Download Practice Exam-Final Fall 2016 W-Ans

CHEM 1411 PRACTICE EXAM I (Chapters 1- 4)
Multiple Choice: Select one best answer.
1. Convert 0.3980 m to mm.
(first see how many significant figures are there in meters and convert to mm and see there are
same significant figures)
A)
B)
C)
D)
E)
2. A scientist obtains the number 0.0470056800 on a calculator. If this number actually has three
(3) significant figures, how should it be written?
A)
B)
C)
D)
E)
0.0470
0.047
0.4700
0.047006
0.47006
3. Perform the indicated operations and express the answer with the proper number of significant
digits. 28.1 cm + 0.53 cm + 75.321 cm =
(When adding or subtracting the least decimal place is used in the product)
A) 104
B) 1.04 x 102
C) 103.95
D) 103.951
E) 104.0
4. How many significant figures are appropriate to show in the result after carrying out the
operation below?
(When multiplying or dividing the least significant numbers are used in the product)
(223.7 + 0.27) ÷ 4.21 =
A) 1
B) 2
C) 3
D) 4
E) 5
5. What is the equivalent temperature for 98.6 Fahrenheit in Kelvin?
First convert oF to oC (oC = 5/9 x (oF-32) and then oC to K (K = oC + 273)
(a) 37
(b) 310
(c) 471.6
(d) 594.8
(e) -273.1
6. An unknown sample has a mass of 13.9g and a volume of 17.4mL What is its density (g/mL)?
Use equation d = m/v
(a) 0.799
(b) 1.04
(c) 3.16
(d) 4.62
(e) 5.07
7. Which of the following is a heterogeneous mixture?
(a) milk
(b) salt, water, and sugar
(c) oil and water
(d) 14-K gold ring
(e) air
8. Which of the following is a physical change?
A) Color of carpet faded (by sun light)
B) Zinc strip dissolves in vinegar (to produce hydrogen gas)
C) Burn a wood block
D) Water freezes
E) Iron rusts
9. How many protons, electrons, and neutrons are there in 72Br- ?
Isotopes with number on front top represent mass number. Mass # = # Protons + Neutrons
Look for atomic number of bromine from periodic table (top one). In a neutral molecule, # of
protons = # of electrons.
(a) 35p, 35e, 72n
(b) 34p, 35e, 37n
(c) 35p, 36e, 72n
(d) 36p, 34e, 72n
(e) 35p, 36e, 37n
10. Which of the following pairs would have similar chemical and physical properties?
(a) Ni, Mg
(b) H, Na
(c) C, Si
(d) Cu, Ca
(e) Cr, Cl
Hint: The elements from the same group (i.e. the vertical column) will have
similar physical and chemical properties.
11. Which of the following is a molecular compound?
(a) KCl
(b) CsF
(c) HCN
(d) AlBr3
(e) NaOH
Hint: Molecular compound is composed by nonmetal reacts with nonmetal atoms.
12. Which of the following is not a correct match?
(a) AlCl3, aluminum chloride
(b) H2O dihydrogen monoxide
(c) CrF3, chromium (III) fluoride
(d) HNO3, nitrous acid
(e) CuSO4.5H2O, copper (II) sulfate pentahydrate
Hint: Follow the rules of naming the compounds. Be able to differentiate the rules
between ionic and molecular compounds.
13. Which of the following pairs is a correct match?
(a) Ba3(PO4)2, barium (II) phosphate
(b) (NH4)2SO3, diammonium sulfate
(c) Na2O2, sodium oxide
(d) Ca(NO3)2, calcium nitrate
(e) CH3COOH, carbonic acid
Hint: Follow the rules of naming the compounds. Be able to differentiate the rules
between ionic and molecular compounds.
14. The atomic masses of 10B and 11B are 10.0129 amu (natural abundance 19.78%) and 11.0093
amu (natural abundance 80.22%), respectively. What is the average atomic mass of B?
(a) 9.467
(b) 9.966
(c) 10.042
(c) 10.504
(e) 10.810
Hint: You must convert the % into decimals first and then apply the formula and
memorize it: M = M1X1 + M2X2 = 10.0129x0.1978 + 11.0093x0.8022 = 10.810.
