired electrons are in a gas phase Co2+ ion
in its ground state?
(C) VO2+ + 2H+ + e – r V3+ + H2O
(A) 2
(D) VO2+ + H2O r VO2+ + 2H+ + e–
41. A solution of aqueous CuSO4 is electrolyzed with a 1.50
ampere current for 30.0 minutes. What mass of copper
metal is deposited?
(A) 0.889 g
(B) 6C
46. Which list gives the symbols of the elements in the order
of increasing first ionization energy?
40. For which half-reaction will a 1.0 unit increase in pH
cause the greatest increase in half-cell potential?
2+
(A) 5B
44. Which set of quantum numbers is NOT allowed?
37. When the reaction: Cl– + ClO3– r Cl2 + H2O is balanced
in acid solution what is the ratio of Cl– to ClO3–?
(A) 1 / 1
43. Which is the symbol for an element whose ground state
atoms have the same total numbers of s electrons and p
electrons?
(B) 1.19 g
(C) 1.78 g
(D) 3.56 g
(B) 3
(C) 4
(D) 5
48. The energy required to ionize a potassium ion is
419 kJ⋅mol –1. What is the longest wavelength of light that
can cause this ionization?
(A) 285 nm
(B) 216 nm
(C) 200 nm
(D) 107 nm
49. Which species has the same electron distribution around
the central atom as SiF 4?
(A) SF4
Page 6
(B) XeF4
(C) ClF4+
(D) BF4–
Not valid as a USNCO National Exam after May 1, 2007
50. Which is/are polar species?
I. SF2
II. SF4
III. SF6
(A) I only
(B) III only
(C) I and II only
(D) II and III only
51. According to the Lewis dot
structure for ozone, what is
the formal charge on the
central oxygen atom?
(A) –2
(B) –1
55. How many unsaturated compounds have the formula
C 4H8?
(A) 3
(B) 4
(C) 5
(D) 6
56. Which compound is least soluble in water?
(A) CH3CH2CH2F
(B) CH3CH2CH2NH2
(C) CH3CH2CH2OH
(D) CH3CH2CH2COOH
57. Which method for characterizing organic compounds
relies on the vibration of atoms in the compound?
(C) 0
(D) +1
52. When the species are arranged in order of increasing
length of the carbon-oxygen bond, which order is
correct?
(A) Na 2CO3 < HCO2Na < CH3ONa
(A) infrared spectroscopy
(B) nuclear magnetic resonance spectroscopy
(C) UV-visible spectroscopy
(D) X-ray diffraction
58. Which substance reacts most rapidly with water?
(B) CH3ONa < HCO 2Na < Na2CO3
(C) HCO2Na < Na2CO3 < CH 3ONa
(A) C 6H5Cl
(B) (CH3)3CCl
(D) Na 2CO3 < CH 3ONa < HCO 2Na
(C) (CH3)2CHCH2Cl
(D) CH3CH2CH2CH2Cl
53. Which ionic solid would require the most energy to form
gaseous ions?
(A) NaF
(B) Na 2O
(C) MgO
(D) MgF2
54. Solid calcium occurs as either cubic closest packing or
hexagonal closest packing. What is the most significant
difference between these two structures?
(A) the placement of layers of calcium atoms
(B) the distance betweeen calcium atoms in a single
layer
(C) the distance between calcium atoms in adjacent
layers
59. What type of compound is formed by the mild oxidation
of 2-pentanol?
(A) acid
(B) aldehyde
(C) ester
(D) ketone
60. Which species is lost during the formation of a
disaccharide from a monosaccharide?
(A) CH2
(B) CH2O
(C) CH2OH
(D) H2O
END OF TEST
(D) the coordination number of the calcium atoms in a
single layer
Not valid as a USNCO National Exam after May 1, 2007
Page 7
NATIONAL OLYMPIAD PART I
2007
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
A
B
D
B
C
A
A
B
A
C
C
A
B
D
A
D
A
B
A
A
D
D
C
B
C
D
C
C
B
D
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Property of the ACS DivCHED Examinations Institute
Answer
B
D
B
B
D
C
D
A
B
D
A
B
C
B
A
C
B
A
D
C
D
C
C
A
B
A
A
B
D
D
2007 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated High School, IL
David W. Hostage, Taft School, CT
Peter E. Demmin (retired), Amherst Central High School, NY
Marian Dewane, Centennial High School, ID
Adele Mouakad, St. John’s School, PR
Jane Nagurney, Scranton Preparatory School, PA
Kimberly Gardner, United States Air Force Academy, CO,
Preston Hayes, Glenbrook South High School, IL
Ronald O. Ragsdale, University of Utah, UT
Todd Trout, Lancaster Country Day School, PA
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for
a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until May 1, 2007, after which
tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the
front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to
use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this
examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials,
scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
Page 1
!
!
1. (12%) Compound X contains 2.239% hydrogen, 26.681% carbon
and 71.080 % oxygen by mass. The titration of 0.154 g of this
compound with 0.3351 M KOH produces the curve shown.
a.
b.
c.
d.
e.
a)
convert masses to moles:
# 1 mol &
2.239 g H " %
( = 2.221 mol (÷2.157) = 1.03
$ 1.008 g '
# 1 mol &
26.681 g C " %
( = 2.157 mol (÷2.157) = 1
$ 12.011 g '
# 1 mol &
71.08 g O " %
( = 4.443 mol (÷2.15) = 2.06
$ 16.00 g '
These numbers are close enough to whole numbers that the empirical formula must be CHO2
b) Obtain molar mass from titration (estimate endpoint at 10.4 mL)
!
!
!
!
!
Determine the empirical formula of the compound.
Calculate its molar mass and give its molecular formula.
When K2Cr2O7 is reacted with X in acidic solution the
products are chromium(III) ions and carbon dioxide.
Describe the color change that accompanies this
reaction.
Write a balanced ionic equation for this reaction.
Find the volume of dry carbon dioxide that could be
collected at 22 ˚C and 738 mm Hg when 0.839 g of
compound X is reacted with an excess of K2Cr2O7.
Mol NaOH = 0.3351 mol/L " 0.0104 L = 0.00348 mol
molar mass = 0.154 g ÷ 0.00348 mol = 44.2 g/mol ( " 2 because titration curve is diprotic) = 88.4 g/mol
The molar mass is 88.4 g/mol, the empirical formula molar mass is 45.02. This value is close to half the value of the
experimentally determined molar mass, so the molecular formula must be C2H2O4.
c)
3+
The color change will be from orange for Cr2O 2–
7 to green for Cr .
+
3+
d) The balanced ionic equation is: Cr2O 2–
+ 6CO 2 + 7H 2O
7 + 3H 2C 2O 4 + 8H " 2Cr
e)
!
Do the stoichiometry for oxalic acid to carbon dioxide, then calculate volume using ideal gas law.
# 1 mol H !C O & # 6 mol H C O &
2 2 4
2 2 4
0.839 g H 2C 2O 4 " %
( "%
( = 0.0186 mol H 2C 2O 4
90.04
g
H
C
O
3
mol
H
$
$
2 2 4'
2C 2O 4 '
(
)
–1
–1
nRT ( 0.0186 mol) 0.0821 L " atm " mol " K ( 295 K)
V=
=
= 0.464 L
$
$ 1 atm ''
P
& 738 mmHg # &
))
% 760 mmHg ((
%
2.
(15%) Coffee cup calorimetry experiments can be used to obtain ∆H f˚ for magnesium oxide.
a. Write a balanced equation for the formation of magnesium oxide, whose enthalpy change is ∆H f˚.
b. To determine the heat capacity of the calorimeter, 49.6 mL of 1.01 M HCl are reacted with 50.1 mL of 0.998 M NaOH. The
solution's temperature increases by 6.40˚C. Determine the heat capacity of the calorimeter. You may assume the solution's
specific heat capacity is 4.025 J·g–1 ·˚C–1 and the enthalpy of neutralization is –55.9 kJ per mole of H2O.
c. When 0.221 g of magnesium turnings are added to 49.9 mL of 1.01 M HCl and 49.7 mL of H2O in the same calorimeter, the
temperature increases by 9.67˚C. Write a balanced equation for the reaction that occurs and calculate the ∆H per mole of
Page 2
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
d.
e.
magnesium. (Assume the solution's specific heat capacity is 3.862 J·g–1 ·˚C–1 and the calorimeter constant is the value
obtained in b.)
When 0.576 g of MgO react with 51.0 mL of 1.01 M HCl and 50.1 mL of H2O in the same calorimeter the temperature rises
4.72˚C. Write a balanced equation for this reaction and calculate its ∆H per mole of MgO using the same assumptions as in
part c.
Use the above results and ∆Hf˚ of H2O(l) (–285.8 kJ·mol–1) to calculate ∆H f˚ of magnesium oxide.
a) Mg(s) + 1 2 O 2 (g) " MgO(s)
b) First, determine the limiting reactant:
Mol HCl = 1.01 mol/L " 0.0496 L = 0.00501 mol HCl
Mol NaOH = 0.998 mol/L " 0.0501 L = 0.0500 mol NaOH , so because it is a 1:1 stoichiometry, NaOH is limiting.
!
Via the enthalpy from the neutralization reaction, HCl(aq) +NaOH(aq) " NaCl(aq) +H 2O(l) #H = –55.9 kJ/mol we can calculate,
# –55.9 kJ &
0.0500 mol NaOH " %
( = – 2.795 kJ
$ 1 mol NaOH '
! so the rest is taken up by the calorimeter:
Account for heat taken up by the solution,
Total volume of solution is 49.6 mL + 50.1 mL = 99.7 mL (no information is provided about density, so the simplest assumption is to
use 1.00 g) so we have 99.7 g solution. Using the given specific heat capacity the heat absorbed by the solution is,
!
!
!
heat = 99.7 g " 4.025 J # g –1#o C –1 " 6.40 o C = 2568 J (heat absorbed by the solution)
Now we can calculate the heat absorbed by the calorimeter: 2795 J – 2568 J = 227 J absorbed by the calorimeter.
So the heat capacity of the calorimeter is 227 J / 6.40 oC = 35.5 J· oC–1
!
c) The reaction of magnesium with an acid is: Mg + 2H + " Mg 2+ + H 2
Total mass is: 99.6 g solution + 0.221 g Mg = 99.821 g
Total heat is heat absorbed by solution + heat absorbed by calorimeter:
heat solution = 99.821 g " 3.862 J #!g –1#o C –1 " 9.67 o C = 3728 J
heat calorimeter = 35.5 J"o C –1 # 9.67 o C = 343 J
Total heat = 3728 J + 343 J = 4071 J
#
&(
This is heat given off by 0.221 g Mg (using molar mass): 0.221 g Mg " %1 mol
24.31 g' = 0.00909 mol Mg
$
–4071 J
Thus,
= –4.479 " 10 5 J # mol–1 = –447.9 kJ # mol–1
0.00909 mol
!
!
!
d) The reaction is: MgO + 2H + " Mg 2+ + H 2O
!
#
&(
First determine moles reacted: 0.576 g MgO " %1 mol
40.31 g' = 0.0143 mol MgO
$
Once again,
! total heat is heat absorbed by solution + heat absorbed by calorimeter: (and solution mass includes MgO)
heat solution = 101.676 g " 3.862 J # g –1#o C –1 " 4.72 o C = 1853 J
! J"o C –1 # 4.72 o C = 168 J
heat calorimeter = 35.5
Total heat = 1853 J + 168 J = 2021 J
–2021 J
Thus,
= –1.413 " 10 5 J # mol–1 = –141.3 kJ # mol–1
0.0143 mol
!
!
!
!
!
e) Now construct a series of reactions that when summed are the formation reaction for MgO:
!
Mg 2+ + H 2O " MgO + 2H +
Mg + 2H + " Mg 2+ + H 2
#H = 141.3 kJ $ mol–1
#H = –447.9 kJ $ mol–1
Summed: These reaction yield:
Mg + H 2O " MgO + H 2
#H = –306.6 kJ $ mol–1
Now combine this reaction with the heat of formation for water to yield the desired result:
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
Page 3
!
Mg + H 2O " MgO + H 2
#H = –306.6 kJ $ mol–1
H 2 + 1 2 O 2 " H 2O
#H = –285.8 kJ $ mol–1
#H = –592.4 kJ $ mol–1
Mg + 1 2 O 2 " MgO
!
!
!
3. (13%) Hydrogen sulfide, H2S, is a weak acid that can be used to precipitate metal ions from solution selectively by controlling
the pH.
Acid Ionization Constants, H2S
K1
5.7×10–8
K2
1.3×10–13
a. Write equations to represent each of the ionization steps of H 2S.
Ksp
b. Write an equation to represent the overall ionization of H2S to form S2– and 2H+
Bi
S
1.6×10–72
2 3
and calculate the equilibrium constant for this process.
MnS
3.0×10–11
c. For a solution with [H2S] = 0.10 M, with [Bi3+] = [Mn2+] = 1.5 mM and
[H+] = 10 mM, give the formula for the metal sulfide which precipitates first and calculate the percentage of it that will
remain in solution at equilibrium.
d. The pH of the solution is raised until the other metal sulfide begins to precipitate. Determine the pH of the solution at which
the second metal sulfide begins to precipitate.
a)
H 2S " H + + HS –
K1 = 5.7 # 10 –8
HS – " H + + S 2–
K 2 = 1.3# 10 –13
b)
H 2S " 2H + + S 2–
K = 7.4 # 10 –21
!
c) Calculate sulfide ion concentration:
!
!
2
K=
[ H 2S]
2
( 0.010) [S 2– ]
= 7.4 " 10 –21 so [S2–] = 7.4×10–18
( 0.1)
Now calculate Q and compare to K for each cation (with sulfide):
Bismuth: Ksp = [Bi3+]2[S2–]3 = 1.6×10–72
Q = (1.5×10–3)2(7.4×10–18)3 = 9.1×10–58
Q > Ksp so there will be a precipitate formed.
Manganese: Ksp = [Mn2+][S2–] = 3.0×10–11
Q = (1.5×10–3)(7.4×10–18) = 1.1×10–20
Q < Ksp so there will not be a precipitate formed.
Thus – the bismuth is the first metal sulfide to precipitate. Now to calculate what percentage will remain in solution:
3+ 2
[Bi ]
=
K sp
2– 3
[S ]
(1.6 " 10 )
(7.4 " 10 )
–72
=
= 3.95" 10 –21 so [Bi3+] = 6.3×10–11
–18 3
The percentage can be calculated using the ratio of the amount remaining in solution divided by the original amount:
!
!
[H + ] [S2– ] =
(6.9 " 10 ) " 100 = 4.2 " 10
(1.5" 10 )
–11
%=
–3
–6
%
d) first determine the concentration of sulfide that will result in precipitation:
Page 2
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
[S ] =
2–
+ 2
[H ]
!
4.
=
K sp
[Mn
2+
(3.0 " 10 ) = 2.0 " 10
=
] (1.5" 10 )
K [ H 2S]
[S ]
2_
–11
–3
–8
Now plug this value into the equation for K from Part (c):
(7.4 " 10 )(0.10) = 3.7 " 10
(2.0 " 10 )
–21
=
–14
–8
(10%) A galvanic cell is based on the half-reactions;
Cr3+ + 3e– r Cr
!
2+
E˚ = –0.744 V
–
Ni + 2e r Ni
a.
b.
c.
d.
e.
f.
!
!
!
!
!
and [H+] = 1.92×10–7 so pH = 6.7
E˚ = –0.236 V
Write the balanced equation for the overall cell reaction.
State which electrode increases in mass as the cell operates. Explain your answer.
Calculate E˚cell
Determine the value of ∆G˚ for the cell reaction at 25˚C.
Calculate the value of K for the cell reaction at 25˚C.
Find the voltage of the cell at 25˚C if [Cr3+] and [Ni2+] are both changed to 0.010 M.
a) 2Cr + 3Ni2+ " 2Cr 3+ + 3Ni
b) The nickel electrode increases in mass as the cell operates because Ni2+ ions in solution are reduced there (it is the cathode) and
are deposited as Ni(s).
c) E ocell = E ored + E oox = –0.236 V + 0.744 V = 0.508 V
d) "G o = –nFE = –(6 mol)(96500 J # mol–1 # V–1)(0.508 V) = – 294000 J = – 294 kJ
"G o = –RT ln K
e) –294100 J = – (8.314 J # mol–1 # K –1)(298 K) lnK
or
K = 10
nE o
0.592
= 10
3.048
0.592
= 3.1" 10 51
ln K = 118.7 and K = 3.62 $ 10 51
2%
"
0.0257 $ [.01] '
0.0257
f) E = E o –
ln$
= 0.508 V –
ln(100
! ) V = 0.508 – 0.0197 V = 0.488 V
3'
6
6
# [.01] &
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Solid ammonium chloride and solid calcium hydroxide are mixed.
b. Excess carbon dioxide gas is bubbled into a sodium hydroxide solution.
c. Sodium sulfite is added to a neutral potassium permanganate solution.
d. Aqueous hydrofluoric acid is placed on a piece of silica.
e. Chlorobenzene is heated with a mixture of concentrated nitric and sulfuric acids.
f. Iodine-131 undergoes radioactive decay.
a) NH 4 Cl + Ca(OH) 2 " CaCl2 + NH 3 + H 2O
b) CO 2 + OH – " HCO –3
!
–
2–
c) SO 2–
3 + MnO 4 " SO 4 + MnO 2
d) HF + SiO 2 " SiF4 + H 2O
!
!
!
e) C 6 H 5Cl + H + + NO –3 "
or
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
!
Page 5
f)
!
131
53 I
"
0
–1#
+ 131
54 Xe or
131
53 I
" 42 He + 12751Sb (for half credit)
6. (12%) The reaction of NO with O2 to give NO2 is an important step in the commercial production of HNO3.
a. Describe an experiment to measure the rate of this reaction.
b. If the rate equation is found to be Rate = k[NO]2[O2], give the effect on the rate of tripling the concentration of.
i. NO
ii. O2
c. These two mechanisms have been proposed for this reaction,
I
2NO + O2 r 2NO2
II 2NO r N2O2
N2O2 + O2 r 2NO2
i. State and explain which of the two mechanisms is more likely.
ii. State and explain which of the two steps in mechanism II must be the slow step if this mechanism is to be consistent with
the rate law in b.
a) The stoichiometry of the reaction is: 2NO(g) + O2(g) r 2NO2(g) so there is a change in the number of moles of gas as the
reaction goes forward. This reaction can be monitored by measuring the total pressure of the system as a function of time.
Alternatively, the appearance of the red color of NO2 can be measured (e.g. with a spectropohotometer.)
b) Because the rate law is Rate = k[NO]2[O2], tripling the concentration of NO will cause the rate to increase by a factor of 9.
Tripling the concentration of O2 will cause the rate to increase by a factor of 3.
c) i. Mechanism is II more likely because mechanism one involves a trimolecular collision. Such a collision is uncommon. By
contrast, II has a pair of bimolecular reactions which are considerably more likely to occur.
ii. Step 2 must be the slow step because it would have a rate law of Rate=k[N2O2][O2], but N2O2 is an intermediate whose
concentration arises from the first step. With a slow second step, the first step achieves equilibrium, so [N2O2] = K[NO]2 and
the overall rate law would be, Rate = k[NO]2[O2].
7.
(16%) Account for the following observations,
a. The bond angle in H2O (104.5˚) is greater than that of H2S (92˚) but less than that in Cl2O (110.8˚).
b. The bond dissociation energy of Cl2 (240 kJ·mol–1) is greater than that of F2 (154 kJ·mol–1) or Br2 (190 kJ·mol–1).
c. The boiling point of NH3 is higher (–33˚C) than that of NF3 (–129˚C) but lower than that of NCl3 (71˚C).
d. SiF4 is tetrahedral while SF4 has a see-saw shape and XeF4 is square planar.
a) The angle in H-O-H is greater than H-S-H because the bonding pairs in H-S-H are further from the S atom (the atomic orbitals
used in S have electron density that is further from the nucleus) so they can be forced closer together by the lone pair electrons
on the S. The Cl-O-Cl bond angle is larger than either of the because the Cl atoms are large which gives rise to steric
interference that forces them apart.
b) The bond in Cl2 is stronger than that in F2 because the F atoms are sufficiently small that the lone pairs on the F atoms repel one
another weakening the bond. The bond in Br2 is weaker than that of Cl2 because the obritals in the larger atom (Br) do not
overlap as efficiently.
c) The boiling point of NH3 is higher than that of NF3 because NH3 molecules can form hydrogen bonds with each other increasing
the attractive forces relative to the dispersion and dipole forces between the NF3 molecules. For NCl3, the dispersion forces are
sufficiently large (because of the large, polarizable Cl atoms) that they are stronger than the hydrogen bonding in NH3.
d) SiF4 is tetrahedral, with four bonding pairs about the central Si atom. SF4 has 5 pairs of electrons (4 bonding and 1 lone pair) so it
has a see-saw shape. XeF4 has 6 pairs of electrons (4 bonding and 2 lone pairs) so it has a square planar shape.
8.
