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Chemistry 104 –1
Professor F. Fleming Crim
Spring, 2005
Name______________________________
TA________________________________
Section_____________________________
Final Exam
1. This exam contains 17 pages of questions and instructions, a page of equations, a list of standard
reduction potentials, a periodic table, and a set of constants and conversion factors.
2. There are 44 multiple choice questions. Circle the one best answer for each.
3. You have 120 minutes to work on the exam.
4. Put your name on each page.
White Exam – Page 2
Name_________________________
1. Which of the following molecules has a chiral center?
(1) 2–methylheptane
(2) 3–methylheptane
(3) 4–methylheptane
a.
b.
c.
d.
e.
1 only
2 only
3 only
1 and 2
2 and 3
2. Which of the following (non–cyclic) hydrocarbons has at least two π bonds?
a.
b.
c.
d.
e.
C4H8
C10H20
C8H18
C5H12
C3H4
3. What is the name of the following compound?
a.
b.
c.
d.
e.
pentane
cyclopentane
pentanone
pentanol
cyclopentene
White Exam – Page 3
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4. Which of the following hydrocarbons can have cis and trans isomers?
a.
b.
c.
d.
e.
butane
2–butyne
2–pentene
2–methylpentane
propene
5. The functional group RCOR' is characteristic of a(n) ________.
a.
b.
c.
d.
e.
ester
alcohol
aldehyde
ketone
amide
6. All of the following compounds contain a carbonyl functional group EXCEPT
a.
b.
c.
d.
e.
amides.
ethers.
aldehydes.
carboxylic acids.
ketones.
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7. Amino acids polymerize in condensation reactions to form a peptide bond between amino acid
molecules. Which of the following are possible dipeptides formed in the reaction of two amino acids
with side groups R and R′ respectively ?
a.
b.
c.
d.
e.
1 only
2 only
1, 2, and 4
1 and 2
1 and 3
8. Rate constants usually
a.
b.
c.
d.
e.
decrease with time.
increase with time.
decrease with increasing temperature.
increase with increasing temperature.
are independent of time and temperature.
White Exam – Page 5
Name_________________________
9. Which of the following relationships is correct for the reaction of peroxydisulfate ion with iodide ion?
S2O82–(aq) + 3 I–(aq) → 2 SO42–(aq) + I3–(aq)
a.
!
"[I – ]
"[I –3 ]
=–
"t
"t
#[I – ] #[I –3 ]
b. "3
=
#t
#t
c. "3
!
d. 3
"[I – ] "[S2O 2–
8 ]
=
"t
"t
e. –
1 "[I – ] 1 "[SO 2–
4 ]
=
3 "t
2 "t
!
!
!
#[I – ]
#[SO 2–
4 ]
=2
#t
#t
10. Given the initial rate data for the decomposition reaction,
A → 2B
determine the rate expression for the reaction.
[A], M
0.125
0.175
0.250
a.
b.
c.
d.
e.
–Δ[A]/Δt M/s
5.14 × 102
1.01 × 103
2.06 × 103
= 4.11 × 103[A]
= 4.11 × 103[A]2
= 3.29 × 104[A]2
= 5.14 × 102
= 3.29 × 104[A]
White Exam – Page 6
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11. For the first–order decomposition of N2O5 at 340 K, where k = 5.8 × 10–3 s–1, calculate the concentration
after 455 seconds if the initial concentration is 0.380 M.
a.
b.
c.
d.
e.
0.0071 M
0.0087 M
0.011 M
0.019 M
0.027 M
12. Which of the following changes generally lead to greater reaction rates?
1. Increasing the temperature
2. Decreasing the concentration of a reactant
3. Adding a catalyst
a.
b.
c.
d.
e.
1 only
2 only
1 and 3
2 and 3
1,2, and 3
13. A chemical reaction that is endothermic by 20 kJ has an activation energy of 180 kJ. What is activation
energy for the reverse reaction?
a.
b.
c.
d.
e.
200 kJ
180 kJ
160 kJ
20 kJ
–160 kJ
White Exam – Page 7
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14. An “old chemist’s tale” says that near room temperature (300 K) a 10°C increase in temperature
typically doubles the reaction rate. What is the activation energy for a reaction whose rate at room
temperature doubles upon an 10°C increase in temperature?
a.
b.
c.
d.
e.