15. What is the molar mass for calcium nitrate, Ca(NO3)2?
(a) 44
(b) 56
(c) 87
(d) 93
(e) 164
Hint: Formula, Ca(NO3)2, indicates that there are 1 Ca, 1x2 = 2 N and 3x2 = 6 O Go to
the periodic table and locate the atomic mass for each atom. Thus the formula mass of
Ca(NO3)2 = 1x40 + 2x14 + 6x16 = 164
16. How many hydrogen atoms are there in 48.0 g of CH4?
(a) 1.81x1023 (b) 7.22x1024 (c) 6.02x1023 (d) 1.20x1025 (e) 4.70x1025
Hint: According to the chemical formula, one mole of CH4 contains 1 mole of C atoms
and 4 moles of hydrogen atoms. Thus, the mole of H = 4 x {mass of CH4/molar mass of
CH4}. When converting moles into atoms, it needs to multiply the Avogadro’s number,
that is, 6.02x1023. Thus the number of H atoms = moles of H atoms x Avogadro’s number
= (6.02x1023)x{4x(48.0/16.0)} = 7.22x1024
17. What is the mass percent (%) for O in SO2?
(a) 38.09
(b) 45.41
(c) 50.00
(d) 53.86
(e) 56.43
Hint: % O = (mass of 2 O/ mass of SO2) x 100% = {2x16/(1x32+2x16)}x100% = 50.0%
18. Which of the following is the empirical formula?
(a) C8H18
(b) CH3COOH
(c) C4H10O2
(d) NH4NO3
(e) B2H6
Hint: Empirical formula is the simplest integral ratio among each atom.
19. Chemical analysis shows the composition of a compound containing carbon, hydrogen,
chlorine, and oxygen, to be 37.84% C, 2.12% H, 55.84% Cl, and 4.20% O. What is its empirical
formula?
(a) CHClO
(b) C2HClO4
(c) C12H8Cl6O
(d) C12H8Cl6O4
Hint: The empirical formula is the simplest integral ratio of moles among each atom.
Here, there are four different kinds of atoms, C, H, Cl and O. Thus mole of C = 37.84/12
= 3.15; mole of H = 2.12/1 = 2.12; mole of Cl = 55.84/35.45 = 1.58; mole of O = 4.20/16
= 0.26.
20. Chemical analysis shows the composition of a compound containing carbon and hydrogen, to
be 80.00% carbon and 20% hydrogen and the molar mass is 30 g. What is its molecular formula?
(a) CH
(b) C2H4
(c) C2H6
(d) C6H12
(e) C10H22
Hint: The molecular formula is an integral multiple of empirical formula. That is, the
molar mass = empirical molar mass x integer. From C: H = 80.00/12 : 20.00/1 = 6.66: 20
= 1: 3. So the empirical formula is CH3 and the empirical molar mass of CH3 = 12x1+1x3
=15. So the integer = 30/15 = 2. Thus there are two empirical formulas in a molecular
formula. Therefore, the molecular formula is C2H6.
21. What are the coefficients respectively when the equation
__PH3 + __O2
→
__P2O5 + __H2O is balanced?
(a) 2, 2, 1, 3
(b) 2, 1, 3, 4
(c) 2, 3, 1, 2
(d) 2, 4, 1, 3
(e) 1, 1, 3, 3
22. In the reaction of Al(OH)3 with H2SO4, how many moles of water can be produced If the
reaction is begun with 5.500 mole of Al(OH)3?
→
2Al(OH)3 + 3H2SO4
(a) 2.50
(b) 4.75
Al2(SO4)3 + 6H2O
(c) 6.32
(d) 7.58
(e) 16.50
23. How many grams of H2O could be formed by the reaction of 16.0g of CH4 with 48.0g of O2?
CH4 + 2O2 →
CO2 + 2H2O
(a) 27.0
(b) 37.3
(c) 46.8
(d) 54.1
24. What is the excess reagent in the above reaction (Q.23) of CH4 and O2? How many grams of
the excess reagent were consumed?
(a) CH4, 12 grams
(b) O2, 12 grams
(c) CH4, 4 grams
(d) O2, 4 grams
(e) CH4, 8 grams
25. A 15.6 grams of benzene (C6H6) is mixed with excess HNO3 to prepare nitrobenzene
(C6H5NO2). After the reaction there are 15.6 grams of nitrobenzene produced. What is the
percent yield of nitrobenzene?
C6H6 + HNO3 → C6H5NO2 + H2O
(a) 34.3%
(b) 47.6%
(c) 58.9%
(d) 63.4%
(e) 71.2%
Hint: The amount of product calculated according to the road map is the theoretical yield
and the one obtained by weighing is the actual yield.