(10%) There are six different isomers with the formula C4H8O2 containing a –CO2 group. When added to water two of the six
are substantially more soluble than the other four.
a. Write structural formulas for the two water-soluble compounds and outline how their structures lead to their greater
solubility.
b. State the name of the class of compounds represented by the other four isomers.
c. Draw structural formulas for any three of the four less soluble isomers.
d. Write an equation for the laboratory synthesis of one of these four isomers and name each of the reactants.
Page 2
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
a) The two more soluble forms have –COOH groups that make them more soluble because that group allows for hydrogen
bonding with water.
b) The less soluble isomers are esters.
c) Any 3 of these four structures could be shown.
d) An example reaction would be:
END OF PART II
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
Page 7
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
NA gas constant
h
R Planck’s constant
°C gram
P
g pressure
c heat capacity
k
Cp rate constant
C hour
Rf
h retention factor
E joule
s
J second
Ea kelvin
c
K speed of light
H kilo- prefix
T
k temperature, K
S liter
t
L time
volt
V
CONSTANTS
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
E = E! –
! k2 $ Ea ! 1 1 $
=
'
" k1 &% R #" T1 T2 &%
" –!H % " 1 %
ln K = $
' $ ' +c
# R & # T&
RT
ln Q
nF
ln #
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Uub
114
Uuq
116
Uuh
118
Uuo
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
(2??)
(2??)
Page 8
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO Olympiad National Exam after May 1, 2007.
2007 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Brookline High School, Brookline, MA
Chair
Linda Weber, Natick High School, Natick, MA
Jim Schmitt, Eau Claire North High School, Eau Claire, WI
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after May 1, 2007
Page 1
2007 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
You have been given two ionic solutions, 0.10 M unknown salt, MClx solution and 0.10 M sodium solution,
NazY. Devise and carry out an experiment to determine the identity of the unknown metal cation and the
unknown anion in these solutions. The possible cations are potassium, zinc, aluminum, or silver. The possible
anions are nitrate, carbonate, phosphate, or sulfide.
You should provide both quantitative and qualitative evidence to support your answers.
Lab Problem 2
LDPE (low density polyethylene, #4) is a petroleum-based polymer used to make flexible bottles, films, and
plastic containers. Given water, ethanol (density = 0.789 g·mL–1), and the equipment provided, devise and
carry out an experiment to precisely determine the thickness of the LDPE samples provided
Page 2
Not valid for use as an USNCO National Examination after May 1, 2007.
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after May 1, 2007
Examiner’s Initials:
Page 3
2. Record your data and other observations.
3. Calculations and Conclusions.
Page 4
Not valid for use as an USNCO National Examination after May 1, 2007.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after May 1, 2007
Examiner’s Initials:
Page 5
2. Record your data and other observations.
3. Calculations and Conclusions.
Your reported thickness of LDPE =
Page 6
Not valid for use as an USNCO National Examination after May 1, 2007.
PERIODIC TABLE OF THE ELEMENTS
2
He
1
H
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Uub
114
Uuq
116
Uuh
118
Uuo
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
(2??)
(2??)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after May 1, 2007
Page 7
2007 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III - ANSWERS
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Lab Problem #1
This problem involves knowledge of solubility rules and precipitation reactions. In addition, the
identification of the unknown cation and anion requires relating the volumes (drops) of the reacting
solutions to the quantity of precipitate produced and hence, to the molar ratios of the reacting ions.
Plan
The plan should include both an intention to gather qualitative information about the individual solutions
and the mixture and quantitative information related to the quantity of precipitate produced upon combining
the two solutions.
Qualitative observations
Both solutions are clear and colorless.
There is no odor from the anion solution.
When mixed, a white precipitate is formed.
No bubbles/gas is produced.
Quantitative observations
When the two solutions are mixed in test tubes so that the ratio of the cation and anion are varied in a
systematic manner the quantity of precipitate should be greatest in the tube with a 3:2 ratio of MClx:NazY.
Excellent Student Results
Student included a range of qualitative observations and reasoning based on them such as;
Clear MClx solution indicates the absence of Ag+ since AgCl is insoluble.
Lack of odor in NazY solution indicates the absence of S2-.
Appearance of precipitate indicates the absence of K+ and NO3- ions since all their compounds are soluble.
Lack of bubbles in NazY solution and upon mixing indicates absence of CO32-.
Possible cations are Zn2+ and Al3+ while the anion is most likely PO43-.
Student provided a clear explanation of the variation of the number of drops to determine the stoichiometry
ratio of MClx:NazY.
Student gave a clear data table with several trials to demonstrate the 3:2 ratio of MClx:NazY.
Student identified the cation as Zn2+ and the anion as PO43-.
Average Student Results
Page 8
Not valid for use as an USNCO National Examination after May 1, 2007.
Some qualitative information is given to demonstrate a knowledge of solubility and precipitate formation.
Student provided evidence of several combinations of the two solutions and may have inferred something
about the relationship between the solution ratio and identity of the salts.
Below Average Student Results
Little or no qualitative information was reported or used to make predictions about the identity of the
unknown cation and anion.
Student either did not report any quantitative information or was unable to use the quantitative information
acquired to infer any information about the reaction stoichiometry from it.
#########################
Lab Problem #2
Excellent Students Results:
Student proposed a clear, detailed procedure for determining the thickness of the LDPE sheet, recognizing
that measuring the volume of such a sheet directly would not be possible because of the small volume.
Excellent procedures invariably involved measuring the density of the plastic; good methods included
making a series of ethanol-water mixtures and interpolating the mixture of neutral buoyancy, or starting
with one liquid and adding the other until neutral buoyancy was achieved. Density of the neutrally buoyant
liquid was measured either by using the weighted average of ethanol or water, or by direct measurement of
the mass of a known volume of the liquid.
Student performed several buoyancy trials, either using a variety of water-ethanol volume ratios in a series
of standards, or by redetermining the point of neutral buoyancy. Results were clearly displayed in a data
table. Area and mass of LDPE piece(s) were measured in duplicate.
Calculations are clearly shown using proper unit measurements and significant figures in final answers.
Student demonstrated knowledge of the assumptions used in calculation (for example, the assumption of
additive volumes if density of the neutrally buoyant mixture was calculated rather than measured directly).
Final value for thickness was within 20% of the accepted value.
Average Student Results:
Measurement of density was proposed, but not clearly thought out; or, less precise procedures for
determining volume directly (e.g., by displacement of liquid in the graduated cylinder) were proposed.
Student made only qualitative (floats in water, sinks in ethanol) or grossly erroneous measurements of
density.
Only one trial was performed.
Final value for thickness was within 40% of the accepted value.
Below Average Student Results:
Procedure was vague or unintelligible.
Calculations were unclear or in error.
Final value for thickness was over 40% off from the accepted value.
Not valid for use as an USNCO National Examination after May 1, 2007
Page 9
2008 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM – PART 1
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, Chair, State University of New York, Cortland
Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired)
William Bond, Snohomish HS, Snohomish, WA
David Hostage, Taft School, Watertown, CT
Peter Demmin, Amherst HS, Amherst, NY (retired)
Marian Dewane, Centennial HS, Boise, ID
Valerie Ferguson, Moore HS, Moore, OK
Paul Groves, South Pasadena HS, Pasadena, CA
Adele Mouakad, St. John’s School, San Juan, PR
Jane Nagurney, Scranton Preparatory School, Scranton, PA
Ronald Ragsdale, University of Utah, Salt Lake City, UT
Kimberly Gardner, US Air Force Academy, Colorado Springs, CO
DIRECTIONS TO THE EXAMINER–PART I
Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 23, 2008, after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators.
DIRECTIONS TO THE EXAMINEE–PART I
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record
of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the
statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the
examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI.
All rights reserved. Printed in U.S.A.
ABBREVIATIONS AND SYMBOLS
A Faraday constant
F molal
atm formula molar mass
M molar
u free energy
G molar mass
A frequency
ν mole
NA gas constant
R Planck’s constant
°C gram
g pressure
c heat capacity
Cp rate constant
C hour
h retention factor
E joule
J second
Ea kelvin
K temperature, K
H kilo– prefix
k time
S
liter
L volt
K
milli– prefix
m
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi– prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
CONSTANTS
m
M
M
mol
h
P
k
Rf
s
T
t
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
1 atm = 760 mmHg
EQUATIONS
E = Eo !
1
1A
1
H
RT
ln Q
nF
"k % E " 1 1 %
ln$$ 2 '' = a $$ ( ''
# k1 & R # T1 T2 &
$ "#H '$ 1 '
ln K = &
)& ) + constant
% R (% T (
PERIODIC TABLE OF THE ELEMENTS
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
1.008
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Uub
114
Uuq
116
Uuh
118
Uuo
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
(2??)
(2??)
Page 2
!
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
DIRECTIONS

When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.

There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.

Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. Which substance has the highest melting point?
(A) Li2O
(B) MgO
2. Which reagents produce a
gas when combined?
(C) CO2
(D) N2O5
I. HCl and Na2SO3
II. NaOH and Al
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
6. A NaOH solution is to be standardized by titrating it
against a known mass of potassium hydrogen phthalate.
Which procedure will give a molarity of NaOH that is too
low?
(A) Deliberately weighing one half the recommended
amount of potassium hydrogen phthalate.
(B) Dissolving the potassium hydrogen phthalate in
more water than is recommended.
(C) Neglecting to fill the tip of the buret with NaOH
solution before titrating.
3. A 1:1 mixture of
pentane and hexane
is separated by
fractional distillation
in the apparatus
shown. At what
temperature does the
first drop of
condensate appear on
the thermometer?
(D) Losing some of the potassium hydrogen phthalate
solution from the flask before titrating.
7. Which solute is least soluble in water?
(A) 1-butanol
(B) ethanol
(C) methanol
(D) 1-propanol
8. The mass of a single molecule of an allotrope of sulfur is
3.20×10–22 g. How many sulfur atoms are present in a
molecule of this allotrope?
Boiling point / oC
pentane
36
hexane
69
(A) 4
(A) less than 36 ˚C
(B) 36 ˚C
(C) between 36 ˚C and 69 ˚C
(D) more than 69 ˚C
4. Which nitrogen halide is least stable thermodynamically?
(A) NF3
(B) NCl3
(C) NBr3
(D) NI3
5. Cyclohexane and water can be separated by using a
separatory funnel. Which property contributes to this
separation?
(A) Cyclohexane and water are immiscible.
(B) Cyclohexane has a lower viscosity than water.
(C) Cyclohexane has a greater molar mass than water.
(D) Cyclohexane has a greater vapor pressure than water.
(B) 6
(C) 8
(D) 12
9. 100. L of carbon dioxide measured at 740. mmHg and
50 ˚C is produced by the complete combustion of a
sample of pentane.
2C5H12 + 16O2 r 10CO2 + 12H2O
What mass of pentane reacted?
(A) 342 g
(B) 265 g
(C) 64.4 g
(D) 53.0 g
10. Which 0.10 M aqueous solution has the smallest change
in freezing point relative to pure water?
(A) HC2H3O2
(B) HCl
(C) CaCl2
(D) AlCl3
11. Magnetite, Fe3O4, can be
Molar Mass / g·mol–1
reduced to iron by heating
Fe3O4
232
with carbon monoxide according to the equation:
Fe3O4 + 4CO r 3Fe + 4CO2
What mass of Fe3O4 is required in order to obtain 5.0 kg
of iron if the process is 88% efficient?
(A) 6.1 kg
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
(B) 6.9 kg
(C) 7.9 kg
(D) 18 kg
Page 3
12. 40.0 g of a solute is dissolved in 500. mL of a solvent to
give a solution with a volume of 515 mL. The solvent has
a density of 1.00 g/mL. Which statement about this
solution is correct?
(A) The molarity is greater than the molality.
(B) The molarity is lower than the molality.
(C) The molarity is the same as the molality.
(D) The molarity and molality cannot be compared
without knowing the solute.
13. In the graph, the
natural log of the
vapor pressures of
two substances
are plotted versus
1/T. What can be
concluded about
the relative
enthalpies of
vaporization
(∆Hvap) of these substances?
18. The atoms in crystals of silver metal are arranged in a
cubic closest packed structure. What is the unit cell in
this structure?
(A) body-centered cubic
(B) face-centered cubic
(C) hexagonal-close packed
(D) simple cubic
19. Use the information provided to calculate the standard
enthalpy of formation of acetylene, C2H2(g), in kJ·mol–1.
C2H2(g) + 5/2O2(g)
r 2CO2(g) + H2O(l)
∆H˚ = –1299.5 kJ
C(s) + O2(g) r CO2(g)
∆H˚ = –393.5 kJ
H2(g) + 1/2O2(g) r H2O(l)
∆H˚ = –285.8 kJ
(A) –1978.8
(B) –1121.4
(C) 226.7
(D) 453.4
20. Which statement is always true for a spontaneous
reaction?
(A) The entropy change for the system is negative.
(B) The enthalpy change for the system is negative.
(C) The entropy change for the universe is positive.
(A) ∆Hvap of I is greater than ∆Hvap of II
(D) The free energy change for the system is positive.
(B) ∆Hvap of I is less than ∆Hvap of II
21. The heat of a reaction is measured in a bomb calorimeter.
This heat is equal to which thermodynamic quantity?
(C) ∆Hvap of I is is equal to ∆Hvap of II
(D) No conclusion can be drawn from this information
alone.
14. For which two gases are the rates of effusion 2:1?
(A) H2 and He
(B) He and O2
(C) Ne and Kr
(D) N2 and Ar
15. Which gas has a density of 0.71 g·L–1 at 0 o C and 1 atm?
(A) Ar
(B) Ne
(C) CO
(D) CH4
16. Supercritical carbon
dioxide exists at which
point on the
accompanying phase
diagram?
(A) ∆E
(B) ∆G
(C) ∆H
(D) ∆S
22. 84.12 g of gold at Specific heat capacities / J.g–1.˚C–1
120.1 ˚C is placed
Au(s)
0.129
in 106.4 g of H2O
H2O(l)
4.184
at 21.4 ˚C. What is the final temperature of this system?
(A) 70.8
(B) 65.0
(C) 27.8
(D) 23.7
23. In order to calculate the lattice energy of NaCl using a
Born-Haber cycle, which value is not needed?
(A) enthalpy of sublimation of Na(s)
(B) first ionization energy of Cl(g)
(C) bond dissociation energy of Cl2(g)
(D) enthalpy of formation of NaCl(s)
(A) A
(B) B
(C) C
17. Which properties increase
with an increase in
intermolecular forces at 25 ˚C?
(D) D
I. surface tension
II. vapor pressure
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
Page 4
24. Liquid bromine boils at 332.7 K.
Estimate the enthalpy of
formation of Br2(g) in kJ·mol–1.
(A) 7.40
(B) 12.1
S˚ / J .mol–1.K–1
Br2(g)
58.6
Br2(l)
36.4
(C) 19.5
(D) 22.2
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
25. A student analyzed the data
from a zero order reaction and
obtained the graph shown.
What labels should be attached
to the X and Y axes,
respectively?
(A) time, concentration
(B) time, 1 / concentration
(C) time, ln (concentration)
(D) 1/time, concentration
26. Under certain conditions the reaction of CO with NO2 to
give CO2 and NO results in the rate law:
rate = k[CO][NO2].
What are the units for the rate constant, k?
(A) mol.L–1.min–1
(B) L.mol–1.min–1
(C) mol2.L–2.min–1
(D) L2.mol–2.min–1
27. For the reaction: X + Y r Z, initial rate data are given in
the table.
[X] / M [Y] / M Rate / mol.L–1.s–1
0.10
0.10
0.020
0.10
0.20
0.080
0.30
0.30
0.540
What is the rate law for this reaction?
(A) Rate = k[X]2
(B) Rate = k[Y]2
(C) Rate = k[X][Y]
(D) Rate = k[X][Y]2
28. The rate of the reaction of chlorine gas with a liquid
hydrocarbon can be increased by all of the changes
except one. Which change will be ineffective?
(A) Use UV light to dissociate the Cl2.
30. For the reaction; A r B, the rate law is rate = k[A]. If the
reaction is 40.0% complete after 50.0 minutes, what is
the value of the rate constant, k?
(A) 8.00×10–3 min–1
(B) 1.02×10–2 min–1
(C) 1.39×10–2 min–1
(D) 1.83×10–2 min–1
31. When 2.00 mol each of H2(g) and I2(g) are reacted in a
1.00 L container at a certain temperature, 3.50 mol of HI
is present at equilibrium. Calculate the value of the
equilibrium constant, Kc.
(A) 3.7
(B) 14
(C) 56
(D) 2.0×102
32. For which equation is the equilibrium constant equal to
Ka for the ammonium ion, NH4+?
(A) NH4+(aq) + OH–(aq) s NH3(aq) + H2O(l)
(B) NH4+(aq) + H2O(l) s NH3(aq) + H3O+(aq)
(C) NH3(aq) + H2O(l) s NH4+(aq) + OH–(aq)
(D) NH3(aq) + H3O+(aq) s NH4+(aq) + H2O(l)
33. What is the pH of a solution prepared by mixing 45.0 mL
of 0.184 M KOH with 65.0 mL of 0.145 M HCl?
(A) 1.07
(B) 1.13
(C) 1.98
(D) 2.92
34. The gas phase reaction
shown is endothermic
as written. Which
change(s) will increase the quantity of CH3CH=CH2 at
equilibrium?
I. increasing the temperature
II. increasing the pressure
(B) Increase temperature at constant pressure.
(A) I only
(B) II only
(C) Divide the liquid into small droplets.
(C) Both I and II
(D) Neither I nor II
(D) Double the pressure by adding He gas.
29. One proposed mechanism of the reaction of HBr with O2
is given here.
HBr + O2 r HOOBr
(slow)
HOOBr + HBr r 2HOBr
(fast)
HOBr + HBr r H2O + Br2
(fast)
What is the equation for the overall reaction?
(A) HBr + O2 r HOOBr
(B) 2HBr + O2 r Br2 + H2O2
(C) 4HBr + O2 r 2H2O + 2Br2
35. The curve represents
the titration of a weak
monoprotic acid. Over
what pH range(s) will
the acid being titrated
serve as a buffer when
mixed with its salt?
I. pH 4 – 6
II. pH 7 – 9
III. pH 12 – 13
(A) I only
(B) II only
(C) I and III only
(D) I, II and III
(D) 2HOBr r 2H2O + Br2
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
Page 5
36. The pH of a saturated solution of Fe(OH)2 is 8.67. What
is the Ksp for Fe(OH)2?
(A) 5×10–6
(B) 2×10–11
(C) 1×10–16
(D) 5×10–17
37. In an operating voltaic cell electrons move through the
external circuit and ions move through the electrolyte
solution. Which statement describes these movements?
(A) Electrons and negative ions both move toward the
anode.
(B) Electrons and negative ions both move toward the
cathode.
(C) Electrons move toward the anode and negative ions
move toward the cathode.
(D) Electrons move toward the cathode and negative
ions move toward the anode.
38. The reduction potentials
I. Ao reduces B2+
for the +2 cations,
II. B2+ oxidizes Co
2+
–
o
III.
Bo oxidizes Do
e.g. A + 2e r A ,
of four metals decrease in the order A, B, C, D. Which
statement(s) is/are true?
(A) II only
(B) III only
(C) I and II only
(D) I and III only
Questions 39 and 40 should be answered with reference to
the reaction:
2Ag+(aq) + M(s) r M2+(aq) + 2Ag
E˚ = 0.940 V
42. A 3.00 amp current is used to electrolyze the molten
chlorides; CaCl2, MgCl2, AlCl3, and FeCl3. The
deposition of which mass of metal will require the
longest electrolysis time?
(A) 100 g Ca
(B) 50 g Mg
(C) 75 g Al
(D) 125 g Fe
43. Which set of quantum numbers corresponds to an
electron in a 4d orbital?
(A) n = 4, l = 1, ml = –1, ms = 1/2
(B) n = 4, l = 2, ml = –2, ms = –1/2
(C) n = 4, l = 3, ml = 3, ms = 1/2
(D) n = 4, l = 3, ml = –1, ms = –1/2
44. What is the energy of a photon from a laser that emits
light at 632.8 nm?
(A) 3.14×10–19 J
(B) 1.26×10–31 J
(C) 2.52×10–33 J
(D) 4.19×10–40 J
45. How many unpaired electrons are in a gaseous Co2+ ion
in its ground state?
(A) 1
(B) 3
(C) 5
(D) 7
46. Which ion is not isoelectronic with Ar?
(A) S2–
(B) K+
(C) Sc2+
(D) Ti4+
47. Which process releases the most energy?
(A) Mg2+ (g) + e– r Mg+(g)
39. What is the value of E˚
for the half reaction,
2+
E˚ / V
Ag (aq) + e r Ag(s) 0.799
+
–
–
M (aq) + 2e r M(s)?
(A) 0.658 V
(B) 0.141 V
(C) –0.141 V
(D) –0.658 V
40. Which change will cause the largest increase in the
voltage of a cell based on the reaction above?
(A) Doubling the [Ag+] from 1M to 2M
+
(C) Doubling the volume of the 1M Ag solution
(D) Reducing the [M2+] from 1M to 0.5M
41. If a voltaic cell has a positive Eo value, what can be
concluded about the values of ∆Go and Keq?
o
(B) ∆G < 0, Keq > 1
o
(D) ∆Go > 0, Keq > 1
(C) ∆G > 0, Keq < 1
Page 6
(C) Na2+(g) + e– r Na+(g)
(D) Na+(g) + e– r Na(g)
48. In which list are the ions arranged in order of increasing
size?