504 kJ/mol
155 kJ/mol
72.5 kJ/mol
53.3 kJ/mol
0.50 kJ/mol
15. A proposed mechanism for the reaction of CH3OH with HBr
CH3OH + H+ + Br– → CH3Br + H2O
involves two steps,
CH3OH + H+
CH3OH2+
+
–
CH3OH2 + Br → CH3Br + H2O
(fast, equilibrium)
(slow)
What is a rate law that is consistent with this mechanism?
a.
b.
c.
d.
e.
rate = k[CH3OH][H+][Br–]
rate = k[CH3OH2+][ Br–]
rate = k[CH3OH2+][ H+]
rate = k[CH3OH][H+]
rate = k[CH3OH][Br–]
16. Which of the following statements concerning equilibrium constants are true?
1. Kinetically fast reactions always have large equilibrium constants.
2. Temperature has no effect on an equilibrium constant.
3. Reactant favored reactions have negative equilibrium constants.
a.
b.
c.
d.
e.
1 only
2 only
3 only
1 and 2
none of above
White Exam – Page 8
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17. Choose the correct expression for the equilibrium constant Ka for aqueous HF.
a.
b.
Ka =
[F – ][H 3O + ]
[HF]
Ka =
[HF]
[F ][H 3O + ]
Ka =
[F – ][H 3O + ]
[HF][H 2O]
Ka =
[HF][H 2O]
[F – ][H 3O + ]
!
c.
!
d.
!
e.
!
–
[F – ]
Ka =
[HF]
18. Write a balanced chemical reaction that corresponds to the following equilibrium constant expression.
!
a.
b.
c.
d.
e.
3 O2(g)
2 SO3(g)
2 SO3(g)
3 O2(g)
2 SO2(g) + O2(s)
2 SO3(g)
2 S(s) + 3 O2(g)
2 SO3(g)
2 SO3(g)
2 S(s) + 3 O2(g)
19. The equilibrium constant at 25ºC for the dissolution of silver iodide is 8.3 × 10–17. If an excess of AgI(s)
is added to water and allowed to equilibrate, what is the equilibrium concentration of Ag+?
a.
b.
c.
d.
e.
8.3 × 10–17 M
4.2 × 10–9 M
8.3 × 10–9 M
9.1 × 10–9 M
1.4 × 10–5 M
White Exam – Page 9
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20. The standard free energy change for the reaction described in the previous problem is
a.
b.
c.
d.
e.
91.7 kJ
39.8 kJ
0 kJ
–39.8 kJ
–91.7 kJ
21. An excess of PbBr2(s) is placed in 750 mL of water. At equilibrium, the solution contains 0.012 M Pb2+.
What is the equilibrium constant for the reaction?
a.
b.
c.
d.
e.
1.7 × 10–6
3.5 × 10–6
6.9 × 10– 6
1.4 × 10–4
5.8 × 10–4
22. The enthalpy of reaction is +26.9 kJ for this reaction
2 ICl(g)
I2(g) + Cl2(g)
At 25°C, Kp = 5.0 × 10–6. If the temperature is increased to 50°C, what change (if any) will occur?
a.
b.
c.
d.
e.
Kp will increase.
Kp will decrease.
Kp will remain unchanged.
[ICl] will increase.
[I2] will decrease.
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23. Addition of which of the following reagents will increase the solubility of Cu(OH)2 in water?
a.
b.
c.
d.
e.
HCl
NaCl
NaOH
Cu(OH)2
Cu(NO3)2
24. What is the conjugate acid of HPO42–(aq)?
a.
b.
c.
d.
e.
H3PO4
H2PO4–
PO42–
H3O+
OH–
25. Molecules or ions that can behave as either a Brønsted–Lowry acid or base are called
a.
b.
c.
d.
e.
amphiprotic.
hydronium.
polyprotic acids or bases.
conjugate acids or bases.
weak electrolytes.
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26. What is the hydroxide ion concentration in 0.094 M HCl at 25ºC?
a.
b.
c.
d.
e.
10.6 M
0.094 M
1.26 × 10–5 M
9.77 × 10–10 M
1.06 × 10–13 M
27. At 50ºC, the water ionization constant, Kw, is 5.5 × 10–14. What is the H3O+ concentration in neutral
water at this temperature?
a.
b.
c.
d.
e.
5.5 × 10–14 M
5.4 × 10–8 M
1.0 × 10–7 M
2.3 × 10–7 M
5.5 × 10–7 M
28. The information above allows us to conclude that the water ionization equilibrium has
a.
b.
c.
d.
e.