% yield = (actual yield/ theoretical yield) x 100%
Now from the road map, we calculate the theoretical yield (note that since the actual
yield given is in unit of grams, so we need to complete the entire road map).
÷ 78 g/mol
x 1C6H5NO2 /1C6H6
x 123g/mol
15.6 gC6H6-------0.2mole C6H6 --------------------- 0.2mole C6H5NO2 ------24.6g C6H5NO2
Thus the percent yield = {15.6 g / 24.6 g} x 100% = 63.41%
26. Which of the following is a weak electrolyte?
A) Sodium hydroxide solution
B) He gas
C) Sugar
D) Water
E) Potassium chloride
others examples of weak electrolytes: Aceti acid, HF, HNO2 , NH3, etc.
27. Which of the following is paired incorrectly?
A)
B)
C)
D)
E)
HF – strong acid
H2SO4 – strong acid
Ba(OH)2 – strong base
HCl – strong acid
NH3 – weak base
Examples of weak acids: aceti acid, HF, HNO2, phosphoric acid, etc.
Examples of weak bases: ammonia, many alkyl amines, , pyridine, etc..
28. Which pair of ions would not be expected to form a precipitate when dilute
solutions of each are mixed?
A)
B)
C)
D)
E)
Al3+, S2–
Pb2+, OH–
Ba2+, PO43–
Mg2+, SO42–
Pb2+, Cl–
All Group IA salts and all metal nitrates will not form a precipitate
29. In the reaction 2Cs(s) + Cl2(g)  2CsCl(s), Cl2 is
A)
B)
C)
D)
the reducing agent
the oxidizing agent
Oxidized
the electron donor
30. Which of the following redox reactions will occur according to activity series?
A) Cu(s) + 2HCl(aq) CuCl2(aq) + H2(g)
B) Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)
C) Fe(s) + ZnSO4(aq) FeSO4(aq) + Zn(s)
D) 2Al(s) + 3Ca(NO3)2(aq) 2Al((NO3)3(aq) + 3Ca(s)
Other example Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
31. How many grams of KOH are present in 35.0 mL of a 5.50 M solution?
A) 5.5
B) 10.8
C) 15.7
D) 17.8
E) 21.3
0.1925 x 56.1 =
31. What is the final concentration (in M) of a solution when water is added to 25.0
mL of a 0.866 M KNO3 solution until the volume of the solution is exactly 500.0 mL?
A) 0.0252
B) 0.0368
C) 0.0117
D) 0.0534
Hint: Mi x Vi = Mf x Vf,
Mf = Mi x Vi/Vf = 25.0ml*0.866M/500.0ml=0.0433M
E) 0.0433
32. What volume (in mL) of a 0.500 M HCl solution is needed to neutralize 10.0 mL of a 0.2000
M Ba(OH)2 solution?
A) 8.00
B) 4.00
C) 2.00
D) 1.00
E) 0.50
2HCl + Ba(OH)2 → 2H2O + BaCl2
M1 x V1 = M2 x V2
V1 = M2 x V2 / M1 = 0.2000 M x 10 ml/500 ml = 4 ml x 2
33. Amphoteric species acts like
a) An acid
b) Base
c) Like an acid or base
d) None of the above
CHEM 1411 PRACTICE EXAM 2 (Chapters 5- 8)
Part I: Multiple Choice: Select one best answer and bubble in the scantron. Part II:
Show your work
Part I: Each 4 points
1. For a particular process q = +14 kJ and w = -17 kJ. Which of the following
statements are false? ΔE = q + w
A) Heat flows from the surroundings to the system.
B) The process is endothermic.
C)
E = -3 kJ
D) The surroundings do work on the system.
E) None of the above are false.
2. What is the change in energy (J) of the gas when a gas expands and does P-V work
on the surroundings equal to 325 J and at the same time absorbs 127 J of heat from
the surroundings?
Hint: ΔE = q + w
A) -198
B) + 198
C) + 157
D) -157
E) + 452
3 kJ
3. C2H5OH(l) + 3O2(g
H=–
2(g) + 3H2O(l
For the combustion of ethyl alcohol as described in the above equation, which of
the following is true?