(A) F– < S2– < Al3+ < Mg2+
(B) F– < S2– < Mg2+ < Al3+
(C) Mg2+ < F– < Al3+ < S2–
(B) Doubling the amount of M(s)
(A) ∆G < 0, Keq < 1
(B) Mg+(g) + e– r Mg(g)
o
(D) Al3+ < Mg2+ < F– < S2–
49. Molecules with non-zero
dipole moments include
which of those listed?
I. H2C=CHCl
II. cis - ClHC=CHCl
III. trans - ClHC=CHCl
(A) I only
(B) III only
(C) I and II only
(D) I, II and III
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
50. Which species is diamagnetic?
(B) N2+
(A) NO
(D) O22–
(C) O2
51. What is the I-I-I bond angle in the I3– ion?
(A) 180o
(B) 120 o
(A) identity of the monomers in the two polymers.
(C) 90 o
o
(D) more than 90 but less than 120
(B) number of monomer units in the two polymers.
o
(C) orientation of the bonds joining the monomers.
52. Which species has the shortest nitrogen-oxygen bond?
(A) NO+
(B) NO2+
60. Cellulose and starch are biological polymers. Humans are
able to digest starch but not cellulose. This difference is
due primarily to a difference in the
(C) NO2–
(D) percentage of carbon in the two polymers.
(D) NO3–
53. Which substance will form hydrogen bonds to water
molecules but will not form hydrogen bonds with its own
molecules?
(A) HF
(B) C2H5OH
(C) CH3NH2
(D) CH3OCH3
END OF TEST
54. In the gas phase PCl5 exists as individual molecules but
in the solid it takes on the ionic structure PCl4+PCl6–.
What are the geometries of these three species
PCl4+
PCl5
PCl6–
(A) trigonal
bipyramidal
see-saw
octahedral
(B) trigonal
bipyramidal
tetrahedral
octahedral
(C) trigonal
bipyramidal
square planar
distorted
octahedral
(D) square
pyramidal
see-saw
square planar
55. Which molecule contains exactly eight carbon atoms?
(A) benzoic acid
(B) 2,3-dimethylhexane
(C) 3-ethylpentane
(D) 3-methyloctane
56. Which formula represents an alkyne?
(Assume all are noncyclic.)
(A) C2H2
(B) C2H4
(C) C5H10
(D) C8H18
57. How many compounds have the formula C2H3Cl3?
(A) 2
(B) 3
(C) 4
(D) 5
58. Which is a condensation polymer?
(A) polyethylene
(B) polyvinylchloride
(C) polystyrene
(D) polyethylene terephthalate
59. What is the number of pi (π) bonds in trans-butenedioic
acid (C4H4O4)?
(A) 1
(B) 2
(C) 3
(D) 4
Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008.
Page 7
Olympiad 2008 National Part I
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
B
C
B
D
A
C
A
B
D
A
C
B
A
C
D
C
A
B
C
C
A
D
B
A
A
B
D
D
C
B
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Answer
D
B
C
D
A
D
D
A
C
A
B
C
B
A
B
C
C
D
C
D
A
A
D
B
B
A
A
D
C
C
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
2008 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, State University of New York, Cortland
Chair
Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired)
Paul Groves, South Pasadena HS, Pasadena, CA
William Bond, Snohomish HS, Snohomish, WA
David Hostage, Taft School, Watertown, CT
Peter Demmin, Amherst HS, Amherst, NY (retired)
Marian Dewane, Centennial HS, Boise, ID
Valerie Ferguson, Moore HS, Moore, OK
Adele Mouakad, St. John’s School, San Juan, PR
Jane Nagurney, Scranton Preparatory School, Scranton, PA
Ronald Ragsdale, University of Utah, Salt Lake City, UT
Kimberly Gardner, US Air Force Academy, Colorado Springs, CO
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for
a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 23, 2008, after
which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the
front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to
use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this
examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials,
scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
Page 1
!
!
!
!
!
!
Key for 2008 National Olympiad (part 2)
1. (14%) Benzene, C6H6, reacts with Br2 in the presence of FeBr3 as a catalyst to give an organic compound with the percentage
composition by mass; C 30.55%, H 1.71%, Br 67.74% and hydrogen bromide.
a. Determine the empirical formula of the compound.
b. When 0.115 g of this compound are dissolved in 4.36 g of naphthalene the solution freezes at 79.51 ˚C. Pure naphthalene
freezes at 80.29 ˚C and has a kf = 6.94 ˚C·m–1. Determine the molar mass and molecular formula of the compound.
c. Write a balanced equation for the reaction.
d. Calculate the theoretical yield for the organic compound when 4.33 g of of benzene is reacted with an excess of bromine.
e. If the actual yield of the reaction is 5.67 g, what is the percentage yield?
f.
i. Write structures for the possible isomers that could be formed in this reaction.
ii. Identify the major isomer(s) formed in this reaction and explain your reasoning.
a)
convert masses to moles:
# 1 mol &
1.71 g H " %
( = 1.70 mol (÷0.848) = 2.00
$ 1.008 g '
# 1 mol &
30.55 g C " %
( = 2.54 mol (÷0.848) = 3.00
$ 12.011 g '
# 1 mol &
67.74 g Br " %
( = 0.848 mol (÷0.848) = 1.0
$ 79.90 g '
These numbers are whole numbers, so the empirical formula must be C3H2Br
b) ΔT = 80.29 – 79.51 = 0.78 oC. Plugging this value into the formula for freezing point depression gives,
"T = k f • m and m = 0.78 o C 6.94 o C / m = 0.11m
mol
0.115 g
0.11
" 0.00436 kg = 0.00048 mol so, MM =
= 240 g # mol-1
kg
0.00048 mol
240 / 117.9 = 2.03 which is approximately 2, so the molecular formula must be C6H4Br2
c)
C6H6 + 2Br2 r C6H4Br2 + 2 HBr
d) The theoretical yield is:
# 1 mol C H & # 1 mol C H Br & # 235.89 g C H Br &
6 6
6 4
2
6 4
2
4.33 g C 6 H 6 " %
( "%
(" %
( = 13.1 g C 6 H 4 Br2
$ 78.11 g C 6 H 6 ' $ 1 mol C 6 H 6 ' $ 1 mol C 6 H 4 Br2 '
" 5.67 g %
Percent yield is: $
' ( 100% = 43.4%
# 13.07 g &
The possible isomers for i. are,
e)
Br
!
Br
Br
Br
ii. the major products are,
Br
Br
Br
Br
Br
Page 2
Br
because –Br is an ortho-para director. The para isomer
should be the most prominent product because of steric
hindrance for the ortho product.
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
!
!
!
!
!
!
!
2.
(10%) Photochemical smog is formed through a sequence of reactions, the first three of which are given below. Smog is formed
when the O(g) produced in reaction (3) reacts with organic molecules.
Bond Dissociation Energy, kJ⋅mol–1
(1)
N2(g) + O2(g) r 2NO(g)
N–N
193
(2)
2NO(g) + O2(g) r 2NO2(g)
N=N
418
(3)
NO2(g) + hν r NO(g) + O(g)
N≡N
941
a. For reaction (1), ∆Hº = +180.6 kJ·mol–1. Calculate the bond
O–O
142
dissociation energy of NO(g).
O=O
498
b. Calculate the entropy change for the first reaction.
So, J⋅mol–1⋅K–1
c. Determine the minimum temperature at which reaction (1) becomes
N2(g)
191.5
spontaneous.
O2(g)
205.0
d. For reaction (3), ∆Hº = +306 kJ·mol–1. If the energy for this
NO(g)
210.6
reaction were provided by sunlight, estimate the wavelength
NO2(g)
240.5
required and specify the region of the spectrum containing this
O(g)
161.0
wavelength.
a)
"H =
The overall enthalpy change can be estimated from the bond dissociation energies via the equation,
# Energy of bonds broken $ # Energy of bonds formed
180.6 kJ = 941 kJ + 498 kJ " 2 # BDE NO
so, BDE NO = 629 kJ " mol–1
b) Similarly,
(
)
"S o = 2S o (NO) # S o (N 2 ) + S o (O 2 )
"S o = 2(210.6) # ( (191.5) + (205)) = 24.7 J $ mol-1 $ K -1
c)
Utilize the equation, "G o = "H o # T"S o , and set ΔGo to zero to find the minimum temperature.
(
)
0 = 180.6 kJ " mol-1 # T $ 0.0247 kJ " mol-1 " K -1 , so T = 180.6 kJ " mol-1 0.0247 kJ " mol-1 " K -1 = 7311 K
d) First convert to energy per molecule,
!
#
&
J
1 mol
J
-19
3.06 " 10 3
"%
. Now calculate wavelength of light with this energy,
( = 5.08 " 10
23
mol $ 6.022 " 10 molecules '
molecule
-34
8
-1
%
h # c ' 6.626 $ 10 J #s 3.0 $ 10 m #s
" =
=
'
E
5.08 $ 10 -19 J # molecule
&
(
)(
) (* = 3.91$ 10
-7
*
)
m (per molecule) = 391 nm (in the ultraviolet).
(12%) Aniline, C6H5NH2, reacts with water according to the equation: C6H5NH2(aq) + H2O(l) s C6H5NH3+(aq) + OH –(aq)
3.
In a 0.180 M aqueous aniline solution the [OH–] = 8.80×10–6.
!
a.
b.
c.
d.
Write the equilibrium constant expression for this reaction.
Determine the value of the base ionization constant, Kb, for C6H5NH2(aq).
Calculate the percent ionization of C6H5NH2 in this solution.
Determine the value of the equilibrium constant for the neutralization reaction;
C6H5NH2(aq) + H3O+(aq) s C6H5NH3+(aq) + H2O(l)
e.
i. Find the [C6H5NH3+(aq)] / [C6H5NH2(aq)] required to produce a pH of 7.75.
ii. Calculate the volume of 0.050M HCl that must be added to 250.0 mL of 0.180 M C6H5NH2(aq) to achieve this ratio.
a) K b =
[C 6H 5NH +3 ][OH -]
[C 6H 5NH 2 ]
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
!
Page 3
Key for 2008 National Olympiad (part 2)
(8.80 " 10 )(8.80 " 10 ) = 4.3" 10
#6
b) K b =
#6
#10
( 0.180)
(8.80 " 10 ) " 100% = 4.9 " 10
#6
c) % ionization =
!
#3
( 0.180)
%
d) C 6 H 5 NH 2 + H 3O + " C 6 H 5 NH +3 + H 2O so K =
!
K b 4.3# 10$10
=
= 4.3# 10 4
K w 1.0 # 10$14
e) (i) For a pH = 7.75, the pOH = 6.25 so [OH - ] = 10"pOH = 5.62 # 10"7 M.
!
4.3 " 10#10 =
[C 6H 5NH +3 ][OH -]!so, [C 6H 5NH +3 ] = 4.3 " 10#10 = 7.65 " 10#4
[C 6H 5NH 2 ]
[C 6H 5NH 2 ] 5.62 " 10#7
(ii) The HCl is a strong acid that will protonate the aniline, so to get the HCl required, we need the amount of C6H5NH2 required
multiplied by the value of the ratio from (i): 7.65" 10#4 " 0.250 L " 0.180M C 6 H 5 NH 2 = 3.44 " 10#5 mol HCl
Now determine the volume !
of reagent: 3.44 " 10#5 mol HCl " 1 L 0.050 mol HCl = 6.88 " 10#4 L = 0.688 mL
!
4.
!
(10%) Gaseous dinitrogen pentoxide,
N2O5, decomposes to form nitrogen dioxide and oxygen gas with the initial rate data at
25 ˚C given in the table.!
[N2O5], M
Rate, mol.L–1.min–1
a.
b.
c.
d.
0.150
3.42×10–4
0.350
7.98×10–4
0.650
1.48×10–3
Write a balanced equation for this reaction.
Use the data provided to write the rate law and calculate the value of k for this reaction. Show all calculations.
Calculate the time required for the concentration of a 0.150 M sample of N2O5 to decrease to 0.050 M.
The initial rate for the reaction of a 0.150 M sample is 2.37×10–3 mol.L–1.min–1 at 40 ˚C. Determine the activation energy for
this reaction.
a) 2N 2O 5 " 4NO 2 + O 2
!
!
!
!
!
b) The rate will be given by the rate law (Rate=k[N2O5]x) in each case, so by taking the ratio, the rate constant cancels,
x
Rate1 [ N 2O 5 ]1
=
Rate 2 [ N 2O 5 ] x
2
x
x
7.98 " 10#4 $ 0.350 '
1.48 " 10#3 $ 0.650 '
x
means
that
2.33
=
2.33
so
x=1.
Checking
with
a
second
set
of
data,
=
=
&
)
&
) leads to
3.42 " 10#4 % 0.150 (
3.42 " 10#4 % 0.150 (
4.33x = 4.33, confirming that the reaction is first order.
Now calculate the rate constant: 1.48 " 10#3 = k(0.650)1 so k = 2.28 " 10 -3 min#1 and we have
Rate = 2.28×10-3 min -1[N2O5].
!
"
%
! law, ln$ [ N 2O 5 ] init ' = kt . Plugging in ln"$ 0.150 %' = ln(3) = (2.28 ( 10)3 min)1)t so t = 481 minutes.
c) Use the integrated rate
$ [N O ] '
# 0.050 &
2 5 t &
#
d) Use the information from the two temperatures given in the Arrhenus equation:
" 1.58 ) 10 -2 %
"k % E " 1
" 1
1 %
Ea
1 %
ln$ 2 ' = a $ ( ! ', so, plugging in values gives : !ln$
)
(
'=
'
-3
-1
-1 $
# 298 K 313 K &
# k1 & R # T1 T2 &
# 2.28 ) 10 & 8.314 J * mol * K
Ea
so ln( 6.93) =
# ( 0.00335 $ 0.00319) and solving for E a gives : E a = 1.00 # 10 5 J " mol-1 = 100 kJ " mol-1 . (Note
8.314 J " mol-1 " K -1
that using rates, rather than rate constants in the argument of the natural log is an alternative, correct method.)
Page 2
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Barium peroxide is added to water.
b. Acidic solutions of potassium iodide and potassium iodate are mixed.
c. A phosphoric acid solution is added to a solution of calcium hydrogencarbonate.
d. Solutions of lead(II) nitrate and potassium chromate are mixed.
e. Concentrated hydrochloric acid is added to an aqueous solution of cobalt(II) nitrate.
f. 2-butanol is heated with concentrated sulfuric acid.
a) BaO 2 + H 2O " Ba 2+ + HO -2 + OH b) I - + IO -3 + H + " I 2 + H 2O
c) H 3 PO 4 + Ca 2+ + HCO"3 # Ca 3 (PO 4 ) 2 + H 2O + CO 2
!
!
!
!
!
d) Pb 2+ + CrO 24 " PbCrO 4
e) Co 2+ + Cl- " CoCl24
f)
H3C CH2
OH
C CH3 + H2SO4
H
H3C CH
C CH3
H
+
+ H2O
H3C CH2 C CH2
H
either isomer counts
6. (12%) The apparatus depicted to the right is often used to demonstrate the electrolysis of
water. Tubes A and B are initially filled with an aqueous solution of H2SO4 or Na2SO4.
a. Describe the purpose of adding the H2SO4 or Na2SO4 rather than using pure water.
b. Give the formula of the gas produced in;
i. tube A
ii. tube B
c. Describe a chemical test that could be used to identify the gas collected in tube A.
Include the procedure and expected observation.
d. Calculate the number of moles of gas expected to be collected in tube B when a
600. milliamp current is applied for 40.0 minutes. (Assume no side reactions occur.)
e. Calculate the volume of the gas produced in part d. for a temperature is 20 ˚C and a
pressure in the laboratory of 735 mmHg. (The vapor pressure of water is 17.5 mmHg.)
f. If H2O2 is formed in a side reaction the quantity of only one of the products is affected.
Identify the product affected and state how its quantity compares with that produced with
no side reaction. Explain your answer.
a)
Because pure water is a poor conductor of electricity, the H2 SO4 or Na2SO4 is added to provide electrolyte (so that
the solution will conduct).
b) i) tube A is the cathode, therefore it is the site of reduction where H2 is produced, while (ii) tube B is the anode,
where oxidation occurs, therefore O2 is produced.
c) Because H2 is flammable, a burning splint can be inserted into the products from Tube A. If there is a “pop”
associated with the reaction, it confirms that the gas is H2.
d) Charge = current × time: 0.600 C·s-1 × 2400 s = 1440 C
# 1 mol e - & # 1 mol O &
2
= 3.73 " 10)3 mol O 2
and: 1440 C " %
(" %
- (
$
'
96500
C
4 mol e
$
'
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
!
Page 5
Key for 2008 National Olympiad (part 2)
e)
First correct for vapor pressure of water: Ptotal = PO 2 + PH 2 O so PO 2 = 735- 17.5 = 717.5 mmHg
nRT (3.73 " 10#3 mol)(0.821 L $ atm $ mol-1 $ K -1)(293 K)
=
= 0.095 L = 95 mL
1 atm
P
717.5 mmHg "
! 760 mm Hg
f) The quantity of O2 would be affected, but the quantity of H2 would not. The yield of O2 would be decreased because
some of the electricity would oxidize H2O into peroxide (H2O2) instead of O2.
V=
!
7.
(16%) Explain the following observations in terms of bonding principles.
a. Carbon dioxide is a gas at room temperature and pressure but silicon dioxide is a high-melting solid.
b. The xenon trioxide molecule has a trigonal pyramidal shape while sulfur trioxide is trigonal planar.
c. In many of its ionic compounds oxygen is present as the O2– ion although the addition of two electrons to an oxygen atom in
the gas phase is an endothermic process.
d. (bmim)+PF6– is a liquid at room temperature while (bmim)+Cl– and Na+PF6– are solids.
Note: (bmim)+ is an abbreviation for N-butyl-N-methylimadazolium ion, CH3N2C3H3C4H9+.
a) Carbon dioxide is small, non-polar molecule. The intermolecular forces between them are small, so CO2 is a gas. Conversely,
silicon dioxide is a network solid. As a network, the connections are covalent bonds, which are quite strong compared to the
intermolecular forces between small molecules, and it requires a great deal of energy to break an SiO2 unit away from the rest
of the solid. Ultimately the key bonding feature in these molecules that gives rise to this difference is that carbon atoms readily
form double bonds, where double bonds to silicon are much less common.
b) Looking at the Lewis structure of the two compounds provides the answer:
O
Xe
O
O
O
S
O
O
There are four charge centers (three bonding, and one lone pair) around Xe in XeO3 leading to a trigonal pyramidal shape, while there
are three charge centers (all of them bonding pairs) in SO3 which is trigonal planar.
c) When the oxide ion, O2–, is in ionic compounds, the 2– charge is interacting with positively charged cations. Thus, even though
the ion formation in the gas phase is endothermic, the oxide ion exists in ionic compounds.
d) Because the (bmim)+ and PF6–ions are quite large, the lattice energy between the two items will be small. The energy available as
heat at room temperature is sufficient to overcome this interaction energy. By contrast (bmim)+ and Cl–and Na+ and PF6– have
one large and one small ion, so they can pack more closely and have larger lattice energy. These larger energies mean the
compounds are solids at room temperature.
8.
(14%) This question deals with the bonding in several organic chemicals.
a. Several different compounds have the formula C2H4O2. Two of these contain –CO2 groups.
i. Give the structures and names of the two compounds with –CO2 groups.
ii. These compounds boil at 31.5 ˚C and 118 ˚C. Assign the two boiling points to the structures in i. and account for the
boiling point difference in terms of their structures.
iii. Sketch the structure of one of the other compounds.
b. Fatty acids are important components of a healthy diet. Three fatty acids are stearic, oleic and linoleic which have the
formulas CH3C16H32COOH, CH3C16H30COOH, and CH3C16H28COOH, respectively.
i. Describe the differences in bonding suggested by the formulas of these compounds.
ii. The compounds melt at –5 ˚C, 13 ˚C and 69 ˚C. Assign these melting points to the respective acids and account for this
behavior in terms of their structures and bonding.
iii. The salts of fatty acids can be used as soaps or detergents. Describe the chemical basis of this behavior.
a)
Page 2
(i) The two structures are:
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
H
H
O
C
C
O
H
OH
C
H
O
C
H
H
H
ethanoic acid
methylmethanoate
ii) 118 o C is ethanoic acid and 31.5 o C is methyl methanoate. The key difference arises from the strength of intermolecular
forces present in ethanoic acid, which can participate in hydrogen bonding, while the strongest intermolecular forces present in
methylmethanoate are dipole-dipole forces.
iii) Possible correct structures include:
H
OH
C
H
C
H
C
OH
HO
C
HO
H
C
H
OH
O
H
C
C
C
H
H
O
H
OH
H
b) (i) CH3C16H32COOH contains a saturated alkyl chain. CH3 C16H30COOH contains a one carbon-carbon double bond.
CH3C16H28COOH contains two carbon-carbon double bonds.
(ii) CH3C16H32COOH melts at 69 oC. It is the highest melting point because the saturated alkyl chain tails are capable of being
closely packed, thereby maximizing the dispersion forces present. Higher intermolecular forces lead to higher melting points.