ΔGº>0, ΔHº>0
ΔGº>0, ΔHº<0
ΔGº=0, ΔHº>0
ΔGº<0, ΔHº<0
ΔGº<0, ΔHº>0
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29. Which of the following weak acids has the strongest conjugate base?
a.
b.
c.
d.
e.
acetic acid, Ka = 1.8 × 10–5
benzoic acid, Ka = 6.3 × 10–5
dihydrogen phosphate ion, Ka = 6.2 × 10–8
formic acid, Ka = 1.8 × 10–4
hydrocyanic acid, Ka = 4.0 × 10–10
30. What is the pH of a 0.75 M solution of sodium cyanide, NaCN? (Ka for HCN =
4.00 × 10–10)
a.
b.
c.
d.
e.
2.36
4.33
9.58
10.04
11.64
31. If you mix equal molar quantities of CH3CO2H (Ka = 1.8 × 10–5) and NaF (Kb =
1.4 × 10–11), the resulting solution will be
a.
b.
c.
d.
e.
acidic because Ka of CH3CO2H is greater than Kb of F–.
acidic because Ka of CH3CO2H is greater than Ka of F–.
basic because Ka of CH3CO2H is greater than Kb of F–.
basic because Ka of CH3CO2H is greater than Ka of F–.
neutral.
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32. What is the effect of adding ammonium chloride to a solution of ammonia?
1. The pH increases.
2. The concentration of NH3 increases.
3. The concentration of H3O+ increases
a.
b.
c.
d.
e.
1 only
2 only
3 only
1 and 2
2 and 3
33. Only 1.06 g of Ca(NO3)2 will dissolve per liter of a solution that is buffered to a pH of 13.00. What is the
value of Ksp for Ca(OH)2? The molar mass of Ca(NO3)2 is 164.1 g/mol.
a.
b.
c.
d.
e.
1.8 × 10–10
4.2 × 10–7
1.1 × 10–6
4.2 × 10–5
6.5 × 10–5
34. The following anions can be separated by precipitation as silver salts: Cl–, Br–, I–, CrO42–. If Ag+ is added
to a solution containing the four anions, each at a concentration of 0.10 M, in what order would they
precipitate?
Compound
AgCl
Ag2CrO4
AgBr
AgI
a.
b.
c.
d.
e.
Ksp
1.8 × 10–10
9.0 × 10–12
3.3 × 10–13
1.5 × 10–16
AgCl→Ag2CrO4→AgBr→AgI
AgI→AgBr→Ag2CrO4→AgCl
Ag2CrO4→AgCl→ AgBr→AgI
Ag2CrO4→AgI →AgBr→AgCl
AgI→AgBr→AgCl→Ag2CrO4
White Exam – Page 14
Name_________________________
35. Thermodynamics can be used to determine all of the following EXCEPT
a.
b.
c.
d.
e.
the direction in which a reaction is spontaneous.
the extent to which a reaction occurs.
the rate of reaction.
the temperature at which a reaction is spontaneous.
the enthalpy change of a reaction.
36. A statement of the second law of thermodynamics is that
a.
b.
c.
d.
e.
spontaneous reactions are always exothermic.
energy is conserved in a chemical reaction.
the entropy of the universe is continually increasing.
the enthalpy of reaction is the difference between product and reactant enthalpies.
the Gibbs free energy is a function of both enthalpy and entropy.
37. All of the following processes lead to an increase in entropy EXCEPT
a.
b.
c.
d.
e.
increasing the temperature of a gas.
freezing a liquid.
evaporating a liquid.
forming mixtures from pure substances.
chemical reactions that increase the number of moles of gas.
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Name_________________________
38. Diluting concentrated sulfuric acid with water can be dangerous. The temperature of the solution can
increase rapidly. What are the signs of H, S, and G for this process?
a.
b.
c.
d.
e.
ΔH < 0, ΔS < 0, ΔG < 0
ΔH < 0, ΔS > 0, ΔG < 0
ΔH < 0, ΔG < 0 but cannot determine ΔS
ΔH > 0, ΔS > 0, ΔG < 0
ΔH > 0, ΔS < 0, ΔG > 0
39. Write a balanced chemical equation for the oxidation of Ag(s) by concentrated nitric acid. Two products
of the reaction are NO2(g) and Ag+(aq).
a.
b.
c.
d.
e.