I. The reaction is exothermic.
II. The enthalpy change would be different if gaseous water was produced.
III. The reaction is not an oxidation–reduction one.
A) I, II
B) I, II, III
C) I, III,
D) III
E) only I
4. The H value for the reaction
is -90.8 kJ. How much
heat is released when 66.9 g Hg is reacted with oxygen? (molar mass of Hg 200.6)
A) 0.333 kJ
B) 6.07 x 103 kJ
C) 30.3 kJ
D) 90.8 kJ
E) none of these
5. What is the heat absorbed (in J) when 2.5 grams of water is heated from 22oC
to 98oC? The specific heat of water is 4.18 J/g. oC.
Hint: q = m x s x Δt = 2.5 x 4.18 x (98-22) = 794 J
A) 794
B) 88 1
C) 974
D) 102
E)
137
6. What is the wavelength of a photon of red light (in nm) whose fre
1014 Hz?
Hint: C = λν or λ = с/ν
A) 659 nm
B) 1.52 x 106 nm
C) 152 nm
D) 551 nm
E) none of these
7. Green light can have a wavelength of 478 nm. The energy of a photon of this light is
Hint: E = h ν or E = h c/λ
A) 1.08 x 10–31 J
B) 5.43 x 10–7 J
C) 4.16 x 10–19 J
D) 5.52 x 1014 J
E) 2.73 x 1018 J
8. Which of the following combinations of quantum numbers (n, l, ml, ms) is not
allowed?
Hint: l can have only 0, 1, 2, 3, ……(n-1), but no negative numbers.
A) 9, 8, -4, 1/2
B) 8, 2, 2, 1/2
C) 6, 5, -5, 1/2
D) 5, -3, -1, 1/2
E) All are allowed.
9. Which of the following could not be a valid ml quantum number for a 5d orbital?
A) 2
B) 3
C) –2
D) 1
E) 0
Hint: For any d, the l value has to be 2 (five orbitals -2, -1, 0, 1, 2). If it is 3 you
will have 7-orbitals which is f type
10. The electron configuration for the oxygen atom is:
A) 1s22s22p4
B) [He]2s4
C) [Ne]2s22p2
D) 1s22p4
E) none of these
Oxygen atomic number is 8, hence it has 8 electrons distributed as above.
11. How many unpaired electrons are there in an atom of sulfur?
A) 0
B) 1
C) 2
D) 3
E) 4
[Ne]3s23p4
Sulfur = 16
 
 

12. Consider the following electron arrangements. Which represents for N?
2s
2p
a)
 









b)


 






c)


 







d)
 








e)
 
A) option a
B) option b
C)
option c
D) option d
E)
option e


 

13. Specify the group of the periodic table in which each of the following elements is
found
(a) [Ne]3s2, (b) [Ne]3s23p2, (c) [Ne]3s23p5,
A) Group 1A, 5A, 8A
B) Group 2A, 3A, 7A
C) Group 2A, 4A, 7A
D) Group 2A, 5A, 8A
E) Group 1A, 3A, 6A
Number of valance electrons of an element is its group number
14. Which of the following atoms has the electron configuration as [Ar] 4s23d2?
A) Sc
B) V
C) Fe
D) Ti
E) Zn
Hint : Look for the 4th element after the Argon in your periodic table sheet.
15. Select the correct electron configuration for Ag (Silver).
A) [Kr]5s24d9
B) [Kr]5s14d10
C) [Kr]5s24p64d3
D) [Kr]5s14d9
E) [Kr]4d11
Ag (47), 11th element after Kr (36). Expected to have [Kr] 5s24d9. Exceptions for
regular electron filling are Cr, Cu, Ag and Au.
16. Which of the following pairs is isoelectronic?
A) Li+ and K+
B) Na+ and Ne
C) I– and Cl–
D) S2– and Ne
E) Al3+ and B3+
Na+ = 1s22s22p6 and Ne = 1s22s22p6
17. Of the following elements, which has the lowest first ionization energy?
A) Ba
B) Ca
C) Si
D) P
E) Cl
Ionization energy (I.E.) is the minimum energy (kJ/mol) required to remove an electron
from a gaseous atom in its ground state. It increases across the periods (L to R) and
decreases for groups from top to bottom.