CH3C16H30COOH with one double bond has additional geometrical constraints due to the relative rigidity of that double bond, so
the tails cannot pack as efficiently, and the melting point is lower, at 13 oC. Finally, for CH3C16H28COOH with two double
bonds, the geometric constraints just noted are even more sizable, so packing is even less efficient. It will, therefore, have the
lowest melting point, –5 oC.
(iii) The key feature is that the molecules have a charged region (often called the head) where the acid group is and an
uncharged and not very polar region (called the tail) where the alkyl chains are. The non-polar tail can interact relatively
strongly with non-polar dirts, oils and greases, leaving the polar/charged head group “sticking out”. This polar/charged group
interacts strongly with polar water molecules. Thus, while polar water molecules do not wash away non-polar dirts and oils by
themselves, taking advantage of the dual behavior of the long-chain fatty acids, the dirt/oil is encapsulated in a micelle-like
structure that can be solvated by water.
Pictorially, showing far too few fatty acids, it would look something like this.
CO2-
CO2-
CO2CO2CO2CO2-
CO2CO2-
END OF PART II
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
Page 7
KEY for 2008 National Olympiad (Part 2)
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
NA gas constant
h
R Planck’s constant
°C gram
P
g pressure
c heat capacity
k
Cp rate constant
C hour
Rf
h retention factor
E joule
s
J second
Ea kelvin
c
K speed of light
H kilo- prefix
T
k temperature, K
S liter
t
L time
volt
V
CONSTANTS
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
E = E! –
! k2 $ Ea ! 1 1 $
=
'
" k1 &% R #" T1 T2 &%
" –!H % " 1 %
ln K = $
' $ ' +c
# R & # T&
RT
ln Q
nF
ln #
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Uun
111
Uuu
112
Uub
114
Uuq
116
Uuh
118
Uuo
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
(2??)
(2??)
Page 8
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO Olympiad National Exam after April 23, 2008.
2008 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Brookline High School, Brookline, MA
Chair
Linda Weber, Natick High School, Natick, MA
John Mauch, Braintree High School, Braintree, MA
Nancy Devino, ScienceMedia Inc., San Diego, CA
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 23, 2008
Page 1
2007 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
You have been given seven pipets that contain solutions of AgNO3, BaCl2, Cu(NO3)2, CuSO4, Pb(NO3)2, KI,
and Na2S2O3, though not necessarily in this order. Using the materials provided, devise and carry out an
experiment to correctly determine the contents of each pipet.
Lab Problem 2
Given a sample of an unknown metal carbonate, MxCO3, and 3.0M hydrochloric acid, HCl(aq), a balloon, and
some laboratory equipment, devise and carry out an experiment by combining these two substances to
determine the volume of the gas produced and the unknown metal. The possible metals are Ba, Ca, Li, or Na.
Room Temp. = 25oC, Standard Pressure = 1 atm
Answer Sheet for Laboratory Practical Problem 1
Page 2
Not valid for use as an USNCO National Examination after April 23, 2008.
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 23, 2008
Examiner’s Initials:
Page 3
2. Record your data and other observations.
3. Based on your observations, write the relevant equations that led to your conclusions.
4. Conclusions
Pipet
Contents
Justification
#1
#2
#3
#4
#5
#6
#7
Page 4
Not valid for use as an USNCO National Examination after April 23, 2008.
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name :________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 23, 2008
Examiner’s Initials:
Page 5
2. Record your data and other observations.
3. Calculations and Conclusions.
4. Conclusions: The volume of gas produced:
The unknown metal:
5. Sources of Error in this experiment (please number):
Page 6
Not valid for use as an USNCO National Examination after April 23, 2008.
PERIODIC TABLE OF THE ELEMENTS
2
He
1
H
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Uub
114
Uuq
116
Uuh
118
Uuo
(223)
226.0
227.0
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
(2??)
(2??)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Not valid for use as an USNCO National Examination after April 23, 2008
Page 7
2008 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Instructions
Directions to the Examiner:
Thank you for administering the 2008 USNCO laboratory practical on behalf of your Local Section. It is
essential that you follow the instructions provided, in order to insure consistency of results nationwide. There
may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important
that you do not lend any assistance or hints whatsoever to the students once they begin work. As in the
international competition, the students are not allowed to speak to anyone until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab station or
table when the students enter the room. The equipment should be initially placed so that the materials used for
Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad
Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a
laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are
instructed to do so.
You will be asked to complete two laboratory problems. Students are to work alone. All the materials and
equipment you may want to use to solve each problem has been set out for you and is grouped by the number of
the problem. Students can use all materials for both lab problems, but each experiment is designed to work best
with equipment and materials provided specifically for each lab problem. You will have one hour and thirty
minutes to complete the two problems. You may choose to start with either problem. You are required to have a
procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that
your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your
station for each safety approval, raise your hand.
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any
chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for
disposing of solutions at the end of this lab practical are:
__________________________________________________________________________________________
____________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. There is a periodic table and constants on the last page.
Are there any questions before we begin?
Page 8
Not valid for use as an USNCO National Examination after April 23, 2008.
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table
on the last page of the booklet. Allow students enough time to read the brief cover directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out
all information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your
cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab
practical with the students. They can learn about possible observations and interpretations and you can acquire
feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a
few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the
Olympiad subcommittee as they prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets
from Part I, and the “Blue Books” from Part II in the UPS return envelope you were provided to this address:
ACS DivCHED Exams Institute
Department of Chemistry
University of Wisconsin – Milwaukee
3210 N Cramer Street
Milwaukee, WI 53211
The label on the envelope should have this address already, you will need only to include your return address
and call United Parcel Service - UPS (1-800-742-5877) for it to be picked up (or it can be dropped in a UPS
collection box). The cost of shipping will be billed to the Exams Institute. You can write down the tracking
number on the label to allow you to track your shipment.
Wednesday, April 23, 2008, is the absolute deadline for receipt of the exam materials at the Examinations
Institute. Materials received after this deadline CANNOT be graded. Be sure to have your envelope picked up
no later than April 21, 2008 for it to arrive on time.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR
GRADING THIS EXAMINATION.
Not valid for use as an USNCO National Examination after April 23, 2008
Page 9
Examiner’s List: 2008 USNCO Lab Practical Equipment and Chemicals
Lab Problem #1: Materials and Equipment
Each student should have available the following equipment and materials:
• Clear acetate sheet
• One grease pencil, used to write in the acetate sheet
• Seven Microtip or thin-stem Beral-style pipets (approx. 2.5-mL volume) to contain
unknown solutions
• One 150-mL or 250-mL beaker to hold the filled pipets
• 3-4 toothpicks for stirring
• Access to distilled water
Lab Problem #1: Chemicals
Each student will need:
• Solution of AgNO3 (mw=170), BaCl2 (mw=208), Cu(NO3)2 (mw=188), CuSO4
(mw=160), Pb (NO3)2 (mw=331), KI (mw=166), and Na2S2O3 (mw=158).
Notes to Coordinators:
DO NOT IDENTIFY FOR THE STUDENTS WHICH SOLUTION IS IN EACH
PIPET!
• All of the solutions should be 0.10M. The solutions should be filled to the maximum
in each pipet and stored upside-down (bulbs down) in the 150- or 250-mL beakers on
the day of the exam. Students DO NOT have access to additional quantities of any of
these solutions.
• The pipets should be labeled using an indelible marker, i.e., a Sharpie®, writing the
letter in a clear, legible capital letter on the bulb portion of each pipet.
• IMPORTANT: The key for this experiment is as follows:
Na2S2O3 = 1
CuSO4 = 2
KI = 3
AgNO3 = 4
Cu(NO3)2 = 5
Pb(NO3)2 = 6
BaCl2 = 7
Page 10
Not valid for use as an USNCO National Examination after April 23, 2008.
Lab Problem #2: Materials and Equipment
Each student will need:
•
•
•
•
•
•
•
•
•
•
One 50-mL beaker (to contain the 3.0 M HCl) labeled ‘3.0M HCl’
One 10-mL graduated cylinder
One balloon (Included)
One scissors
One metric ruler with mm precision
One length of string approximately 30 cm (12”) in length
2-3 sheets of waxed weighing paper
One metal scoopula
Access to a 0.01 or better electronic balance
Access to distilled water
Lab Problem #2: Chemicals
• Approximately 1.5 g sample CaCO3 in an unlabeled, capped 10- or 20-mL capacity vial
• 20 – 25-mL of 3.0 M HCl
Quick Check to be sure this experiment works for your examinees:
Notes to Coordinators:
• You can pour the HCl into the 50- mL beaker the day of the exam.
• The CaCO3 should be powdered, not granular or in rock form. Be sure that the CaCO3
vial is not labeled and is capped. For your reference, below is CaCO3 “Fisher Science
Education Catalog” product number:
• S719221 Calcium Carbonate: Reagent Grade - Powder; 100g Calcium Carbonate,
White, Application: CO2 generation, CaCO3 ($7.10)
• S71922 Calcium Carbonate: Reagent Grade - Powder; 500g Calcium Carbonate,
White, Application: CO2 generation, CaCO3 ($8.70)
• Please give each balloon a few stretches prior to placing at the student lab bench to
ensure that the balloon will adequately inflate.
Not valid for use as an USNCO National Examination after April 23, 2008
Page 11
USNCO 2008 Part III Answers
Lab Problem #1
This lab problem involves knowledge of precipitation and solubility. The focus of this problem in
qualitatively determining each of the seven unknown solutions is to apply knowledge and understanding
of predicted reactions from observations made combining the solutions with one another using the
provided acetate sheet as a spot plate.
Procedure
Students should have constructed some kind of data table that examines combinations of solutions with
one another. The various color changes and formed precipitates providing evidence which students must
use to form conclusions about each unknown solution.
Qualitative Evidence:
The two copper solutions are blue.
To distinguish between both blue solutions of copper sulfate and copper nitrate, the copper (II) sulfate
will precipitate with barium but not with copper (II) nitrate; both copper (II) solutions will precipitate
with the KI.
Copper (II) ions will form a redish-brown precipitate with iodide.
Silver nitrate forms white solid silver chloride and darkish silver sulfate.
Barium forms insoluble whitish barium sulfate and barium iodide.
Nitrates are soluble and will thus not form precipitates in combination with any of the other cations.
Lead reacts with iodide to form yellow lead iodide, with chloride to form white lead chloride, and with
sulfate to form whitish lead sulfate.
Thiosulfate will react with silver to form a darkish precipitate that then dissolves on further addition of
thiosulfate.
Answers: A = Na2S2O3, B = CuSO4, C = KI, D = AgNO3, E = Cu(NO3)2, F = Pb(NO3)2, G =BaCl2.
Lab Problem #2
This lab problem involves identification of an unknown metal carbonate (here, CaCO3) by quantitatively
reacting a measured mass of the solid sample to 3.0M HCl, capturing the CO2 gas evolved, and
determining the volume of the gas collected in the balloon provided to then calculate the molar mass of
the metal carbonate and conclude which metal carbonate was present from the choices given. One
experiment might have been to weigh the gas volume produced or measure the gas volume by water
displacement (though materials provided were insufficient to approximate the volume produced by
water displacement). Another avenue to determine the unknown metal, though not implied with the
materials and chemicals provided for this lab problem, could have been to react the solid white calcium
carbonate with several of the identified solutions from lab problem #1, but there is no quantitative
evidence here, and the lab problem specifies not only to identify the unknown metal, M, but to also
determine the volume of the gas produced.
Students needed to figure out a method of gas collection without the volume of HCl(aq) occupying the
balloon. One possibility was to insert a measured mass of the solid carbonate into the balloon, then
carefully pull the lip of the balloon over the top of the provided graduated cylinder (filled with the HCl.
By shaking the solid carbonate into the cylinder with the balloon still over the top of the cylinder, the
Page 12
Not valid for use as an USNCO National Examination after April 23, 2008.
two chemical reactants can combine to produce CO2 gas to fill the balloon. One method for measuring
the volume could have been to wait for the completed reaction, then carefully twist shut the balloon at
its lip, tie off the balloon to make a more spherical shape, then use the string and ruler to measure the
balloon’s circumference to then calculate its approximate spherical volume using the formula for
circumference and volume of a sphere. Students might have also measured the circumference, emptied
the balloon of the products, then refilled it with water to the approximate volume when filled with CO2,
and finally measure the volume of water using the graduated cylinder. The distinguish between the
possibility of Na2CO3 (molar mass = 106) vs. CaCO3 (molar mass = 100), students might have taken a
small solid sample remaining and test its relative solubility in water. Though not explicitly encouraged,
students might also have tested the unknown carbonate with the known Cu (II) solutions from
Experiment #1 to observe precipitates form.
Points of error abound, including assuming standard pressure (1 atm), room temperature (25 oC), a
perfect sphere, CO2 collected as a ‘wet’ gas, completed reaction, and some air and uncaptured CO2 gas
remaining in the graduated cylinder.
Sample Student Data:
General reaction: MxCO3 + HCl Æ MCly + H2O(l) + CO2(g), where the mol ratio of MxCO3 : CO2
is 1 : 1.
Mass of MxCO3 used
= 1.00 g
Circumference of balloon
= 23.9 cm
Radius of balloon, using C = 2πr
= 3.8 cm
Calculation of volume from a sphere, V = 4/3πr3
= 240 cm3
Use ideal gas law, PV = nRT, to find moles of CO2 = (1)(0.240) = n(.0821)(298), n = 0.00981
Molar mass of MxCO3: 0.00981 mol = 1.00g/mol. mass
M = Ca, therefore unknown compound is CaCO3 .
= 102 or close to 100, mol. mass if
Not valid for use as an USNCO National Examination after April 23, 2008
Page 13
USNCO 2008 Part III Grading Results
Lab Problem #1
Excellent Students Results:
Student provided clear explanation of procedure that included an organized and methodical plan to react
solutions with each other and obtain qualitative data from which to make conclusions.
Student showed an organized data table with all possible trial combinations and complete descriptions of
each reaction.
Student showed at least four relevant equations that follow a logical conclusion based on student
observations. Written equations were all balanced and indicated both precipitates and aqueous ions.
Student clearly based their conclusions on chemical knowledge about solubility and color of possible
precipitates. Student used deduction and analysis to infer the identity of the unknowns using the
qualitative data they obtained in this experiment.
Average Student Results:
Student conducted most of the trial combinations and made use of chemical knowledge to form
conclusions about the unknowns, with some incorrect and missing information.
Not all solutions correctly identified.
Written equations correctly identified precipitates but were either not completely balanced or did not
correctly show aqueous ions. Only several of the relevant equations were included or the minimum were
included but not al balanced correctly.
Some qualitative information was given to demonstrate an understanding of precipitate formation and
solubility
Below Average Student Results:
Student was unable to apply adequate chemical knowledge from solution combinations.
Written equations were incomplete.
Little or no qualitative information was used to make predictions about the identity of the unknown
solutions.
Lab Problem #2
Excellent Students Results:
Student proposed a clear, detailed procedure for determining the unknown metal by gas collection and
measured gas volume.
Data was organized and calculations were clear and followed a logical plan based on the student’s
experiment.
A detailed and complete listing of points of error was given following the experiment.
Multiple trials were performed.
Credibly creative alternative methods to determine the unknown carbonate were used, including testing
solubility and precipitation with known solutions from Exp. #1
Average Student Results:
Student made adequate measurements and had a general understanding of connection between gas
collection, volume measurement, and identifying the unknown metal by determining the carbonate’s
molar mass.
Student incorrectly identified the unknown metal or miscalculated the volume of gas due to one or more
sources of error.
Page 14
Not valid for use as an USNCO National Examination after April 23, 2008.
Only 1-2 points of error were given following the experiment.
Below Average Student Results:
Student was not able to connect the volume of gas collected to the calculation of the unknown
carbonate’s molar mass.
An experiment was performed to collect the gas but both reactants were placed in the balloon together.
One or no points of error were given.
Not valid for use as an USNCO National Examination after April 23, 2008
Page 15
2009 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, Chair, State University of New York, Cortland, NY
James Ayers, Mesa State College, Grand Junction, CO
Paul Groves, South Pasadena HS, South Pasadena, CA
Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired) Preston Hays, Glenbrook South HS, Glenbrook, IL
Seth Brown, University of Notre Dame, Notre Dame, IN
David Hostage, Taft School, Watertown, CT
Peter Demmin, Amherst HS, Amherst, NY (retired)
Marian Dewane, Centennial HS, Boise, ID
Valerie Ferguson, Moore HS, Moore, OK
Adele Mouakad, St. John’s School, San Juan, PR
Jane Nagurney, Scranton Preparatory School, Scranton, PA
Ronald Ragsdale, University of Utah, Salt Lake City, UT
Kimberly Gardner, US Air Force Academy, Colorado Springs, CO
DIRECTIONS TO THE EXAMINER–PART II
Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for
a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the
examination period. All testing materials including scratch paper should be turned in and kept secure until April 29, 2009, after
which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”,
Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the
front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest-breaks between parts.
Part I
Part II
Part III
60 questions
8 questions
2 lab problems
single-answer multiple-choice
problem-solving, explanations
laboratory practical
1 hour, 30 minutes
1 hour, 45 minutes
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to
use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same
identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides
of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this
examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials,
scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
Page 1
1. (12%) Butanoic (butyric) acid, C3H7COOH, is a monoprotic acid with Ka = 1.51×10–5. A 35.00 mL sample of 0.500 M
butanoic acid is titrated with 0.200 M KOH.
a. Calculate the [H+] in the original butanoic acid solution.
b. Calculate the pH after 10.00 mL of KOH have been added.
c. Determine the pH at the half-equivalence point of the titration.
d. Find the volume of KOH solution needed to reach the equivalence point for the titration.
e. Calculate the pH at the equivalence point.
(a.) Let HA = C3H7COOH, and A–= C3H7COO–
Ka
H + ][ A - ]
[
=
[ HA]
Let [H+] and [A-] = x. Plugging in we get, 1.51" 10#5 =
x2
. Solving for x gives, [H+] = 2.73×10–3
(0.500 # x)
(b.) Determine the initial number of moles of acid: 0.03500 L " 0.500 M = 0.0175 mol HA
Determine the number of moles of NaOH added:
! 0.0100 L " 0.200 M = 0.0020 mol OH added
!
0.0155 mol HA
Determine the molarity of H+: 0.0175!– 0.0020 = 0.0155 mol HA remain, in 0.045 L, so the molarity is
= 0.344 M
0.045 L
!
0.0020 mol A The [A-] is changed only through dilution, [A-] =
= 0.0444 M
0.045 L
!
Plug these values into the equilibrium constant expression and solve for [H+].
[H ](0.0444) ,and
+
1.51"10#5 =
(0.344)
! "10
[H ] = 1.16
+
-4
, so pH = log(1.16 "10 -4 ) = 3.93
(c.) At the half equivalence point: [HA] = [A-] and [H+] = Ka = 1.51×10–5. So, pH = log(1.51×10–5) = 4.82
!
(d.) The initial moles HA (from part b.) = 0.0175 mol HA, so the equivalence point is reached when we have 0.0175 mol OH– added.
0.0175mol OH - "
1L
= 0.0875 L .
0.200 mol OH -
(e.) Total volume at the equivalence point is 0.0875 L + 0.0350 L = 0.1225 L
[ ]
At the equivalence point all HA is converted to A-, so: A" =
!
0.0175 mol
= 0.143 M
0.1225 L
The A- is a base according to the equation, A– + H2O ⇔ HA + OH–, and K b =
!
Let [OH–] and [HA] = x. Plugging in we get, 6.63 "10#10 =
K w 1.0 " 10#14
=
= 6.63" 10#10
K a 1.51" 10#5
x 2 . Solving for x = [OH–] = 9.73×10-6.
(0.143)
!
-6
So, pOH = -log(9.73×10 ) = 5.01 and pH = 14.00 – 5.01 = 8.99
2.
!
(12%) Chromium metal reacts with acid to produce Cr3+ ions and hydrogen gas.
a. Write a balanced equation for this reaction.
b. When a sample of chromium metal is reacted with excess acid, 94.7 mL of gas is collected over water at 745 mm Hg and
20˚C. Assuming ideal gas behavior determine the mass of metal reacted. (The vapor pressure of water at 20˚C is 24 mmHg.)
c. State and explain how the volume of gas would change if Cr2+ (rather than Cr3+) ions were formed in this reaction.
d. Determine the number of molecules of water vapor that would be present in the volume of gas produced assuming ideal
behavior.
e. Calculate the ratio of the average molecular velocity of hydrogen to the average molecular velocity of water vapor at the same
temperature.
f.
Page 2
(
"
%2 +
*)
# V & -,
n
The van der Waals equation for real gases is *P + a$ ' - . [V / nb] = nRT .
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
!