HNO3(aq) + Ag(s) → NO2(g) + Ag+(aq)
HNO3(aq) + Ag(s) → NO2(g) + Ag+(aq) + OH–(aq)
HNO3(aq) + Ag(s) + H+(aq) → NO2(g) + Ag+(aq) + H2O(l)
2 HNO3(aq) + 2 Ag(s) → 2 NO2(g) + 2 Ag+(aq) + H2(g) + O2(g)
2 HNO3(aq) + Ag(s) + H2O(l) → NO2(g) + Ag+(aq) + H+(aq)
40. All of the following statements concerning voltaic cells are true EXCEPT
a.
b.
c.
d.
e.
the two half–cells are connected by a salt bridge.
electrons flow from the anode to the cathode.
oxidation occurs at the cathode.
voltaic cells can be used as a source of energy.
a voltaic cell consists of two–half cells.
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41. Consider the following half–reactions:
Fe3+(aq) + e– → Fe2+(aq)
Sn2+(aq) + 2 e– → Sn(s)
Fe2+(aq) + 2 e– → Fe(s)
Al3+(aq) + 3 e– → Al(s)
Mg2+(aq) + 2 e– → Mg(s)
Eº = +0.77 V
Eº = –0.14 V
Eº = –0.44 V
Eº = –1.66 V
Eº = –2.37 V
Which of the above metals or metal ions are able to oxidize Al(s)?
a.
b.
c.
d.
e.
Fe3+ and Sn2+
Fe3+, Sn2+, and Fe2+
Fe2+, Sn, and Fe
Mg and Mg2+
Mg2+ only
42. Calculate the cell potential, at 25°C, based upon the overall reaction
3 Fe3+(aq) + Al(s) → 3 Fe2+(aq) + Al3+(aq)
if [Fe3+] = 0.300 M, [Fe2+] = 0.150 M, and [Al3+] = 0.300 M. The standard reduction potentials are
Fe3+(aq) + e– → Fe2+(aq)
Al3+(aq) + 3 e– → Al(s)
a.
b.
c.
d.
e.
+0.87 V
+2.26 V
+2.46 V
+3.01 V
+4.21 V
Eº = +0.771 V
Eº = –1.66 V
White Exam – Page 17
Name_________________________
43. Given the following standard reduction potentials,
Pb2+(aq) + 2 e– → Pb(s)
PbSO4(s) + 2 e– → Pb(s) + SO 42–(aq)
Eº = –0.126 V
Eº = –0.356 V
determine the Ksp for PbSO4(s) at 25ºC.
a.
b.
c.
d.
e.
1.7 × 10–8
8.5 × 10–8
5.3 × 10–6
4.2 × 10–4
2.1 × 10–4
44. Calculate the standard reduction potential for the following reaction at 25ºC,
Ag(NH3)2+(aq) + e– → Ag(s) + 2 NH3(aq)
given the following thermodynamic information.
Ag+(aq) + e– → Ag(s)
Ag+(aq) + 2 NH3(aq) → Ag(NH 3)2+(aq)
a.
b.
c.
d.
e.
–1.2 V
–0.54 V
+0.37 V
+0.54 V
+1.2 V
Eº = 0.80 V
K f = 1.7 × 107
Potentially Useful Equations
ln
!
[A]t
= "kt
[A]0
1
1
"
= kt
[A]t [A]0
t1/ 2 =
!
ln2 0.693
=
k
k
k(T) = Ae"E a / RT
ln k(T) = ln A "
Ea
RT
!
[C]c [D]d
K(T) =
[A]a [B]b
!
Kp = Kc(RT)(c+d–a–b)
!
Kw = [H3O+][OH–] = 1.0 × 10–14 (at 25˚C)
!
k2
E # 1 1&
=" a% " (
k1
R $ T2 T1 '
!
pH = –log[H3O+]
!
ln k 2 " ln k1 = ln
pOH = –log[OH–]
Ka =
[H 3O + ][A" ]
[HA]
Kb =
[BH + ][OH" ]
[B]
!
pK a = pH - log
pKw = 14.00 = pH + pOH
[A– ]
[HA]
pKa = –logKa
!
Ka Kb = Kw
ΔS = qrev/T
ΔSuniv = ΔSsys + ΔSsurr
ΔG° = ΔH° – TΔS°
ΔG = ΔG° + RT ln Q
o
o
o
E cell
= E cathode
" E anode
!
ΔG° = –nFEo
ΔG° = –RT ln K
E = Eo "
lnK =
!
!
RT
0.0257V
lnQ = E o "
lnQ (at 25°C)
nF
n
nF o
n
E =
E o (at 25°C)
RT
0.0257V