18. Which of the following arrangements is in order of increasing size?
A) Ga3+ < Ca2+ < K+ < Cl– < S2–
B) S2– > Cl– > K+ > Ca2+ > Ga3+
C) Ga3+ > S2– > Ca2+ > Cl– > K+
D) Ga3+ > Ca2+ > S2– > Cl– > K+
E) Ga3+ > Ca2+ > S2– > K+ > Cl–
Cations are smaller and anions are bigger. Among cations, higher the charge
smaller the cation. Among anions higher the charge bigger the anion.
19. Which of the following is a correct order of lattice energy?
A) Ga2O3 < CaO < K2O
B) LiF < LiCl < LiBr
C) Na2O < MgO < Al2O3
D) NaCl < AlCl3 < MgCl2
E) LiCl < NaCl < KCl
Lattice Energy
• Elattice increases as Charge on ions increases
• Elattice decreases as inter-ionic distances increase
20. Which of the following bonds is least polar?
A) C—O
B) H—C
C) S—Cl
D) Cl—Cl
E) They are all nonpolar.
Nonpolar bond: two identical atoms form a bond & equally share the
bond’s electron pair and there is no polarization of the bond.
21. Consider the following molecules. Which molecules violate the octet rule?
I.
II.
III.
IV.
V.
A)
B)
C)
D)
E)
BF3
CHBr3 (C is the central atom)
Br2
BeCl2
SF4
I, II, IV
I, III, IV, V
III, V, V
I, IV, V
I, II, IV, V
22. The formal charges in Nitromethane CH3NO2
A) Positive charge on oxygen and negative charge on nitrogen
B) Positive charge on nitrogen and negative charge on oxygen
C) Only negative charge on oxygen and zero charge on nitrogen
D) Only positive charge on nitrogen and zero charge on oxygen
E) none of them
23. Which of the following have resonance structure
A) NH4 +
B) CO2
C) AlCl3
D) H2O
E) NO3-
24. Which of the following is an ionic bond?
(a) The HN bond in NH3.
(b) The SiSi bond in Cl3SiSiCl3.
(c) The CaF bond in CaF2.
(d) The OO bond in O2.
(e) The NN bond in H2NNH2.
Part II: Show your work: Each 7 points
25. Calculate the heat of combustion (kJ/mol) for the following reaction:
2H2S(g) + 3O2(g)  2H2O(l) + 2SO2(g)
The standard enthalpies of formation for H2S(g), H2O(l), SO2(g) and O2(g) are
-20.15, -285.8, -296.4 and 0.0, respectively.
ΔH0rxn = ∑ n ΔH0f (products) - ∑ n ΔH0f (reactants)
= [2 H2O + 2SO2 ] – [2H2S]
= [(2 x -285.8) + (2 x -296.4)] – [2 x -20.13]
= [(-571.6) + (-592.8)] – [-40.3] = -1124
A) +562
B) -1124
C) -1205
D) -562
E) +281
26. From the following data,
C(graphite) + O2(g)  CO2(g)
∆H = -393.5 kJ/mol
H2(g) + 1/2O2(g)  H2O(l)
∆H = -285.8 kJ/mol
2C2H6(l) + 7O2(g)  4CO2(g) + 6H2O(l) ∆H = -3119.6 kJ/mol
Calculate the enthalpy change (kJ/mol) for the reaction
2C(graphite) + 3H2(g)  C2H6(g)
4C(graphite) + 4O2(g)  4CO2(g)
4∆H = -393.5 kJ/mol
6H2(g) + 3 O2(g)  6H2O(l)
6∆H = -285.8 kJ/mol
4CO2(g) + 6H2O(l) 2C2H6(l) + 7O2(g) ∆H = +3119.6 kJ/mol
4C(graphite) + 6H2(g)  2C2H6(g)
∆H = -169.5 kJ/2 mol
A) -84.6
B) -42.3
C) +84.6
D) +42.3
E) -67.2
27. What is the wavelength of light associated with the n = 4 to n = 1 electron transition?
A)
B)
C)
D)
E)
3.97 x 10–25 nm
8.21 x 102 nm
9.75 x 106 cm
1.94 x 10–18 m
9.73 x 10–8 m
E = −hcRH (
1
1
2 –
2
nf
ni
)
E = h ν = hc/ λ or λ = hc/E
28. What is the ∆Hrxn (kJ) for the reaction H2(g) + I2(g)  2HI(g)?
The bond enthalpies for H-H, I-I and H-I are 436.4, 151.0 and 298.3 kJ/mol,
respectively.