The coefficients, a and b, for the hydrogen gas are a = 0.242 atm.L2.mol–2 and b = 0.0266 L.mol–1. The corresponding values
of a and b for sulfur dioxide are 6.714 and 0.05636, respectively.
i. Identify the molecular property that corresponds to the a coefficient and account qualitatively for the difference between
its values for these two gases.
ii. Identify the molecular property that corresponds to the b coefficient and account qualitatively for the difference between
its values for these two gases.
a. 2Cr + 6H+  2Cr3+ + 3H2
b. The pressure of H2(g) produced is: 745 mmHg – 24 mmHg = 721 mmHg
" 721
%
atm'( 0.0947 L)
$
pV
# 760
&
Determine number of moles of hydrogen: n =
=
= 0.00373 mol H 2
RT
0.0821 L ( atm mol ( K ( 293 K)
2 mol Cr 52.00 g Cr
"
= 0.129 g Cr
Now get mass of chromium: 0.00373 mol H 2 "
3 mol H 2 1 mol Cr
(
)
c. The volume of gas would be only!2/3 as much (63.1 mL) becase each Cr atom would release only two electrons to reduce the
H+ ions (rather than 3.)
" 24!
%
atm'( 0.0947 L)
$
pV
# 760
&
d. n =
=
= 1.24 ) 10 -4 mol H 2
L
(
atm
RT
0.0821
mol ( K ( 293 K)
(
)
and 1.24 " 10 -4 mol H 2 "
!
e.
!
vH2
vH2O
=
MM H 2 O
MM H 2
=
6.02 " 10 23 molecules H 2
= 7.49 " 1019 molecules H 2
1 mol H 2
18
= 3 so the ratio of velocities is 3:1
2
!
f. (i.) The “a” coefficient is part of the term that is a correction factor for the attractive forces between molecules. SO2 has a
larger value for “a” because SO2 molecules have stronger forces (due to it being both larger than H2, and polar.)
(ii.) The “b” coefficient is part of the term that is a correction factor for the molecular volume. Because SO2 is a larger
molecule, (more volume) the value for “b” is larger.
3. (14%) The reaction of bromate and bromide ions in acid solution is
represented by the equation,
5Br–(aq) + BrO3–(aq) + 6H3O+(aq) r 3Br2(aq) + 9H2O(l)
In order to measure the rate of the reaction, stock solutions
were prepared as shown in the table:
Stock Solution Concentrations
Stock Br– solution
1.37 M
–3
Stock BrO3– solution
7.10×10 M
Stock H3O+ solution
0.573 M
Reaction mixtures were prepared by mixing the volumes of solutions listed below, and the initial rate of disappearance of
bromate ion was measured.
Expt.
1
2
3
4
a.
b.
c.
Vol. Br– stock
(mL)
0.100
0.200
0.100
0.200
Vol. BrO3- stock
(mL)
0.500
0.500
1.000
0.500
Vol H3O+ stock
(mL)
1.000
1.000
1.000
0.700
Vol H2O
(mL)
1.400
1.300
0.900
1.600
Initial rate of BrO3–
disappearance (mol·L–1·s–1)
5.63×10–6
1.09×10–5
1.13×10–5
5.50×10–6
Calculate the rate of appearance of Br2(aq) in experiment one.
Write the rate law for this reaction and give the value of the specific rate constant, k.
The following mechanism is proposed for the reaction:
(I)
BrO3–(aq) + H3O+(aq) r HBrO3(aq) + H2O(l)
(II)
HBrO3(aq) + H3O+(aq) r H2 BrO3+(aq) + H2O(l)
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
Page 3
(III)
H2BrO3+(aq) r BrO2+(aq) + H2O(l)
(IV)
BrO2+(aq) + Br–(aq) r BrOBrO(aq)
(V)
BrOBrO(aq) + Br–(aq) r Br2 (aq) + BrO2–(aq)
Subsequent reactions of BrO2-(aq) are fast.
i. Draw a Lewis structure of BrO2+ and predict its geometry.
ii. Given the rate law you determined in b, which of the steps (I)-(V) could potentially be rate-limiting?
Justify your answer.
a.
"[ Br2 ]
"[ Br2 ]
= 3#
= 3(5.63 # 10$6 ) = 1.69 # 10$5 mol
L %s
"t
"t
b.
• Looking at experiments 2 and 1, the volume of Br– is doubled and the rate increases by 1.94 (essentially doubles) so the
reaction is first order in Br–.
• Looking at experiments 3 and 1, the volume of BrO3– is doubled and the rate doubles so the reaction is first order in BrO3–.
• Looking at experiments 2 and 4, the volume ratio of H3O+ is (1.00/0.700 = 1.4) and the rate ratio is (1.09 / 0.55 = 1.98) so
the reaction is second order in H3O+ because (1.4)2 = 1.96
Thus, the rate law is : rate = k[Br–][ BrO3–][ H3O+]2 and
5.63 " 10#6 mol L $ s
3
rate
k=
=
= 2.86 L
2
2
mol3 $ s
& 0.1
)& 0.5
)
#3 mol )& 1
Br - BrO 3- H 3O +
mol
mol
%
1.37
%
7.1
"
10
%
0.573
(
L +*(' 3
L +*(' 3
L +*
' 3
!
[ ][
][
]
c. (i.) One resonance structure is shown below, and the shape determined by VSEPR is bent.
O
!
Br
O
(ii.) To have the rate law determined in part (b), the rate limiting step of the mechanism must depend on the concentrations of
[Br–], [BrO3–], and [H3O+]2. The only way for this to be true is for Step IV to be the rate limiting step. One way to confirm
this is to determine what the rate law would be for each step as the rate limiting step.
If Step I is limiting, the rate would vary with [BrO3–][ H3O+]
If Step II is limiting, the rate would vary with [BrO3–][ H3O+]2
If Step III is limiting, the rate would vary with [BrO3–][ H3O+]2
If Step IV is limiting, the rate would vary with [BrO3–][ Br–][ H3O+]2
If Step V is limiting, the rate would vary with [BrO3–][ Br–]2[ H3O+]2
4.
(12%) There is great current interest in developing fuel cells based on the reaction,
2CH3OH(l) + 3O2(g) r 2CO2(g) + 4H2O(l)
a. Write a balanced equation for the half-reaction that occurs in acid solution for such a fuel cell at the;
i. anode.
ii. cathode.
b. If the E˚ value for the cell reaction is 1.21 V, calculate the value of ∆G˚.
c. The E˚ value for the O2(g) half reaction is 1.23 V in 1 M H+, calculate the E˚ value expected in 1 M OH–.
d. State two advantages of carrying out this reaction in a fuel cell rather than burning methanol and converting the heat into
electricity.
a. i. anode: CH3OH + H2O → CO2 + 6H+ + 6e–
ii.) cathode: O2 + 4H+ + 4e– → 2H2O
o
o
b. ΔG = –nFE = –(12 mol)(96500 J/V⋅mol)(1.21 V) = –1.40×103 kJ
#
&
RT % 1 (
o
ln
c. Use the Nernst equation: E = E "
nF %% H + 4 ((
$
'
'
(8.314 J/mol # K)(298 K) $
1
0.0257
ln& "14 4 ) = 1.23"
ln 10 56 = 0.40 V
so, E = 1.23"
(4)(96500)
4
% (10 ) (
[ ]
( )
Page 2
!
!
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
d. 1. No wasted heat. 2. No energy lost during conversion.
5.
(12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and
omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances.
You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated.
a. Solid calcium is heated in nitrogen gas.
b. Solid sodium ethoxide is added to water.
c. Solutions of magnesium sulfate and barium hydroxide are mixed.
d. An acidic potassium permanganate solution is added to a solution of sodium sulfite.
e. Radium-222 undergoes alpha decay.
f. 2-propanol is heated with concentrated sulfuric acid.
a. Ca(s) + N 2 ( g) "
"# Ca 3 N 2 ( s)
b. NaOCH 2CH 3 ( s) + H 2O "
"# Na + ( aq) +OH - ( aq) + CH 3CH 2OH
!
!
!
c. Mg 2+ ( aq) + SO 42- ( aq) + Ba 2+ ( aq) + 2OH - ( aq) "
"# BaSO 4 (s) + Mg(OH) 2 (s)
d. MnO 4- ( aq) + H + ( aq) + SO 32- ( aq) "
"# Mn 2+ ( aq) + H 2O + SO 42- ( aq)
e.
222
"# 42 He
88 Ra"
+
218
86 Rn
H 2SO 4
f. CH 3CH(OH)CH 3 " " "# CH 3CH = CH 2 + H 2O
!
6.
! (12%)
a. Explain why many chemical reactions that are nonspontaneous, with ∆G˚ > 0 at room temperature, proceed to a
!
significant extent at that temperature.
b. Account for the fact that standard enthalpies of formation of compounds at 25˚C may be either positive or negative.
c. Explain why all elements and compounds have positive S˚ values at 25˚C.
d. Give an example of a chemical species that does not have a positive S˚ value at 25 ˚C and explain why its standard
entropy is not positive.
a. ΔGo values refer to standard conditions including 1 M concentrations. Reactions that are nonspontaneous under these
conditions may be caused to occur by increasing the concentration of the reactants and/or decreasing the concentrations of
the products.
b. ΔHfo values of compounds are relative to their elements in standards states (for which ΔH fo = 0). Depending on the
compound, formation may either release energy (ΔHfo < 0) or absorb energy (ΔHfo > 0).
c. The standard for entropy, S˚, is a perfect crystal at 0 K, which by the Third Law of Thermodynamics is zero. As
temperature increases, entropy increases, so S˚ is positive at 25˚C.
d. S˚ values of many ions (such as F–, Cl–, PO43–) are less than zero, This occurs because the reference for aqueous ions is
the standard entropy for H+, which is set to zero. Some ions, like those listed, may organize the solvent molecules more
than the hydrogen ions, so their standard entropy will be negative.
7.
(12%) Account for the following observations in terms of atomic/ionic/molecular properties.
a. Sodium fluoride melts at a higher temperature than potassium chloride.
b. Titanium(III) chloride is a solid at room temperature but titanium(IV) chloride is a liquid at room temperature..
c. N2O3 is an acidic anhydride but Bi2O3 is a basic anhydride.
d. Lithium chloride is much more soluble in ethanol than is sodium chloride.
a. The internuclear distance in NaF is less than that in KCl so the lattice energy of NaF is greater. Overcoming larger lattice
energy leads to higher melting points.
b. TiCl3 has Ti3+ ions at the center and is an ionic compound whereas for TiCl4, the smaller Ti4+ ion causes the Ti–Cl bonds to
have more covalent character. Covalent molecules typically melt at lower temperatures than ionic ones.
c. N2O3 reacts with H2O to form HONO, the high electronegativity of N draws electrons from H of H-O bond to give H+.
Bi2O3 reacts with H2O to form Bi(OH)3. The less electronegative Bi3+ bonds less strongly to O so OH– is released.
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
Page 5
d. The smaller Li+ ion has a higher charge density than the larger Na+ ion. This makes LiCl more covalent than NaCl (for
the same reasons noted in part (a)). Covalent compounds are more soluble in C2H5OH which has lower polarity.
8.
(14%) Four compounds with a molar mass of 59 have the formula C3H9N and the structures:
CH3CHNH2
CH3CH2CH2NH2
CH3CH2NHCH3
(CH3)3N
CH3
a. Name the class to which these compounds belong.
b. The boiling points of the four compounds vary from 3˚C to 46˚C. Identify the lowest and highest boiling compounds and
account for the difference in terms of the intermolecular forces in each.
c. Each of the four compounds is basic. For one of the compounds draw a structural formula for the conjugate acid formed
with H+. Account for the observation that all of these compounds are more basic than ammonia.
d. There are two amides with the formula C2H5NO and the same molar mass as the above compounds.
i. Draw structural formulas for these two compounds.
ii. State whether these compounds have boiling points above or below 46˚C. Rationalize your prediction.
iii. State whether these compounds are more or less basic than those with the structures given above. Rationalize your
prediction.
a. These molecules are amines.
b. (CH3)3N is the lowest boiling, CH3CH2CH2NH2 is the highest boiling. (CH3)3N is lowest because it has no
hydrogen bonding interactions. The remaining three all have hydrogen bonding, so the highest boiling will have
the largest dispersion forces, which the longer chain alkane provides.
c. Any of these four drawings would count…
CH3
H
H
H
CH3
CH2
CH2
N
CH3
H
CH
N
H
CH3
CH2
N
H
CH3
N
H
H
CH3
CH3
CH3
H
These compounds are more basic than NH3 because the carbon containing groups are better electron donors than
hydrogen. This inductive effect causes the lone pairs on the nitrogen to be donated to H+ more readily.
d. i.
O
O
CH3
C
N
H
H
C
N
CH3
H
H
or
ii. These compounds have higher boiling points because they will hydrogen bond more strongly.
iii. They will be less basic. The C=O functional group will draw electrons from the nitrogen making the lone pairs
less available.
END OF PART II
Page 2
Not valid for use as an USNCO Olympiad National Exam after April 29, 2009.
amount of substance
ampere
atmosphere
atomic mass unit
atomic molar mass
Avogadro constant
Celsius temperature
centi- prefix
coulomb
electromotive force
energy of activation
enthalpy
entropy
ABBREVIATIONS AND SYMBOLS
n equilibrium constant
K measure of pressure mmHg
A Faraday constant
F milli- prefix
m
atm formula molar mass
M molal
m
u free energy
G molar
M
A frequency
ν mole
mol
NA gas constant
h
R Planck’s constant
°C gram
P
g pressure
c heat capacity
k
Cp rate constant
C hour
Rf
h retention factor
E joule
s
J second
Ea kelvin
c
K speed of light
H kilo- prefix
T
k temperature, K
S liter
t
L time
volt
V
CONSTANTS
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
USEFUL EQUATIONS
E = E! –
! k2 $ Ea ! 1 1 $
=
'
" k1 &% R #" T1 T2 &%
" –!H % " 1 %
ln K = $
' $ ' +c
# R & # T&
RT
ln Q
nF
ln #
PERIODIC TABLE OF THE ELEMENTS
1
H
2
He
1.008
4.003
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Uun
111
Uuu
112
Uub
114
Uuq
(223)
(226)
(227)
(261)
(262)
(263)
(262)
(265)
(266)
(269)
(272)
(277)
(2??)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
237.0
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)
Page 7
2009 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Brookline High School, Brookline, MA
Chair
Linda Weber, Natick High School, Natick, MA
John Mauch, Braintree High School, Braintree, MA
Mathieu Freeman, Greens Farms Academy, Greens Farms, CT
Erling Antony, Arrowhead Union High School, Hartland, WI
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for
safety procedures and the proper disposal of chemicals at your examining site.
Not valid for use as an USNCO National Examination after April 29, 2009
Page 1
2009 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD
PART III — LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection at all times during this laboratory practical. You
also must follow all directions given by your examiner for dealing with spills and with disposal of wastes.
Lab Problem 1
You have been given six numbered pipets containing 0.50M solutions of the sodium salts Na2CO3, NaHCO3,
NaHSO3, NaH2PO4, Na2HPO4, Na3PO4, not necessarily in this order, a 50-mL beaker containing 0.40M HCl,
and a pipet containing methyl orange indicator. Devise and carry out an experiment to determine to contents
of each pipet, providing both qualitative and quantitative data to justify your conclusions.
Lab Problem 2
You have been given a thermometer, styrofoam cup with lid, a beaker, a graduated cylinder, and access to
room temperature water, heated water and ice cubes. Using these materials, design and carry out an
experiment to determine the heat of fusion, Hf, for water.
Page 2
Not valid for use as an USNCO National Examination after April 29, 2009
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School: _______________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name: ________________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Examiner’s Initials:
2. Record your data and other observations.
Not valid for use as an USNCO National Examination after April 29, 2009
Page 3
(data and observations – continued)
3. Based on your observations, write the relevant equations that led to your conclusions:
4. Conclusions
Pipet #
Page 4
Contents
Justification
Not valid for use as an USNCO National Examination after April 29, 2009
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School: _______________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name: ________________________________ Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
Before beginning your experiment, you must get
approval (for safety reasons) from the examiner.
Not valid for use as an USNCO National Examination after April 29, 2009
Examiner’s Initials:
Page 5
2. Record your data and other observations.
3. Calculations and Conclusions.
4. Conclusions: The Hf for water is: _____________
5. Sources of Error in this experiment (please number)
Page 6
Not valid for use as an USNCO National Examination after April 29, 2009
2009 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
Examiner's Instructions
Directions to the Examiner:
Thank you for administering the 2009 USNCO laboratory practical on behalf of your Local Section. It is
essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be
considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do
not lend any assistance or hints whatsoever to the students once they begin work. As in the international competition,
the students are not allowed to speak to anyone until the activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab station or
table when the students enter the room. The equipment should be initially placed so that the materials used for Lab
Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad
Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a
laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are
instructed to do so.
You will be asked to complete two laboratory problems. Students are to work alone. All the materials and equipment
you may want to use to solve each problem has been set out for you and is grouped by the number of the problem.
Students can use all materials for both lab problems, but each experiment is designed to work best with equipment
and materials provided specifically for each lab problem. You will have one hour and thirty minutes to complete the
two problems. You may choose to start with either problem. You are required to have a procedure for each problem
approved for safety by an examiner. (Remember that approval does not mean that your procedure will be
successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety
approval, raise your hand.
Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any
chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for
disposing of solutions at the end of this lab practical are:
We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the
directions on the front page. There is a periodic table and constants on the last page.
Are there any questions before we begin?
Not valid for use as an USNCO National Examination after April 23, 2009
Page 1
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table
on the last page of the booklet. Allow students enough time to read the brief cover directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all
information at the top of the answer sheets. Are there any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your
cooperation during this test.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required information
on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the
students. They can learn about possible observations and interpretations and you can acquire feedback as to what
they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete
the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad subcommittee as they
prepare next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from
Part I, and the “Blue Books” from Part II in the UPS return envelope you were provided to this address:
ACS DivCHED Exams Institute Department of Chemistry Iowa State
University0213 Gilman Hall Ames, IA 50011
The label on the envelope should have this address already, you will need only to include your return address and
call United Parcel Service -UPS (1-800-742-5877) for it to be picked up (or it can be dropped in a UPS collection
box). The cost of shipping will be billed to the Exams Institute. You can write down the tracking number on the
label to allow you to track your shipment.
Wednesday, April 29, 2009, is the absolute deadline for receipt of the exam materials at the Examinations
Institute. Materials received after this deadline CANNOT be graded. Be sure to have your envelope picked up no
later than April 28, 2009 for it to arrive on time.
THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR
GRADING THIS EXAMINATION.
Not valid for use as an USNCO National Examination after April 23, 2009
Page 2
Examiner’s List: 2009 USNCO Lab Practical Equipment and Chemicals
USNCO 2009 PART III: EXAMINER’S NOTES
Lab Problem #1: Materials and Equipment
Lab Problem #1 Each student will have:
Materials
• Six Beral-style pipets to be filled with the unknown sodium salts.
• One 150-mL or 250-mL beaker for holding the six pipets containing the unknown solutions.
• A twelve-hole clear well plate
• Several toothpicks for stirring
• Two empty Beral-style pipets
• One Beral-style pipet to be filled with methyl orange indicator.
• Access to sink and running tap water.
Chemicals
• Six filled Beral-style transfer pipets each containing 0.50M solution of the following
sodium salts: Na2CO3, NaHCO3, NaHSO3, NaH2PO4, Na2HPO4, Na3PO4
IMPORTANT: Use this key to number the pipets:
Na3PO4
NaHSO3
Na2CO3
NaH2PO
#1
#2
#3
#4
NaHCO3
Na2HPO4
#5
#6
4
Vial
• A single Beral-style transfer pipet of methyl orange indicator, filled and labeled ‘methyl orange’.
Methyl Orange Indicator: 0.1% aqueous solution.
• Approximately 20-30 mL of 0.40M HCl, in a 50-mL beaker labeled ‘0.40M HCl’.
Notes to Coordinators
• Place the numbered pipets upside down in the small beaker at the student’s desk.
• Make sure the solutions are freshly made (especially the NaHSO3 solution) prior to the lab.
• Obviously, do NOT label the pipets with the chemical formulas.
Not valid for use as an USNCO National Examination after April 23, 2009
Page 3
Lab Problem #2 Each student will have:
Materials
• One Celsius thermometer (-20oC to 100 oC range is sufficient)
• One standard size (8 oz.) styrofoam cup with tight-fitting lid.
• One 25-mL graduated cylinder
• One 250-mL beaker
Chemicals
• Access to tap water, ice cubes, and hot water.
Notes to Coordinators
• Make sure that students can easily access the ice cubes and hot water. Ideally, they can obtain each and return to
their work areas using their 250-mL beakers. The hot water should be heated to approximately 60-70 oC.
• Students might inquire about the need for an electronic balance to determine the mass of the water, but assuming
that the density of water is 1 g/mL eliminates the need for determining masses here. Do NOT tell students this – it is
for them to figure out.
Note that the examiner will need to initial each student’s experimental plan. Please do not comment on
the plan other than looking for any potentially unsafe practices.
Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A
lab coat or apron for each student is desirable but not mandatory. You will also need to give students
explicit directions for handling spills and for disposing of waste materials, following approved safety
practices for your examination site. Please check and follow procedures appropriate for your site.
Not valid for use as an USNCO National Examination after April 23, 2009
Page 4
2009 USNCO Part III Lab Practicals Answers
Lab Problem #1
Key:
Vial
Na3PO4
#1
NaHSO3
#2
Na2CO3
#3
NaH2PO4
#4
NaHCO3
#5
Na2HPO4
#6
This problem tests students’ understanding of acids and bases, titration, differentiation of mono-protic, di-protic, and
tri-protic acids in titration, and some qualitative observations in acid-base neutralization.
A sample data table quantifying the HCl added might look like this:
sample
# drops HCl
Justifications / Observations
Na3PO4
33
No bubbles, color ∆ from red to orange-yellow
Na2HPO4
18
No bubbles, color ∆ from red to orange-yellow
NaH2PO4
1
No bubbles, color ∆ from red to orange -yellow
NaHSO3
8
Sharp, biting odor, very slight bubbling
NaHCO3
15
Bubbles, red to orange-yellow color change
Na2CO3
29
Bubbles, red to orange-yellow color change
Excellent Student Results:
• Student presented an organized plan to add a fixed number of drops of each of the unknowns, a fixed number of
drops of the methyl orange indicator to each unknown, and then add HCl drop wise until a color change occurs.
Student planned to note the color changes, drops added, and a detectable odor for the sulfite ______ solution.
• Student showed a carefully constructed data table OR written account of data collected including all color
changes, number of drops added to reach color changes, extent of bubbling including size of bubbles, and odor
produced. Multiple trials were performed.
• Student included each of the six ionic equations that correlated to the reactions performed from data table,
correctly indicating ion charges, states symbols, and stoichiometric relationship.
• Student correctly identified each of the six unknowns and gave justifications that were consistent with the data
taken in this experiment.
Average Student Results:
• Student presented a plan but might not have indicated the odor to identify the sulfite or was not clear about how
to distinguish the bisulfite from the sulfite.
• Student might have written most of the equations correctly but neglected to include ion charges or states
symbols. One trial was performed.
• Most of the identifications were correct; most of the justifications were valid.
Below Average Student Results:
• Student did not connect the procedure with the conclusions. Plan was vague or unclear about how the added HCl
would be used to conclude unknowns. Plan was difficult to follow.
• Data was unorganized or difficult to follow.
Not valid for use as an USNCO National Examination after April 23, 2009
Page 5
•
•
Student neglected to include or incorrectly wrote chemical equations.
A majority of the unknowns were incorrectly identified.
Notes: The sulfite solution should have been made fresh. Bisulfite is easily oxidized in water and older solutions
that were used take little or no acid added to cause the color change and distinguishing odor.
Lab Problem # 2
This is a calorimetry problem. The novelty in this question was the intended lack of access to a balance. Students
were to make assumptions about the mass of ice based on measured volumes of water prior to adding ice, and the
final volume after the ice had been added. Students also had to account for the heat required to increase the
temperature of the cold water from the melted ice.
Conservation of energy applies such that q(ice) + q(cold water) = q(hot water). Once the ice begins to melt (don’t
spill, and put the cap on quickly to minimize heat loss) in the calorimetry cup, the hot water becomes colder. Since
you don’t end at 0 oC, the equation breaks into:
[mice x Hfus] + [mice-water x Cp x Tfinal – initial ] = [mhot water x Cp x Tfinal – initial ].
Then solve for Hfus
Below are sample student data:
10.0 gice Hfus + 10.0 gice-water (4.184) 27.0 – 0.0) = -[ 100ghotwater (4.184) (27.0 – 40.0)
10.0 gice Hfus + 1130 J = 5440 J
10.0 gice Hfus = 4310 J
Hfus = 431 J/g, about a 30% error based on 334 J/g as a correct value.
The question did not specify the units for the reported answer. Other acceptable values include: 80.0 cal/g, 6.01
J/mol, 1440 cal/mol, 6.01 kJ/mol, or 1.44 kcal/mol.
Excellent Student Results:
• Student included initial and final temperatures, determined quantity of ice used, water used, carried out replicate
determinations. The coffee cup/lid were constructed as a calorimeter in this experiment.
• Data and calculations clearly used the overall heat lost-heat gained equations.
• Students provided a cogent discussion of the major sources of error such as heat lost or gained by the system,
inaccuracies of the measuring equipment used.
Average Student Results
• Student did not perform multiple determinations.
• Student did perform more than one trial, but averaged observations rather than results.
• Student might have neglected the heating of the water from the melted ice to the Tfinal in their calculations.
Below Average Student Results
• Student neglected to provide any measure (volume or converted mass) for ice used.
• Student did not clearly use the heat lost-heat gained equation to solve for heat of fusion.
• Student did not provide ‘heat loss’ or uncertainty in temperature measurement as sources of error, or provided
frivolous sources such as ‘spilled water’.
Not valid for use as an USNCO National Examination after April 23, 2009
Page 6
2010 U.S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM PART I
Prepared by the American Chemical Society Chemistry Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, Chair, State University of New York, Cortland, NY
James Ayers, Mesa State College, Grand Junction, CO
Sherry Berman-Robinson, Consolidated HS, Orlando Park, IL (retired)
Seth Brown, University of Notre Dame, Notre Dame, IN
Peter Demmin, Amherst HS, Amherst, NY (retired)
Marian Dewane, Centennial HS, Boise, ID
Xu Duan, Queen Anne School, Upper Marlboro, MD
Valerie Ferguson, Moore HS, Moore, OK
Julie Furstenau, Thomas B. Doherty HS, Colorado Springs, CO
Kimberly Gardner, United States Air Force Academy, CO
Regis Goode, Ridge View HS, Columbia, SC
Paul Groves, South Pasadena HS, South Pasadena, CA
Preston Hayes, Glenbrook South HS, Glenbrook, IL (retired)
David Hostage, Taft School, Watertown, CT
Dennis Kliza, Kincaid School, Houston, TX
Adele Mouakad, St. John's School, San Juan, PR
Jane Nagurney, Scranton Preparatory School, Scranton, PA
Ronald Ragsdale, University of Utah, Salt Lake City, UT
DIRECTIONS TO THE EXAMINER-PART I
Part I of this test is designed to be taken with a Scantron answer sheet on which the student records his or her responses. Only this
Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the
student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until
April 26, 2010, after which tests can be returned to students and their teachers for further study.
Allow time for students to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer
sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has
elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper.
There are three parts to the National Chemistry Olympiad Examination. You have the option of administering the three parts in any
order, and you are free to schedule rest breaks between parts.
Part I
60 questions
single answer, multiple-choice
1 hour, 30 minutes
Part II
8 questions
problem-solving, explanations
1 hour, 45 minutes
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron
answer sheet to be scored. Be sure to write your name on the answer sheet, an ID number is already entered for you. Make a record
of this ID number because you will use the same number on Parts II and III. Each item in Part I consists of a question or an
incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes
the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on
the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you
complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron
answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today.
Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036
All rights reserved. Printed in U.S.A.
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
amount of substance
ampere
atmosphere
atomic mass unit
Avogadro constant
Celsius temperature
centi– prefix
coulomb
density
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
1
1A
1
H
1.008
3
Li
n
A
atm
u
NA
°C
c
C
d
E
Ea
H
S
K
ABBREVIATIONS AND SYMBOLS
Faraday constant
F molar mass
free energy
G mole
frequency
ν Planck’s constant
gas constant
R pressure
gram
g rate constant
hour
h reaction quotient
joule
J second
kelvin
K speed of light
kilo– prefix
k temperature, K
liter
L time
measure of pressure mm Hg vapor pressure
milli– prefix
m volt
molal
m volume
molar
M
CONSTANTS
M
mol
h
P
k
Q
s
c
T
t
VP
V
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
PERIODIC TABLE OF THE ELEMENTS
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
18
8A
2
He
4.003
10
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
19
K
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
11
1B
29
Cu
12
2B
30
Zn
26.98
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
113
114
115
116
117
118
(223)
(226)
(227)
(261)
(262)
(266)
(264)
(277)
(268)
(281)
(272)
(Uut)
(Uuq)
(Uup)
(Uuh)
(Uus)
(Uuo)
58
Ce
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
(277)
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
232.0
Page 2
59
Pr
Uub
231.0
238.0
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
DIRECTIONS
ƒ
When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2
pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully.
ƒ
There is only one correct answer to each question. Any questions for which more than one response has been blackened will not
be counted.
ƒ
Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question.
1. A student prepares a 100 mL aqueous solution
containing a small amount of (NH4)2SO4 and a second
100 mL solution containing a small amount of NaI, then
mixes the two solutions. Which statement describes what
happens?
6. Four elements were tested in the laboratory and gave the
results in the table below. Which element is a metalloid?
Element
Appearance
Conductivity
A
High
B
Slight
luster
Shiny
(B) Both compounds dissolve initially but NH4I
precipitates when the solutions are mixed.
C
Dull
None
(C) Both compounds dissolve initially but Na2SO4
precipitates when the solutions are mixed.
D
Shiny
High
(A) Both compounds dissolve and remain in solution
when the two solutions are mixed.
(D) The NaI dissolves but the (NH4)2SO4 does not.
There is no change upon mixing.
2. A colored gas is observed with which combination?
(A) calcium hydride and water
(B) lead metal and nitric acid
Low
(A) Element A
(B) Element B
(C) Element C
(D) Element D
7. What is the molarity of Na+ ions in a solution made by
dissolving 4.20 g of NaHCO3 (M = 84.0) and 12.6 g of
Na2CO3 (M = 126) in water and diluting to 1.00 L?
(C) sodium carbonate and sulfuric acid
(A) 0.050 M
(B) 0.100 M
(D) zinc sulfide and hydrochloric acid
(C) 0.150 M
(D) 0.250 M
3. Mixing which pair of 0.10 M solutions produces two
precipitates that cannot be separated from one another by
filtration?
(A) aluminum chloride and copper(II) nitrate
(B) strontium bromide and lead(II) acetate
(C) magnesium perchlorate and lithium carbonate
(D) barium hydroxide and copper(II) sulfate
(B) O2
(C) CO2
(D) CH4
5. For aqueous solutions of which of the following
substances could the concentration be determined by
visible spectrophotometry?
I Cr(NO3)3
II KMnO4
III Zn(NO3)2
(A) I only
(B) III only
(C) I and II only
(D) I, II, III
Page 3
8. Which solute has the greatest solubility (in mol/L) in
water at 25 ˚C and 1 atm?
(A) CH4
(B) NH3
(C) AgCl
(D) CaSO4
9. Which 2.00 M solution can be used to separate Al3+ from
Fe3+ in an aqueous solution?
(A) HCl
(B) H2SO4
(C) NaCl
(D) NaOH
10. The percent composition of the high explosive HNS is
4. Which gas turns limewater, a saturated solution of
Ca(OH)2, cloudy?
(A) H2
Behavior
with HCl
Bubbles
slowly
No
reaction
No
reaction
Bubbles
rapidly
C
H
N
O
37.35%
1.34%
18.67%
42.65%
The molar mass of HNS is 450.22. What is the molecular
formula of HNS?
(A) C13H4N7O12
(B) C14H6N6O12
(C) C15H10N6O11
(D) C16H12N5O11
11. A student prepares four 0.10 M solutions, each
containing one of the solutes below. Which solution has
the lowest freezing point?
(A) CaCl2
(B) KOH
(C) NaC2H3O2
(D) NH4NO3
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
12. What is the molarity of a hydrochloric acid solution if
20.00 mL of it neutralizes 18.46 mL of a 0.0420 M
Ba(OH)2 solution?
19. Calculate ∆E when one mole of liquid is vaporized at its
boiling point (80 ˚C) and 1 atm pressure.
(A) 0.0194 M
(B) 0.0388 M
[∆Hvap = 30.7 kJ/mol]
(C) 0.0455 M
(D) 0.0775 M
(A) 33.6 kJ
13. Ar and He are both gases at room temperature. How do
the average molecular velocities (V) of their atoms
compare at this temperature?
(A) VHe = 10VAr
(B) VAr = 10VHe
(C) VHe = 3VAr
(D) VAr = 3VHe
14. Lithium reacts with water to produce hydrogen gas and
lithium hydroxide. What volume of hydrogen collected
over water at 22˚C and 750 mm Hg pressure is produced
by the reaction of 0.208 g of Li? [VPH2O = 19.8 mm Hg]
(A) 367 mL
(B) 378 mL
(C) 735 mL
(D) 755 mL
15. Correct statements about samples of ice and liquid water
at 0 ˚C include which of the following?
I Molecules in ice and liquid water have the same
kinetic energy.
II Liquid water has a greater entropy than ice.
III Liquid water has a greater potential energy than ice.
(A) I and II only
(B) I and III only
(C) II and III only
(D) I, II, and III
16. A sample of a volatile liquid is introduced to an
evacuated container with a movable piston. Which
change occurs as the piston is raised? (Assume some
liquid remains.)
I The fraction of the molecules in the gas phase
increases
II The pressure in the container decreases
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
17. The kinetic energy of the molecules in a sample of H2O
in its stable state at –10 ˚C and 1 atm is doubled. What
are the initial and final phases?
(A) solid → liquid
(B) liquid → gas
(C) solid → gas
(D) solid → solid
18. Barium metal crystallizes in a body-centered cubic lattice
with barium atoms only at the lattice points. If the density
of barium metal is 3.50 g/cm3, what is the length of the
unit cell?
(A) 3.19 × 10–8 cm
(B) 4.02 × 10–8 cm
(C) 5.07 × 10–8 cm
(D) 6.39 × 10–8 cm
Page 4
(B) 31.4 kJ
(C) 30.0 kJ
(D) 27.8 kJ
20. Use the following data to calculate the molar enthalpy of
combustion of ethane, C2H6.
2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l) ∆H = –2511 kJ/mol
C2H2(g) + 2H2(g) → C2H6(g)
∆H = –311 kJ/mol
2H2(g) + O2(g) → 2H2O(g)
∆H = –484 kJ/mol
(A) –1428 kJ/mol
(B) –2684 kJ/mol
(C) –2856 kJ/mol
(D) –3306 kJ/mol
21. A 10.00 g piece of metal is heated to 80.00 ˚C and placed
in 100.0 g of water at 23.00 ˚C. When the system has
reached equilibrium the temperature of the water and
metal are 23.50 ˚C. What is the identity of the metal?
[Specific heat capacity of H2O = 4.184 J/g ˚C]
(A) Ag (Cp 0.236 J/g ˚C)
(B) Cu (Cp 0.385 J/g ˚C)
(C) Fe (Cp 0.449 J/g ˚C)
(D) Al (Cp 0.901 J/g ˚C)
22. For a reaction at constant pressure to be spontaneous,
which relationship must be correct?
(A) ∆Hrxn < 0
(B) ∆Grxn < 0
(C) ∆Srxn < 0
(D) ∆Suniv < 0
23. Tungsten is obtained commercially by the reduction of
WO3 with H2 according to the equation:
WO3(s) + 3 H2(g) → W(s) + 3 H2O(g)
The following data related to this reaction at 25 ˚C are
available.
H2O(g)
WO3(s)
∆H˚ kJ/mol
–840.3
–241.8
∆G˚ kJ/mol
–763.5
–228.5
The temperature at which this reaction is at equilibrium at
1 atm is closest to which of the following?
(A) 124 K
(B) 213 K
(C) 928 K
(D) 2810 K
24. The gaseous compound NOBr decomposes according to
the equation
NOBr(g)
NO(g) + 1/2 Br2(g)
At 350 K the equilibrium constant, Kp, is 0.15. What is the
value of ∆G˚?
(A) –5.5 × 103 J/mol
(B) –2.4 × 103 J/mol
(C) 2.4 × 103 J/mol
(D) 5.5 × 103 J/mol
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
25. The rate of decomposition of a certain compound in
solution is first order. If the concentration of the compound
is doubled, what happens to the reaction's half-life?
(A) It doubles
30.
(B) It decreases to ½ of the original value
(C) It decreases to ¼ of the original value
(D) It remains the same
26. Consider the reaction: 2 ICl(g) + H2(g) → 2 HCl(g) + I2(g)
At a certain temperature the rate constant is found to be
1.63 x 10–6 L/mol.s. What is the overall order of the
reaction?
(A) zero
(B) first
(C) second
(D) third
27. For the reaction: N2O4(g) → 2 NO2(g)
the number of moles of N2O4(g) is
Time, min
Moles
N2O4(g)
0
5
10
0.200
0.170
0.140
What is the number of moles of NO2(g) at t = 10 min?
(Assume moles of NO2(g) = 0 at t = 0.)
(A) 0.280
(B) 0.120
(C) 0.110
(D) 0.060
28. A compound decomposes with a first-order rate constant
of 0.00854 s–1. Calculate the concentration after 5.0
minutes for an initial concentration of 1.2 M.
(A) 0.010 M
(B) 0.093 M
(C) 0.92 M
(D) 1.1 M
29. Ozone in the earth's atmosphere decomposes according
to the equation: 2 O3(g) → 3 O2(g)
This reaction is thought to occur via the two-step
mechanism:
O2(g) + O(g) Fast, reversible
Step 1 O3(g)
Step 2 O3(g) + O(g) → 2 O2(g) Slow
What rate law is consistent with this mechanism?
(A) –∆[O3]/∆t = k[O3]
(B) –∆[O3]/∆t = k[O3]2
(C) –∆[O3]/∆t = k[O3]2/[O2]
(D) –∆[O3]/∆t = k[O3]2/[O2]3
The rates of many substrate reactions catalyzed by
enzymes vary with time as shown. Which factor(s) best
account(s) for the constant reaction rate after a certain
time?
I The enzyme's active sites are filled.
II The amount of substrate is constant.
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
31. Consider the system at equilibrium:
NH4HS(s)
NH3(g) + H2S(g)
∆H > 0
Factors which favor the formation of more H2S(g)
include which of the following?
I adding a small amount of NH4HS(s) at constant
volume
II increasing the pressure at constant temperature
III increasing the temperature at constant pressure
(A) I only
(B) III only
(C) I and II only
(D) I and III only
32. A 2.0 L container is charged with a mixture of 6.0 moles
of CO(g) and 6.0 moles of H2O(g) and the following
reaction takes place:
CO2(g) + H2(g)
CO(g) + H2O(g)
When equilibrium is reached the [CO2] = 2.4 M. What is
the value of Kc for the reaction?
(A) 16
(B) 4.0
(C) 0.25
(D) 0.063
33. Determine K for the reaction:
H2C2O4(aq) + 2OH–(aq) Æ C2O42–(aq) + 2 H2O(l)
H2C2O4(aq) Ka1 = 6.5 × 10–2
H2O Kw = 1.0 × 10–14
Ka2 = 6.1 × 10–5
(A) 4.0 × 10–34
(B) 4.0 × 10–6
(C) 4.0 × 106
(D) 4.0 × 1022
34. Which range includes the value of the equilibrium
constant, Keq, for a system with ∆G˚ << 0?
(A) –1 < Keq < 0
(B) 0 < Keq < 1
(C) Keq < –1
(D) 1 < Keq
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
Page 5
35. What volumes of 0.200 M HNO2 and 0.200 M NaNO2
are required to make 500. mL of a buffer solution with
pH = 3.00? [Ka for HNO2 = 4.00 x 10–4]
(A) 250. mL of each
(B) 143 mL of HNO2 and 357 mL of NaNO2
(C) 200. mL of HNO2 and 300. mL of NaNO2
(D) 357 mL of HNO2 and 143 mL of NaNO2
36. A sample of sparingly soluble PbI2(s) containing
radioactive I-133 is added to 0.10 M KI(aq) and stirred
overnight. Observations about this system include which
of the following?
I The radioactivity of the liquid phase increases
significantly.
II The concentration of the I– ion in solution increases
significantly.
(A) I only
(B) II only
(C) Both I and II
(D) Neither I nor II
37. An unknown metal, M, and its salt, M(NO3)2, are
combined with a half-cell in which the following reaction
occurs:
Ag+(aq) + e– → Ag(s) [E˚red = 0.80 V]
If E˚cell = 1.36 V, what is E˚ red for M2+(aq) + 2e– → M(s)?
(A) 0.56 V
(B) 0.24V
(C) –0.24V
(D) –0.56V
38. Given the standard reduction potentials:
O2 + 4H+ + 4e– → 2H2O
Br2 + 2e– → 2Br –
2H+ + 2e– → H2
Na+ + e– → Na
E˚ = 1.23 V
E˚ = 1.08 V
E˚ = 0.00 V
E˚ = –2.71 V
What products are formed in the electrolysis of 1 M
NaBr in a solution with [H3O+] = 1 M?
(A) Na(s) and O2(g)
(B) Na(s) and Br2(g)
(C) H2(g) and Br2(g)
(D) H2(g) and O2(g)
39. According to the standard reduction potentials:
Pb2+(aq) + 2e– Æ Pb(s)
Fe2+(aq) + 2e– Æ Fe(s)
Zn2+(aq) + 2e– Æ Zn(s)
E˚ = –0.13 V
E˚ = –0.44 V
E˚ = –0.76 V
Which species will reduce Mn3+ to Mn2+ [E˚ = 1.51 V]
but will NOT reduce Cr3+ to Cr2+ [E˚ = –0.40 V]?
(A) Pb only
(B) Zn only
(C) Pb and Fe only
(D) Pb, Fe, and Zn
40. Zn(s) / Zn2+(aq) // H+(aq) / H2(g) E˚ = 0.76 V
What must be the pH in the hydrogen compartment of the
cell designated above if the cell voltage is 0.70 V?