A) -9.2
B) +9.2
C) -289.1
D) +289.1
E) 0.0
Bonus Question:
Write the Lewis structures of NCS- ion and show the preferred structure. (5 points)
Key to Practice Exam-2 (Chapters 5-8):
1-D,
2-A,
3-A,
4-C,
5-A,
6-D,
7-C,
8-D,
9-B,
10-A, 11-C,
12-A, 13-C, 14-D, 15-B, 16-B, 17-A, 18-A, 19-C, 20-D, 21-D, 22-B,
23-E, 24-C, 25-B, 26-A, 27-E, 28-A,
CHEM 1411 PRACTICE EXAM 3 (Chapters 9- 12)
Multiple Choice Questions: Select one best answer. Each 3 points
Essay Questions: Write best answer. Each 7 point
Bonus Question (1): 8 points
1. The molecular geometry of NCl3 is
A)
B)
C)
D)
trigonal pyramidal
octahedral
trigonal planar
bent
Tetrahedral
Trigonal pyramidal
2. What hybridization is predicted for the nitrogen atom in the NO3– ion?
A)
B)
C)
D)
E)
sp2
sp3
dsp3
d2sp3
none of these
# of atoms + # lone pairs on central atom = # of hybrid orbital (sum of the superscripts!)
or
# of electron groups on central atom = # of hybrid orbital
3. Which of the following has the electron pair geometry identical to its molecular
geometry?
A) CH4
B) SF4
C) XeF4
D) H2S
E) NH3
No lone pair electrons on carbon atom. Therefore, the electron pair geometry is
identical to its molecular geometry
ED Trigonal bipyramidal
Molecula geometry ---Square planar
3 2
ED geometry SP d ----Octahedral
4. How many pi bonds and sigma bonds are there in the tetracyanoethylene
molecul
N
C
C
N
C
N
C=C
N
A) 4, 9
C
B) 4, 5
C) 9, 4
D) 9, 9
E) 5, 4
5. Which of the following has the bond angle as 120o?
A) NH3 (107) TPy
B) PF3 (>107) TPy
D) AlCl3 (120)TPl
C) SF6 (90)OH
E) XeF4
6. What is the pressure of the gas (in atm) when 5.0 moles of CO gas are present in
a container of 20.0 L at 27 oC? Hint: PV = nRT Or P = nRT/V
A) 3.25
B) 6.15
C) 7.40
D) 9.30
E) 10.55
7. What is the volume (in L) of 88.0 g of carbon dioxide gas at STP?
Hint: PV = nRT Or V = nRT/P
A) 44.8
B) 53.7
C) 62.1
n = mass/molar mass (88/44) = 2
D) 74.6
E) 83.2
8. What volume (in L) does a sample of air occupy at 6.6 atm when 1.2 atm, 3.8 L of
air is compressed at constant temperature?
P1 x V1 = P2 x V2 Or V2 = P1 x V1 / P2 = 1.2 atm. X 3.8 L / 6.6 atm. = 0.69
A) 0.34
B) 0.57
C) 0.69
D) 0.77
E) 0.86
9. What is the final temperature (in K), under constant-pressure condition, when a
sample of hydrogen gas initially at 88oC and 9.6 L is cooled until its finial volume is
3.4 L?
V1/T1 = V2/T2 Or T2 = V2 x T1/V1 = 3.4L x 361/9.6 = 128
A) 31
B) 68
C) 94
D) 128
E) 260
10. What is the molar mass (g/mol) of 7.10 grams of gas whose volume is 5.40 L at
741 torr and 40 oC? (7.48g).
Hint: PV = nRT = mRT or M = mRT
M
A) 34.6
B) 70.3
C) 86.2
PV
D) 94.6
E) 102.3
= 7.10g*0.082 *(40+273.15)K/[(741/760)atm*5.40L]= 34.6
11. The combustion process for methane, major component of natural gas, is
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
If 15.0 moles of methane are reacted, what is the volume of carbon dioxide (in L)
produced at 23.0 oC and 0.985 atm?
Hint: PV = nRT Or V = nRT/P = 15 x 0.0821 x 296 / 0.985
A) 370.0
B) 430.7
C) 510.8
D) 630
E) 720
12. What is the total pressure (in atm) of the mixture when a 2.5-L flask at 15 oC
contains a mixture of nitrogen, helium, and neon gases at partial pressure of 2.0 atm
for nitrogen, 0.5 atm for helium, and 2.5atm for neon?
PT = P 1 + P 2 + P 3 + .