(Assume that both the [Zn2+] and the H2(g) pressure are
at standard values and T = 25 ˚C.)
(A) 0.51
(B) 1.01
(C) 2.50
(D) 3.21
41. The equilibrium constant, K, is 2.0 × 1019 for the cell
Ni(s) / Ni2+(aq) // Hg22+(aq) / Hg (l)
The value of E˚ at 25 ˚C for this cell is closest to
(A) –1.14V
(B) –0.57V (C) 0.57V
(D) 1.14V
42. In a battery with a zinc anode, what is the minimum mass
of zinc required if a current of 250 mA is drawn for 12.0
minutes?
(A) 0.0610 g
(B) 0.122 g
(C) 0.244 g
(D) 1.02 g
43. Which set of quantum numbers (n, l, ml, ms) is possible for
the outermost electron in a strontium atom in its ground
state?
(A) 5, 0, 0, –1/2
(B) 5, 0, 1, 1/2
(C) 5, 1, 0, 1/2
(D) 5, 1, 1, –1/2
44. How many orbitals are in an f sublevel (l = 3)?
(A) 3
(B) 5
(C) 7
(D) 14
45. What is the energy of photons with a wavelength of 434
nm?
(A) 2.76 × 105 kJ/mol
(B) 2.76 × 102 kJ/mol
(C) 2.76 × 10–1 kJ/mol
(D) 2.76 × 10–4 kJ/mol
46. In which choice are the species listed in order of
increasing radius?
(A) Na+, Mg2+, Al3+
(B) Cl–, S2–, P3–
(C) Ar, K+, Cl–
(D) Cl–, Ar, K+
47. Which element has the highest melting point?
(A) Na
(B) K
(C) Mg
(D) Ca
48. Which element forms a compound with the formula
H3XO4?
(A) As
(B) Cl
(C) N
(D) S
49. Which molecule contains the smallest F-S-F angle?
(A) SF2
(B) SOF2
(C) SO2F2
(D) SF6
50. Which species has the longest N–O bond?
(A) NO
Page 8
(B) NO+
(C) NO2
(D) NO2+
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
51. How many pi bonds and how many lone pairs are in the
Lewis structure of hydrazine, N2H4?
(A) 2 pi bonds, 0 lone pairs
59. All of the following statements concerning benzene, C6H6,
are correct EXCEPT
(A) Each carbon atom forms three sigma bonds.
(B) 1 pi bond, 0 lone pairs
(B) Each carbon is sp2 hybridized.
(C) 1 pi bond, 1 lone pair
(C) Pi electrons are delocalized over all 6 carbon atoms.
(D) 0 pi bonds, 2 lone pairs
(D) Benzene forms cis and trans isomers when it reacts.
52. In the Lewis structure of nitrous acid:
60. Which functional group is not commonly found in
nucleic acids?
What is the formal charge on nitrogen?
(A) –1
(B) 0
(C) +1
(D) +3
53. How many isomers of octahedral Co(NH3)3Cl3 are there?
(A) 2
(B) 3
(C) 4
(A) alcohol
(B) amine
(C) carboxylic acid
(D) dialkyl phosphate
END OF TEST
(D) 5
54. The bond angle in H2O is approximately 105˚ while the
bond angle in H2S is approximately 90˚. Which
explanation best accounts for this difference?
(A) H–S bonds are longer than H–O bonds.
(B) H–S bonds are less polar than H–O bonds.
(C) S has d orbitals available for bonding, O does not.
(D) O uses sp3 hybrid orbitals for bonding, S uses its 3p
orbitals.
55. Which compound exists in optically active forms?
(A) CH3CFCClCH3
(B) CH2FCH2CH2Cl
(C) CH2FCHClCH3
(D) CHF2CH2CH2Cl
56. What product results when 2-butene reacts with
chlorine?
(A) 2-chlorobutane
(B) 1,2-dichlorobutane
(C) 2,2-dichlorobutane
(D) 2,3-dichlorobutane
57. Which chloroalkane undergoes substitution with OH–
exclusively by an SN1 mechanism?
(A) (CH3)2CHCH2Cl
(B) (CH3)3CCl
(C) CH3CH2CHClCH3
(D) CH3CH2CH2CH2Cl
58. Which is a monosaccharide?
(A) fructose
(B) lactose
(C) maltose
(D) sucrose
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
Page 9
2010 U.S. National Chemistry Olympiad
National Exam Part I
KEY
Number
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Answer
A
B
D
C
C
B
D
B
D
B
A
D
C
B
D
A
C
C
D
A
B
B
C
D
D
C
B
B
C
A
Number
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
Not for use as an USNCO National Exam after April 26, 2010
Answer
B
A
D
D
D
A
D
C
A
B
C
A
A
C
B
B
D
A
D
C
D
B
A
D
C
D
B
A
D
C
2010 U.S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM - PART II
Prepared by the American Chemical Society Olympiad Examinations Task Force
OLYMPIAD EXAMINATIONS TASK FORCE
Arden P. Zipp, Chair, State University of New York, Cortland, NY
James Ayers, Mesa State College, Grand Junction, CO
Sherry Berman-Robinson, Consolidated HS, Orlando Park, IL (retired)
Seth Brown, University of Notre Dame, Notre Dame, IN
Peter Demmin, Amherst HS, Amherst, NY (retired)
Marian Dewane, Centennial HS, Boise, ID
Xu Duan, Queen Anne School, Upper Marlboro, MD
Valerie Ferguson, Moore HS, Moore, OK
Julie Furstenau, Thomas B. Doherty HS, Colorado Springs, CO
Kimberly Gardner, United States Air Force Academy, CO Regis Goode, Ridge View HS, Columbia, SC
Paul Groves, South Pasadena HS, South Pasadena, CA
Preston Hayes, Glenbrook South HS, Glenbrook, IL
David Hostage, Taft School, Watertown, CT
Dennis Kliza, Kincaid School, Houston, TX
Adele Mouakad, St. John's School, San Juan, PR
Jane Nagurney, Scranton Preparatory School, Scranton, PA
Ronald Ragsdale, University of Utah, Salt Lake City, UT
DIRECTIONS TO THE EXAMINER - PART II
Part II of this test requires that student answers be written in a response booklet with blank pages. Only this "Blue Book" is
graded for a score on Part II. Testing materials, scratch paper, and the "Blue Book" should be made available to the student only
during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 26, 2010,
after which tests can be returned to students and their teachers for further study.
Allow time for the student to read the directions, ask questions, and fill in the required information on the "Blue Book". When the
student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the "Blue Book", Part II
of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the
"Blue Book," and that the same identification number used for Part I has been used again for Part II.
There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and
you are free to schedule rest breaks between parts.
Part I
60 questions
single-answer multiple-choice
1 hour, 30 minutes
Part II
8 questions
problem-solving, explanations
1 hour, 45 minutes
Part III
2 lab questions
laboratory practical
1 hour, 30 minutes
A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use
non-programmable calculators.
DIRECTIONS TO THE EXAMINEE - PART II
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving
problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name,
the name of your school, and your identification number in the spaces provided on the "Blue Book" cover. (Be sure to use the same
identification number that was coded onto your Scantron sheet for Part I.) Answer all of the questions in order, and use both sides of
the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination.
When you complete Part II (or at the end of one hour and forty-five minutes) you must turn in all testing materials, scratch paper, and
your "Blue Book ". Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036 All rights reserved. Printed in U.S.A. Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 amount of substance
ampere
atmosphere
atomic mass unit
Avogadro constant
Celsius temperature
centi– prefix
coulomb
density
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
n
A
atm
u
NA
°C
c
C
d
E
Ea
H
S
K
ABBREVIATIONS AND SYMBOLS
Faraday constant
F molar mass
free energy
G mole
frequency
ν Planck’s constant
gas constant
R pressure
gram
g rate constant
hour
h reaction quotient
joule
J second
kelvin
K speed of light
kilo– prefix
k temperature, K
liter
L time
measure of pressure mm Hg vapor pressure
milli– prefix
m volt
molal
m volume
molar
M
CONSTANTS
M
mol
h
P
k
Q
s
c
T
t
VP
V
V
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
PERIODIC TABLE OF THE ELEMENTS
1
1A
1
H
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
1.008
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
113
114
115
116
117
118
(223)
(226)
(227)
(261)
(262)
(266)
(264)
(277)
(268)
(281)
(272)
(Uut)
(Uuq)
(Uup)
(Uuh)
(Uus)
(Uuo)
Page 2 Uub
(277)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 1.
(12%) 5.60 g of solid carbon is placed in a rigid evacuated 2.5 L container. Carbon dioxide is added to the container to a final pressure
of 1.50 atm at 298 K.
a. Calculate the number of moles of each reactant in the container originally.
2 CO(g) ∆H˚ = 173 kJ
b. The container is heated to 1100 K and the following reaction occurs: C(s) + CO2(g)
i. Calculate the pressure in the container at this temperature before the reaction takes place.
ii. When equilibrium is reached the pressure inside the container is 1.75 times that calculated in b.i. Determine the
equilibrium partial pressures of CO2(g) and CO(g).
iii. Write the equilibrium expression for this reaction, Kp.
iv. Calculate the value of Kp for this reaction at 1100 K.
c. Predict the effect on the number of moles of carbon monoxide of each of the following changes made to this system at
equilibrium. Give reasons for your predictions.
i. The volume of the container is increased to 5.0 L.
ii. The pressure inside the container is increased by adding helium.
iii. The temperature of the system is increased to 1200 K.
iv. The amount of solid carbon is increased to 6.00 g.
2.
(14%) Green plants utilize sunlight to convert CO2 and H2O to glucose (C6H12O6) and O2.
a. Write a balanced equation for this process.
b. Use the information in the accompanying table to calculate
i. ∆H˚
ii. ∆S˚
iii. ∆G˚ at 298 K for this reaction.
Substance
CO2(g)
H2O(l)
C6H12O6(s)
O2(g)
∆Hf˚ kJ/mol
–393.5
–285.8
–1273.3
S˚ J/mol⋅K
213.2
69.9
212.1
205.0
c. Comment on the spontaneity of this reaction at 25˚C and other temperatures.
d. Green plants use light with wavelengths near 600 nm for this process. Calculate
i. the energy of a 600 nm photon,
ii. ∆G˚ for the formation of one molecule of glucose by the reaction in 2a,
iii. the minimum number of 600 nm photons required to make one molecule of glucose by the reaction in 2a.
e. All of the photosynthesis on earth in a year stores 3.4 × 1018 kJ of solar energy.
i. Use the ∆G˚ for the photosynthetic reaction to calculate the number of moles of CO2 removed from the atmosphere by
photosynthesis each year.
ii. Determine the mass of carbon that is fixed annually by photosynthesis.
3.
(14%) A 0.125 g piece of vanadium reacts with nitric acid to produce 50.0 mL of a yellow solution of vanadium ions in their highest
oxidation state.
a. Calculate the number of moles of vanadium dissolved and the molarity of vanadium ions in this solution.
b. Write the electron configuration of a neutral gaseous vanadium atom.
c. Give the oxidation state of vanadium in the yellow solution and outline your reasoning.
d. A 25.0 mL portion of this yellow solution is reduced with excess zinc amalgam under an inert atmosphere to give a violet
–
solution. A 10.0 mL aliquot of this violet solution is titrated with a solution of 2.23 × 10 2 M KMnO4 in acid forming Mn2+.
–
A volume of 13.20 mL of the MnO4 solution is required to convert the vanadium back to yellow. Determine the:
–
i. number of moles of MnO4 used in this titration,
–
ii. mole ratio of vanadium ions to MnO4 ions in this titration,
iii. oxidation number change for vanadium in this titration and the oxidation state of vanadium ions in the violet solution.
e. When 2.00 mL of the violet solution are mixed with 1.00 mL of the original yellow solution, a green solution results. When this
ratio is reversed a bright blue solution is formed. Determine the oxidation states of the green and blue vanadium ions. Support your
answers with calculations.
Property of ACS USNCO ­Not for use as an USNCO National Exam after April 26, 2010 Page 3 4.
(12%) The reaction NO(g) + O3(g) → NO2(g) + O2(g) is first order in each reactant with an activation energy, Ea, of 11.7 kJ/mol and a
– –
rate constant of k = 1.2 × 1010 .L mol 1 s 1 at 25 ˚C.
–
a. Calculate the value of the pre-exponential factor, A, in the equation k = Ae Ea/RT.
b. Would the A factor for the chemical reaction NO(g) + N2O(g) → NO2(g) + N2(g) be expected to be larger or smaller than the A
factor in the above reaction if each reaction occurs in a single step? Outline your reasoning.
c. Calculate the rate constant for this reaction at 75 ˚C.
d. The following two-step mechanism has been proposed for this reaction:
Step 1
O3(g) → O2(g) + O(g)
NO(g) + O(g) → NO2(g)
Step 2
State and explain whether this mechanism is consistent with the observed rate law.
5. (12%) Write net equations for each of the reactions below. Use appropriate ionic and molecular formulas and omit formulas for all
ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the
equations.
a. Solutions of hydrochloric acid and silver acetate are mixed.
b. A small piece of potassium is added to water.
c. Concentrated hydrochloric acid is added to a solution of cobalt(II) sulfate.
d. An acidified potassium dichromate solution is added to a tin(II) chloride solution
e. Methyl ethanoate (methyl acetate) is reacted with a sodium hydroxide solution.
f. Carbon-14 undergoes beta decay.
6. (12%) Account for the following observations on the basis of electrochemical principles. The Standard Reduction Potentials are
provided.
–
E˚ = 1.61 V
2 HOCl(aq) + 2 H+(aq) + 2 e → Cl2(g) + 2 H2O(l)
–
–
E˚ = 1.36 V
Cl2(g) + 2 e → 2 Cl (aq)
–
O2(g) + 4 H+(aq) + 2 e → 2 H2O(l)
E˚ = 1.23 V
–
Cu (aq) + 2 e → Cu(s)
E˚ = 0.34 V
–
Sn (aq) + 2 e → Sn(s)
E˚ = –0.14 V
–
Fe (aq) + 2 e → Fe(s)
E˚ = –0.44 V
–
Zn (aq) + 2 e → Zn(s)
E˚ = –0.76 V
2+
2+
2+
2+
In a voltaic cell made with Cu metal in a 1.0 M CuSO4 and Zn metal in 1.0 M ZnSO4 the Zn is the anode and the cell
potential is more than 1.0 V. When aqueous sodium sulfide is added to the CuSO4 solution the cell potential decreases
substantially.
b. Iron metal corrodes readily in moist air but this corrosion can be prevented when iron is coated with tin or zinc. Corrosion is
prevented when the zinc coating is intact or broken. In contrast, corrosion is prevented by coating iron with tin only as long as the
tin coating remains intact but actually occurs faster when there is a break in the tin coating.
c. In acid solution chloride and hypochlorite ions react to form chlorine gas whereas in basic solution chlorine gas reacts to
form chloride and hypochlorite ions.
a.
7. (12%) Two stable allotropes of oxygen are dioxygen (O2) and ozone (O3).
a. Describe the geometry of ozone and state the hybridization of each of the oxygen atoms.
b. Ozone has a nonzero dipole moment. Account for this fact and predict the direction of the dipole moment.
c. Dioxygen is weakly attracted to strong magnetic fields (i.e. is paramagnetic), while ozone is weakly repelled by magnetic fields
(i.e. is diamagnetic). Account for these observations in terms of the bonding in the two molecules.
d. The most stable allotrope of sulfur is the cyclic S8 molecule while S2 is a highly unstable gas. In contrast, O2 is the most
stable allotrope of oxygen and O8 is unknown. Account for these differences in the relative stability of the allotropes of
these two elements.
8. (12%) There are four structural isomers with the formula C4H9Cl, one of which exists in optically active forms.
a. Write structural formulas for these four isomers.
b. Identify the isomer that exists in optically active forms and describe the difference in behavior of these two forms.
–
c. Each of these isomers reacts with OH ions to eliminate a molecule of HCl.
i. Give the name and molecular formula for the family of compounds formed by this elimination reaction.
ii. Write a structural formula for each of the elimination products.
iii. Identify the elimination product that can exist in different isomeric forms and draw structures for these forms.
Page 4 Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
2010 U.S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM - PART II - KEY
1
1mol
= .466mol
12.01g
CO2 n = PV RT n = (1.50atm )(2.5L ) / 298K (.0821) n = 0.153mol
(1100K )
P2 = 5.54atm
b. (i). P2 = P1 T2 / T1
P2 = 1.50atm
298K
(9.70atm)(2.5L)
n = PV RT
(ii). PT = 1.75(5.54) = 9.70atm
n=
= 0.268mol
1100K (.0821)
0.268 = 0.153 − x + 2x
x = 0.268 − 0.153
x = .115
nRT
(.038)(.0821)1100
= 1.3atm
n CO = 0.230
PCO 2 =
PCO 2 =
n CO 2 = 0.038
V
2.5L
(.230)(.0821)1100
PCO = 8.31atm
PCO =
= 8.31atm
2.5L
⎛ (.0821)(1100) ⎞
9.70atm = (0.153 − x ) + 2x⎜
0.269 = 0.153 + x
x = .116
⎟
2 .5
⎠
⎝
a. 5.60gC ×
(
(iii). K p =
PCO 2
)
PCO 2
(iv). K p = (8.31)
2
K p = 50.4
1.37
c. (i). nCO will increase. As V is ↑ , P ↓ s so system shifts → .
(ii). nCO does not change. He is not in K p so has no effect.
(iii). nCO will increase.. ΔΗ is positive so ↑ T will favor → .
(iv). nCO will not change. Solids do not affect equilibrium.
2
a. 6CO2+6H2O → C6H12O6+6O2
b. (i). ∆H˚ = –1273.3+0–[6(–393.5)+6(–285.8)]
= –1273.3–[–2361–1714.8] = –1273.3+4075.8 = 2802.5 kJ
(ii). ΔS˚ = 212.1+6(205.0)–[6(213.2+6(69.9)
= 212.1+1230.0–[1279.2+419.4] = 1442.1–1698.6 = –256.5 J/mol K
(iii). ΔG˚ = ΔH˚–TΔS˚
ΔG˚ = 2802.5 kJ–298(–.2565 kJ/mol)
ΔG˚ = 2802.5 kJ+76.44 = 2878.9 kJ/mol
c. Reaction is not spontaneous at 25˚C because ΔG˚>0
Reaction is not spontaneous at other Ts because ∆H˚>0 and ΔS˚<0
d. (i). Ε = hν
c = νλ
Ε = hc λ
(6.626 × 10 −34 )(3 × 108 )
600 × 10 − 9
E = 3.31 × 10 −19 J
kJ
1mol
×
= 4.78 × 10 − 21 kJ mol ec = 4.78 × 10 −18 J molec
(ii). ΔG o / molecule = 2878.9
mol 6.022 × 10 23 molec
J
1phot
×
= 14.4photons
(iii). # of photons = 4.78 × 10 −18
molec 3.31 × 10 −19 J
1molC6 H12 O 6
6molCO 2
e. (i). 3.4 × 1018 kJ yr ×
×
= 7.08 × 1015 molCO 2
2.88 × 103 kJ 1molC6 H12 O 6
E=
(ii). 7.08 × 1015 molC × 12.01 g mol = 8.50 × 1016 gC
3
1mol
= .00245mol
50.94g
.00245mol
M=
= 0.049M
.050L
b. V Z = 23 1s2 2s2 2p6 3s2 3p6 4s2 3d 3
c. V is in +5 oxid st. due to loss of 4s and 3d electrons.
d. (i). mol MnO 4− = 2.23 × 10 −2 mol L × .01320L = 2.94 × 10 −4 mol
a. 0.125gV ×
(0.010L)(.0491 mol L) 4.91 × 10 −4 1.67
=
=
1
2.94 × 10 − 4 mol
2.94 × 10 − 4
5 1.67 = 3.0 Δ for V
(iii). Mn goes from +7 → +2 Δ = 5
(ii).
e. 2.00( x − 2) = 1.00(5 − x )
1.00( x − 2) = 2.00(5 − x )
4
a. k = Αe
− Ea
2x − 4 = 5 − x
3x = 9
x = 3 green
x − 2 = 10 − 2x
3x = 12
x = 4 blue
11700 J
1.2 × 1010 = Αe 8.−314
( 298)
RT
V 2 + violet
1.2 × 1010 = Αe −4.722
Α = 1.2 × 1010 / .008894
Α = 1.35 × 1012
b. The A factor for NO and N2O would be smaller than that for NO and O3 because there are fewer geometric
arrangements involving NO and N2O molecules that could lead to a successful reaction. The probability of
successful reactions are lower for NO and N2O.
E ⎛1
k
1 ⎞
c. ln 2 = a ⎜⎜ − ⎟⎟
k1
R ⎝ T1 T2 ⎠
ln
k2
11700J ⎛ 1
1 ⎞
=
−
⎜
⎟
10
J
1.2 ×10
⎝ 298 348 ⎠
8.314
MolK
= 1407.3(.0033557 − .0028736 )
(
ln
k2
1.2 × 1010
)
= 1407.3 4.821 × 10 −4
k2
= 0.6785
= 1.971
1.2 × 1010
k 2 = 2.37 × 1010
d. This mechanism would give either R = k[O 3 ] if first step is the slow one or R = k
[ NO][O 3 ]
if second step
[O 2 ]
is slow since neither of these rate laws R = k[ NO][O 3 ] this can’t be the mechanism.