A) 3.0
. . .
B) 1.0
C) 4.0
D) 2.0
E) 5.0
13. A mixture of gases contains 0.31 mol CH4, 0.25mol C2H6 and 0.29 mol C3H8.
The total pressure is 1.50atm. Calculate the partial pressure of the CH4, C2H6 and
C3H8 gases
Pi = Xi PT
mole fraction (Xi ) = ni/nT
A) 0.23,0.35,0.56atm
B) 0.54,0.44,0.51atm
C) 0.5,0.1,0.5atm
D) 1, 0.25, 0.75atm
E) 2, 1, 3atm
14. In a body centered cubic lattice, how many atoms will be there in a unit cell
Hint: 8 x 1/8 +1 = 2
A. 1
B. 2
C. ½
D. 4
15. Which of the following molecules can form a hydrogen bond?
A. SiH4
B. C2H5OH
C. H2Te
D. CH4
Molecules in which a hydrogen is attached to highly electronegative atoms such as
N, O, and F will exhibit hydrogen bonding.
16. Nonpolar substances will condense mainly because of
A. dipole-dipole interactions
B. ion-dipole
C. Dispersion forces
D. hydrogen bonds
Nonpolar molecules have only dispersion forces. It is one of the intermolecular
forces.
17. Arrange CH4, CF4, CCl4, CBr4, CI4 (carbon tetraiodide) molecules in the order of
increasing melting points
A. CH4>CF4>CCl4>CBr4>CI4
B. CH4<CF4<CCl4<CBr4<CI4
C. CI4>CBr4>CCl4>CF4>CH4
D. CI4<CBr4<CCl4<CF4<CH4
Nonpolar molecules will have only dispersion forces and the dispersion forces
increase with increasing molar mass resulting in increasing melting point.
18. What type(s) of intermolecular forces exist between I2 and NO3A.
B.
C.
D.
Ion induced dipole forces and dispersion forces
Dipole-dipole induced forces only
Ion induced dipole forces only
D. Dispersion forces only
Nonpolar molecules will have only dispersion forces. However, if ions are added,
they will have ion induced dipole forces as well along with dispersion forces.
19. The hydrogen containing molecules of group 6A elements are H2O, H2S, H2Se,
and H2Te. Which one of these has highest boiling point?
A. H2O
B. H2S
C. H2Se
D. H2Te
Generally, it is expected that boiling point will increase with increasing molar
mass. However, if there is any hydrogen bonding is possible even the low
molecular weight compounds will have high boiling point compared to high
molecular weight compounds with no hydrogen bonding.
20. Some insects can walk on water because of the phenomenon, called
A. Viscosity
B. Surface tension
C. Velocity
D. A&C
Surface tension is an elasticity tendency of liquids that forces liquids to have
minimum surface area. High intermolecular forces result in high surface tension.
21. Arrange the following molecules in order of increasing boiling points
1. CH3CH2OH
2. H2O
3. CH3OCH3
4. CH3CH2CH3
Generally, hydrogen bonding in a molecule will result in high boiling point. Water
has one H-bond for a molar mass of 18, ethanol has one H-bond for a molar
mass of 46. Ether has no H-bonds but is a dipole molecule. Propane is nonpolar.
A. 1<2<3<4
B. 4<3<1<2
C. 2<3<1<4
D. 4<3<2<1
22. Diamonds are extremely stable owing to which of the following kinds of
bonding?
A. ionic bonding
B. network covalent bonding
C. hydrogen bonding
D. Dispersion forces
23. Ice floats on water because of
A. London forces
B. Surface tension
C. Hydrogen bonding
D. Covalent bonding
24. Water rises in capillary tube because
A.
B.
C.
D.
Adhesion grater than cohesion
cohesion grater than adhesion
Cohesion and adhesion are same
Non of the above
25. List the substances Ar, Cl2, CH4, and CH3COOH in order of increasing strength
of intermolecular attractions.
(a)
(b)
(c)
(d)
(e)
CH4 < Ar < CH3COOH < Cl2
Cl2 < CH3COOH < Ar < CH4
CH4 < Ar < Cl2 < CH3COOH
CH3COOH < Cl2 < Ar < CH4
Ar < Cl2 < CH4 < CH3COOH
26. Which of the following elements, if doped into silicon, would yield an n-type
semiconductor:
a) Zn
b) Ga
C) As
d) C
27. Polyacetylene is what type of material
a) an ionic solid
polymer
b) an alloy
c) addition polymer
d) condensation
28. In which of these substances is hydrogen bonding likely to play an important role
in determining physical properties:
a) methane (CH4)
c) methyl fluoride (CH3F)
b) hydrazine (H2NNH2)
d) hydrogen sulfide (H2S)
Show your work (Essay Questions):
1. Metallic iron crystallizes in a cubic lattice. The unit cell edge length is 287 pm.
The density of iron is 7.87 g/cm3. How many iron atoms are within a unit cell?
d=m/v
d = 7.87 g/cm3
Unit cell volume, v = (287 X 10-10cm)3 = 23639903 x 10-30 =2.364 x 10-23 cm3
m=dxv
m = 7.87g/cm3 x 2.364 x 10-23 cm3 = 1.86 x 10-22 g
Molar mass of Iron = 55.85 g
1.86 x 10-22 g x 1 mole
55.85g
x 6.023x1023 atoms
= 2 atoms
1 mole
2. Write the structures for following molecules and describe the hybridization,
molecular shapes, and bond angles
a) Methane CH4
SP3
b) Ethylene CH2=CH2
SP2
c) Acetylene CH=CH
SP
Total # of electron groups = number of hybrid orbital (sum of the superscripts!)
3. Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4. Explain why?
O
H
S
H
dipole moment
polar molecules
O
O
H
H
O
C
O
no dipole moment
nonpolar molecules
C
H
H
4. Explain the following and give examples:
a) Sigma bond
b) Pi bond
c) Resonance structure
Sigma bond: bonding density is along the internuclear axis-head to head
overlap of orbitals
Pi bond: bonding density is above and below the internuclear axis-Sideways
overlap of orbitals
Benzene, NO3- ion
5. Write the electron pair geometry and molecular geometry for the following molecules
and describe the shape of molecules, bond angles, and polarity.
a) BeCl2
b)BF3
c)CH4
d) NH3
6. What is the bond order for H2, He2, N2 molecules. Explain if He2 molecule can form.
# of bonding e  # of antibonding e
Bond order =
2
7. Define the (a) Ionization Energy and (b) Electron Affinity (c) electronegativity
What is the overall trend of these in the perodic table?
Ionization energy (I.E.) is the minimum energy (kJ/mol) required to remove an
electron from a gaseous atom in its ground state
Electron affinity (EA) is the energy released when an electron is accepted by an
atom in the gaseous state to form an anion.
Electronegativity is the ability of an atom to attract toward itself the electrons in
a chemical bond.
8. Write resonance structures for Carbonate-2 ion and benzene molecule.
Total # of Valence electrons = 4+ (3x6) +2 = 24
Or
9. Calculate the enthalpy change upon converting 1.00 mol of ice at –25 °C to steam at
125 °C under a constant pressure of 1 atm. The specific heats of ice, liquid water,
and steam are 2.03, 4.18, and 1.84 J/g-K, respectively. For H2O, ΔHfus = 6.01 kJ/mol
and ΔHvap = 40.67 kJ/mol.
AB: ΔH = (1.00 mol)(18.0 g/mol)(2.03 J/g-K) (25 K) = 914 J = 0.91 kJ
BC: ΔH = (1.00 mol)(6.01 kJ/mol) = 6.01 kJ
CD: ΔH = (1.00 mol)(18.0 g/mol)(4.18 J/g-K)(100 K) = 7520 J = 7.52 kJ
DE: ΔH = (1.00 mol)(40.67 kJ/mol) = 40.7 kJ
EF: ΔH = (1.00 mol)(18.0 g/mol)(1.84 J/g-K)(25 K) = 830 J = 0.83 kJ
The total enthalpy change is the sum of
the changes of the individual steps:
ΔH = 0.91 kJ + 6.01 kJ + 7.52 kJ + 40.7 kJ + 0.83 kJ = 56.0 kJ
10.
Determine the packing efficiency of a body centered cubic cell (a = 4r/√3)
assuming the edge length is 1
3
# of atoms x 4πr /3 x 100 = 68%
packing efficiency =
a
3
Key to Practice Exam-3, Fall 2016
1. A
7. A
2. A
8. C
3. A
9. D
4. D
10. A
5. D
11. A
6. B
12. E
13. B
14. B
15. B
16. C
17. B
18. A
19. A
20. B
21. B
22. B
23. C
24. A
25. C
26. C
27. C
28. B