5
a. H + + Cl − + Ag + + C 2 H 3O 2− → AgCl + HC 2 H 3O 2
b. K + H 2 O → K + + OH − + H 2
c. H + + Cl − + Co 2 + + SO 24 − → CoCl 24 − + HSO −4
d. Cr2 O 72− + H + + Sn 2+ → Cr 3+ + Sn 4+ + H 2 O
e.
O
O
H3CCOCH3 + OH- → H 3CCO -
f.
6
+ H3COH
14
14
0
6 C→ 7 N + −1 β
E cell = 0.76 + 0.34 = 1.08V
a. Zn + Cu2+ → Zn2+ + Cu
22+
When S is added to Cu /Cu half cell CuS forms reducing [Cu2+], shifting the reaction to the left and
decreasing Ecell
b. Fe → Fe2+ +2e– is 0.44V so oxidation (corrosion) is spontaneous.
Covering surface with Sn or Zn prevents reaction with O2.
If Zn coating is broken, Zn will still oxidize preferential.
If Sn coating is broken, Fe will oxidize more readily
c. 2HOCl + 2H+ + 2Cl– →2Cl2 + 2H2O
Ε o = 0.25 spontaneous in acid.
+
In basic solution [H ] is very low so reaction shifts to the left and Cl2 forms Cl– and OCl–.
7
a.
O
O
↔
O
O
O
+
O
2
O
O
O
2
O3 is bent. Central O is Sp hybridized. All 3 are Sp hybridized in delocalized structure.
b. Formal charge for central O is +1, –½ for each of terminal Os in delocalized structure. (-1 and O in each
resonance form). DM has + end on central O.
c. O3 is diamagnetic because all e– are paired.
O2 is paramagnetic because it has 2 unpaired e– (M.O theory). KΚ 4 σ 2 σ*22s π 22 p π 22 p σ*2 π*1π*1
d. P orbitals overlap better in smaller O atoms. So double bond in O2 is stronger that 2 single bonds / e–-e–
repulsion between O atoms weakens single bonds. S is larger than O so p orbitals don’t overlap as well S-S .
Double bond is weaker that 2 single bonds. S-S bonds are longer so e–-e– repulsion is lower.
8
Cl
Cl
a.
C
C
C
C
Cl
C
C
C
C
C
C
C
C
CH 3
CH 3
Cl
b.
C
C
C
C
Occurs in optically active forms.
Cl
Cl
C 2 H5
C
Vs
CH 3
CH3
C
C2 H 5
H
H
They differ in the direction they rotate plane polarized light.
c. (i). Alkenes C4H8.
(ii).
C
C
C
C
C
C
C
C
C
(iii).
C
C
C
C
Exists in isomeric forms
CH 3
CH 3
C
H
CH 3
and
C
H
C
C
H
C
CH3
C
C
C
C
C
Cl
2010 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM – PART III
Prepared by the American Chemical Society Olympiad
Laboratory Practical Task Force
OLYMPIAD LABORATORY PRACTICAL TASK FORCE
Steve Lantos, Chair, Brookline High School, Brookline, MA
Linda Weber, Natick High School, Natick, MA
John Mauch, Braintree High School, Braintree, MA
Mathieu Freeman, Greens Farms Academy, Greens Farms, CT
Erling Antony, Arrowhead Union High School, Hartland, WI
Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL
DIRECTIONS TO THE EXAMINER–PART III
The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the
format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner.
This gives explicit directions for setting up and administering the laboratory practical.
There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop
between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the
examiner before the student begins that procedure.
Part III
2 lab problems
laboratory practical
1 hour, 30 minutes
Students should be permitted to use non-programmable calculators.
DIRECTIONS TO THE EXAMINEE–PART III
DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE
INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED.
There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do
them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to
have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You
are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that
you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner
for safety procedures and the proper disposal of chemicals at your examining site.
Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036
All rights reserved. Printed in U.S.A.
Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010
amount of substance
ampere
atmosphere
atomic mass unit
Avogadro constant
Celsius temperature
centi– prefix
coulomb
density
electromotive force
energy of activation
enthalpy
entropy
equilibrium constant
1
1A
1
H
n
A
atm
u
NA
°C
c
C
d
E
Ea
H
S
K
ABBREVIATIONS AND SYMBOLS
Faraday constant
F molar
free energy
G molar mass
frequency
ν mole
gas constant
R Planck’s constant
gram
g pressure
hour
h rate constant
joule
J reaction quotient
kelvin
K second
kilo– prefix
k speed of light
liter
L temperature, K
measure of pressure mm Hg time
milli– prefix
m vapor pressure
molal
m VP
volt
volume
CONSTANTS
M
M
mol
h
P
k
Q
s
c
T
t
R = 8.314 J·mol–1·K–1
R = 0.0821 L·atm·mol–1·K–1
1 F = 96,500 C·mol–1
1 F = 96,500 J·V–1·mol–1
NA = 6.022 × 1023 mol–1
h = 6.626 × 10–34 J·s
c = 2.998 × 108 m·s–1
0 °C = 273.15 K
V
V
PERIODIC TABLE OF THE ELEMENTS
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
K
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
1.008
4.003
10
Ne
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
8
8B
26
Fe
9
8B
27
Co
10
8B
28
Ni
11
1B
29
Cu
12
2B
30
Zn
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
113
114
115
116
117
118
(223)
(226)
(227)
(261)
(262)
(266)
(264)
(277)
(268)
(281)
(272)
(277)
(Uut)
(Uuq)
(Uup)
(Uuh)
(Uus)
(Uuo)
Page 2
Uub
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
2010 U. S. NATIONAL CHEMISTRY OLYMPIAD
PART III – LABORATORY PRACTICAL
Student Instructions
Introduction
These problems test your ability to design and carry out laboratory experiments and to draw conclusions
from your experimental work. You will be graded on your experimental design, on your skills in data collection,
and on the accuracy and precision of your results. Clarity of thinking and communication are also components
of successful solutions to these problems, so make your written responses as clear and concise as possible.
Safety Considerations
You are required to wear approved eye protection and gloves at all times during this laboratory
practical. You also must follow all directions given by your examiner for dealing with spills and with disposal
of wastes. In particular, special precautions are required when using the silver nitrate solution in Problem #2.
Do not get it on your skin or clothing as it will cause stains. Please wash off any chemicals spilled on your skin
or clothing with large amounts of tap water. On Problem #2, neither solution ‘AgNO3 solution’ nor ‘K2CrO4
solution’ can be disposed of down the drain. Both solutions need waste beakers (provided) in which the used
pipettes should be placed.
Lab Problem 1
You have been given a well plate, several test tubes and pipets, a concentrated ammonia solution, access to
distilled water, and four numbered vials containing iron (III) chloride hexahydrate, cobalt (II) sulfate
heptahydrate, copper (II) chloride dihydrate, and potassium oxalate monohydrate, though not necessarily in
this order. Devise and carry out an experiment to produce at least FIVE new different complex compounds,
using your understanding of coordination compound geometry and qualitative evidence in your results.
Lab Problem 2
You have been given a sample of seawater, a pipet that contains 5.0 × 10–4 mole/gram AgNO3, a pipet that
contains K2CrO4(aq), several small test tubes, and access to an electronic balance. Devise and carry out an
experiment to determine the percentage of chloride ion, Cl–(aq), in seawater.
Ksp at 25 °C
AgCl = 1.77 × 10–10
Ag2CrO4 = 1.12 × 10–12
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
Page 3
Answer Sheet for Laboratory Practical Problem 1
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name:_________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
2. Data and Observations.
Vials
#1
Contents as correctly written formulas
Evidence
#2
#3
#4
Before beginning your experiment, you must get
Approval (for safety reasons) from the examiner
Page 4
Examiner’s Initials:
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
3. Show any relevant reactions and sketch the possible geometry of each formed compounds.
4. Conclusions
Reactant(s) Used
Proposed coordinate complex formula
Evidence
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
Page 5
Answer Sheet for Laboratory Practical Problem 2
Student's Name: __________________________________________________________________________
Student's School:________________________________________ Date: ___________________________
Proctor's Name: _________________________________________________________________________
ACS Section Name:_________________________________Student's USNCO test #: ________________
1. Give a brief description of your experimental plan.
2. Record your data and all relevant balanced chemical equations.
Before beginning your experiment, you must get
Approval (for safety reasons) from the examiner
Page 6
Examiner’s Initials:
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
3. Calculations:
4. Conclusion. The percentage of chloride ion present in seawater =
5. List the assumptions that you made in this experiment:
Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010
Page 7
2010 U. S. NATIONAL
CHEMISTRY OLYMPIAD
NATIONAL EXAM—PART III
Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force
Examiner’s Instructions
Directions to the Examiner:
Thank you for administering the 2010 USNCO laboratory practical on behalf of your Local Section. It is
essential that you follow the instructions provided, in order to insure consistency of results nationwide.
There may be considerable temptation to assist the students after they begin the lab exercise. It is
extremely important that you do not lend any assistance or hints whatsoever to the students once they
begin work. As in international competition, the students are not allowed to speak to anyone until the
activity is complete.
The equipment needed for each student for both lab exercises should be available at his/her lab station or
table when the students enter the room. The equipment should be initially placed so that the materials used
for Lab Problem 1 are separate from those used for Lab Problem 2.
After the students have settled, read the following instructions (in italics) to the students.
Hello, my name is ________ . Welcome to the lab practical portion of the U.S. National Chemistry
Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to
reason through a laboratory problem and communicate its results. Do not touch any of the equipment in
front of you until you are instructed to do so.
You will be asked to complete two laboratory problems. All the materials and equipment you may
want to use to solve each problem has been set out for you and is grouped by the number of the problem.
You may use equipment from one problem to work on the other problem, but the suggested ideal
equipment and chemicals to be used for each problem has been grouped for you. You will have one hour
and thirty minutes to complete the two problems. You may choose to start with either problem. You are
required to have a procedure for each problem approved for safety by an examiner. (Remember that
approval does not mean that your procedure will be successful – it is a safety approval.) When you are
ready for an examiner to come to your station for each safety approval, please raise your hand.
Safety is an important consideration during the lab practical. You must wear goggles and gloves
at all times. In particular, special precautions are required when using the silver nitrate solution in
Problem #2. Do not get it on your skin or clothing as it will cause stains. Please wash off any chemicals
spilled on your skin or clothing with large amounts of tap water. On Problem #2, neither solution
‘AgNO3 solution’ nor ‘K2CrO4 solution’ can be disposed of down the drain. Both solutions need waste
beakers (provided) in which the used pipettes should be placed.
The appropriate procedures for disposing of solutions at the end of this lab practical are:
____________________________________________________________________________________
____________________________________________________________________________________
____________________________________________________________________________________
We are about to begin the lab practical. Please do not turn the page until directed to do so, but
read the directions on the front page. Are they any questions before we begin?
Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes
Page 1
Distribute Part III booklets and again remind students not to turn the page until the instruction is given.
Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic
table on the second page of the booklet. Allow students enough time to read the brief cover directions.
Do not turn to page 2 until directed to do so. When you start to work, be sure to fill out all of the
information at the top of the answer sheets. Are they any additional questions?
If there are no further questions, the students should be ready to start Part III.
You may begin.
After one hour and thirty minutes, give the following directions.
This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for
your cooperation during this portion of the exam.
Collect all the lab materials. Make sure that the student has filled in his or her name and other required
information on the answer sheets. At this point, you might wish to take a few minutes to discuss the lab
practical with the students. They can learn about possible observations and interpretations and you can
acquire feedback as to what they actually did and how they reacted to the problems. After this discussion,
please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely
useful to the USNCO subcommittee as they prepare for next year’s exam.
Please remember to return the post-exam Questionnaire, the answer sheets form Part III, the Scantron
sheets from Part I, and the ‘Blue Books” from Part II in the overnight return envelope you were provided
to this address:
American Chemical Society
U.S. National Chemistry Olympiad Office
1155 16th Street, NW
Washington, DC 20036
The label on the UPS Express Pak envelope should have this address and your return address already. The
cost of the shipping is billed to ACS - USNCO. You can keep copy of the tracking number to allow you to
track your shipment.
Wednesday, April 28, 2010, is the absolute deadline for receipt of the exam material. Materials received
after this deadline CANNOT be graded. Be sure to have your envelope sent no later than Tuesday, April
27, 2010 for it to arrive on time.
THERE WIL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE
FOR GRADING THIS EXAMINATION.
Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes
Page 2
Lab Problem #1: Materials and Equipment
Each student should have available the following equipment and materials:
Materials
• Three standard size (18 or 20 × 150mm) test tubes
• One 150 or 250 mL beaker to hold the test tubes
• Four 5 mL capacity Beral-style pipets
• Four pieces of waxed weighing paper
• One 100 or 150 mL Erlenmeyer flask with stopper
• One 12-hole white spot plate (porcelain or polyurethane)
• Several wooden stirring sticks, the kind used for stirring coffee
• Access to distilled water, preferably individual bottles at each lab station
Chemicals
• Four capped numbered vials (20–30 mL capacity) containing approximately 3–5 g of each of the
following solids: FeCl3⋅6H2O, CoSO4⋅7H2O, CuCl2⋅2H2O, and K2C2O4⋅H2O
Vial
#1
#2
#3
#4
Substance
FeCl3⋅6H2O
CoSO4⋅7H2O
CuCl2⋅2H2O
K2C2O4⋅H2O
Obviously, do not identify or label the contents of each vial!
• 25–30 mL of concentrated ammonia solution per student. To make this solution for a group of 25–30
students, using stock (14.8M) ammonium hydroxide, NH4OH(aq) mix 250 mL of the ammonium
hydroxide with 500 mL water to make a total volume of 750 mL ammonia solution. The 100 or 150 mL
flask should be labeled ‘Ammonia Solution’.
Lab Problem #1 Notes
• Be sure that the solids are powdered and dry. In the case of the iron (III) chloride you may want to
grind it fine as it has a tendency to clump.
• Students will be warned that the ammonia solution is pungent and to take caution as it is used.
Safety: It is your responsibility to ensure that all students wear safety goggles and gloves during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to
give students explicit directions for handling spills and for disposing of waste materials, following
approved safety practices for your examination site. Please check and follow procedures appropriate for
your site.
Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes
Page 3
Lab Problem #2: Materials and Equipment
Each student should have available the following equipment and materials:
Materials
• Three small (10 × 75 mm) test tubes
• One 50 mL beaker to hold the test tubes
• One 100 or 150 mL beaker to hold the labeled silver nitrate, seawater, and potassium chromate pipets
• Three 5 mL Beral–style pipet (graduated or ungraduated, thin or regular stem)
Chemicals
• Approximately 30 mL of seawater in a covered 100 or 150 mL flask labeled ‘seawater’
• Two 5 mL capacity Beral-style pipets filled with 5.0 ×10–4 mol AgNO3 per gram solution. To make this
solution, dissolve 2.13 g of solid AgNO3 in 25.0 mL DI water. Store in a dark, covered container until
ready to apportion to students. Pipets should be labeled ‘AgNO3 solution’. It should be clearly labeled
"Caution - Do not get on Skin or Clothes - Will cause Stains"
• One Beral-style pipet containing potassium chromate solution
• The concentration of this solution should be approximately 1M (Mr K2CrO4 = 194). This pipet should
be labeled ‘K2CrO4 solution’. To make this solution, dissolve 0.4 g K2CrO4 / 2 mL per student or 10 g
in 50 mL for 25 students.
• Neither solution ‘AgNO3 solution’ nor ‘K2CrO4 solution’ can be disposed of down the drain. Both
solutions need waste beakers in which the used pipettes are placed.
Lab Problem #2 Notes:
• For the ocean water, if you live near the coast, obtain a sample directly from the ocean. Let any solids
settle before apportioning to students. If you do not have an available source of ocean water, you may
simulate a sample by dissolving 2.80 g of NaCl in 100 mL of distilled water (don’t use tap water as it
likely already contains measurable amounts of chloride). You may also use common commercial
laboratory seawater preparations that include approximate percentages of minerals found in ocean
water.
• Make sure to use DI water in making the silver nitrate solution. Ideally, this solution is made just prior
to use the day of the exam.
• The concentration of the potassium chromate solution is not critical but should be approximately 1M.
Safety: It is your responsibility to ensure that all students wear safety goggles and gloves during the lab
practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to
give students explicit directions for handling spills and for disposing of waste materials, following
approved safety practices for your examination site. Please check and follow procedures appropriate for
your site.
Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes
Page 4
2010 USNCO Part III Answers
Lab Problem 1
Students were expected to create different complex compounds by combinations of the
hydrate compounds provided with the aqueous ammonia solution. Students might have
thought also to react the hydrates with an aqueous oxalate solution or to slowly add water to
samples of each hydrate. For example, adding H2O slowly to a small amount of copper(II)
chloride in one of the test tubes produces a color change as the dihydrate of this salt forms a
tetrahydrate.
Vial
1
2
3
4
Compound
Yellow-Orange, FeCl3 . 6H2O
Pinkish, CoSO4 . 7H2O
Green, CuCl2 . 2H2O
White, K2C2O4 . H2O
Compound
.
CuCl2 2H2O + H2O
CuCl2 . 2H2O +
ammonia solution
.
CoSO4 7H2O +
ammonia solution
.
FeCl3 6H2O +
ammonia solution
.
K2C2O4(aq) + FeCl3 6H2O
.
K2C2O4(aq) + CuCl2 2H2O
Possible Formula
.
CuCl2 4H2O
Cu(NH3)4
Observation
green Æ blue
deep blue color change
Co(NH3)6
bright blue-green color change
Fe(NH3)6
Brown-purple gelatinous ppt.
K3[Fe(C2O4)3]
K2[Cu(C2O4)2]
greenish color change
greenish color change?
Ammonia replaces water in these hydrated compounds to produce complexes that include
NH3 surrounding the central metal cation. Possible formulas might also include ligand
combinations of ammonia, water, and chloride or sulfate as part of the complex.
Possible sketches of these structures show the central Cu, Fe, and Co cations surrounded by
the NH3 ligands in tetrahedral, square planar geometry, or octahedral geometries.
Excellent Results:
Students attempted as many matches as possible, creating a legible and systematically
organized table to communicate their results. The nomenclature of the possible complexes was
clear and the formulas used the proper coordinate complex compound was accurate. Drawings
for the possible structures included three-dimensional representations of the complexes and
indicated how and where isomers might also be formed. It was drawn/noted that the oxalate
ion was a bidentate and three ions were able to coordinate with the iron (III) cation.
1
Average Results:
Students attempted to create and draw only the requested five complexes. A listing of these
combinations and the observed results was provided. At least one possible complex geometry
was shown.
Below Average Results:
Students created fewer than the requested five complexes. No systematic table or chart was
provided to indicate the possible compounds formed. One or no possible geometries were
provided. Errors in the nomenclature and formulas were evident.
Lab Problem 2
Sample Data:
Initial mass of the seawater and pipet
Final mass of the seawater and pipet
Initial mass of silver nitrate and pipet
Final mass of silver nitrate and pipet
3.121 g
2.243 g
2.820 g
1.918 g
Sample Student Calculations:
3.121 g – 2.243 g = 0.878 g seawater used
2.820 g – 1.918 g = 0.902 g silver nitration solution used
Molar masses:
NaCl = 58.44 g/mol
AgNO3
= 169.87 g/mol
0.902 g AgNO3 x 5.0 x 10-4 mol/g AgNO3 = 4.51 x 104 mole AgNO3
4.51 x 104 mole AgNO3 = mole Ag+ ; mole Cl4.51 x 104 mole Cl- x 35.5 g/mol Cl- = 1.60 x 10-2 g Cl1.60 x 10-2 g Cl-/0.878 g seawater x 100 = 1.82% Cl- in solution by mass
(A value of 1.80-1.90% is generally accepted)
Excellent Results:
At least two trials were attempted using both of the filled AgNO3 pipets and averaging results.
These titrations were completed in the small test tubes provided. A clear and organized table
showed the initial and final masses of both the seawater and silver nitrate pipets. Observations
of the silver chloride and the silver chromate formation were noted. Calculations were legible,
organized, and followed a logical sequence in order to determine the mass of chloride ion
2
present and the final mass percentage present. Use of the Ksp values was made to clearly show
the AgCl then Ag2CrO4 selective order of precipitation. Assumptions included the complete
precipitation of AgCl from solution, the equivalent molar ratio of Ag+: Cl-, the presence of a
drop or two of the chromate solution to the seawater to be sufficient to create a white/red color
change endpoint. Error included the subjective judgment of the endpoint to distinguish the
complete precipitation of the chloride from the beginning of the precipitation of the chromate,
and the additional drop or two overshooting of the endpoint to turn the solution entirely
reddish.
Average Results:
Only one trial was attempted. Leftover materials were not utilized to complete additional
experiments. Observations were made to determine the endpoint. Calculations show a logical
sequence to solve this problem, but may not have been clearly organized. Only one point was
included for the assumptions and error.
Below Average Results:
Not all of the materials were used. Calculations did not show evidence of logical sequence in
problem solving. Observations were incomplete. One or no points were made regarding error
or assumptions.
3