Chemistry 104 –1
Professor F. Fleming Crim
Spring, 2006
Name______________________________
TA________________________________
Section_____________________________
Exam 3
1. This exam contains 12 pages of questions and instructions, a page of equations, a periodic table, and
a set of constants and conversion factors.
2. There are 32 multiple choice questions. Circle the one best answer for each.
3. You have 90 minutes to work on the exam.
4. Put your name on each page.
White Exam – Page 2
Name_________________________
1. What is the effect of adding 10 mL of 0.1 M NaOH(aq) to 100 mL of 0.2 M NH4+(aq)?
1. The pH will decrease.
2. The concentration of NH3 will increase.
3. The concentration of NH4+ will decrease.
a.
b.
c.
d.
e.
1 only
2 only
3 only
2 and 3
1, 2, and 3
2. What is the pH of an aqueous solution of 0.30 M HF and 0.15 M F–? (Ka of HF = 7.2 × 10–4)
a.
b.
c.
d.
e.
1.83
2.84
3.14
3.44
10.86
3. Which of the following combinations would be the best to buffer an aqueous solution at a pH of
5.0?
a.
b.
c.
d.
e.
H3PO4 and H2PO4–, Ka1 = 7.5 × 10–3
HNO2 and NO2–, Ka = 4.5 × 10–4
CH3CO2H and CH3COO–, Ka = 1.8 × 10–5
H2PO4– and HPO42–, Ka2 = 6.2 × 10–8
NH4+ and NH3, Ka = 5.7 × 10–10
White Exam – Page 3
Name_________________________
4. All of the following statements concerning buffers are true EXCEPT
a. buffers resist pH changes upon addition of small quantities of strong acids or
bases.
b. a buffer solution is always basic.
c. the pH of a buffer is close to the pKa of the weak acid from which it is made.
d. buffers contain appreciable quantities of a weak acid and its conjugate base.
e. buffers resist changes in pH when diluted with water.
5. If the ratio of acid to base in a buffer increases by a factor of 100, the pH of the buffer
a.
b.
c.
d.
e.
decreases by 1.
decreases by 2.
increases by 2.
increases by 1.
remains unchanged.
6. Which of the following equations is the solubility product for magnesium iodate, Mg(IO3)2?
a.
b.
c.
d.
e.
Ksp = [Mg2+][I–]2[O2–]6
Ksp = [Mg2+][I–]2[3O2–]2
Ksp = [Mg2+][IO3–]
Ksp = [Mg2+]2[IO3–]
Ksp = [Mg2+][IO3–]2
White Exam – Page 4
Name_________________________
7. To make a buffer using acetic acid one would add
a.
b.
c.
d.
e.
carbonic acid.
sodium acetate.
sodium chloride.
ammonium chloride.
ammonium phosphate.
8. A buffer solution is one that
a. contains more than the expected amount of solute for a particular temperature and
is therefore unstable.
b. contains the maximum amount of solute possible for a particular temperature.
c. changes color upon addition of strong base.
d. contains an equal number of hydronium and hydroxide ions.
e. resists changes in pH upon addition of acid or base.
9. The Ksp of CaSO4 is 4.9 × 10–5 at 25 °C. What is the concentration of Ca2+ in a saturated solution
of CaSO4?
a.
b.
c.
d.
e.
2.4 × 10–9 M
2.5 × 10–5 M
4.9 × 10–5 M
7.0 × 10–3 M
3.7 × 10–2 M
White Exam – Page 5
Name_________________________
10. What is the molar solubility of Ag2CrO4 in 0.0028 M K2CrO4 at 25 °C? The value of Ksp for
Ag2CrO4 is 5.4 × 10–12 at 25 °C.
a.
b.
c.
d.
e.
1.6 × 10–5 mol/L
2.2 × 10–5 mol/L
3.3 × 10–5 mol/L
6.6 × 10–5 mol/L
1.1 × 10–4 mol/L
11. What is the molar solubility of Fe(OH)3(s) in a solution that is buffered at pH 2.75 at 25 °C? The
Ksp of Fe(OH)3 is 6.3 × 10–38 at 25 °C.
a.
b.
c.
d.
e.
1.1 × 10–29 mol/L
1.1 × 10–26 mol/L
2.0 × 10–15 mol/L
2.2 × 10–10 mol/L
3.5 × 10–4 mol/L
12. Using the data from the question above, calculate the standard free energy change for the
dissolution of Fe(OH)3(s).
a.
b.
c.
d.
e.
–92.1 kJ/mol
–17.8 kJ/mol
17.8 kJ/mol
92.1 kJ/mol
212 kJ/mol
White Exam – Page 6
Name_________________________
13. 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
will they precipitate?
Compound
AgCl
Ag 2 CrO 4
AgBr
AgI
a.
b.
c.
d.
e.
K sp
1.8 × 10 –10
5.4 × 10–12
5.4 × 10 –13
8.5 × 10 –17
AgCl → Ag2CrO4 → AgBr → AgI
AgI → AgBr → Ag2CrO4 → AgCl
Ag2CrO4 → AgCl → AgBr → AgI
Ag2CrO4 → AgI → AgBr → AgCl
AgI → AgBr → AgCl → Ag2CrO4
14. Consider the reaction
Zn(OH)2(s) + 2 OH–(aq) ∫ Zn(OH)42–(aq) K = 8.7 × 10–2
If Ksp for Zn(OH)2 is 3.0 × 10–17, what is the value of the formation constant, Kform, for the
reaction below?
Zn2+(aq) + 4 OH–(aq) ∫ Zn(OH)42–(aq)
a.
b.
c.
d.
e.
2.6 × 10–18
3.4 × 10–16
2.9 × 1015
3.3 × 1016
3.8 × 1017
15. Which of the following occurs in a solution where the initial concentrations of Pb(NO3)2(aq) and
KI(aq) are both 0.004 M? The value of Ksp for PbI2 is 1.4×10–8.
a.
b.
c.
d.
e.
NO2 evolves from the solution.
PbI2(s) precipitates.
KNO3(s) precipitates.
No net reaction occurs.
Ksp changes to 9×10–9.
White Exam – Page 7
Name_________________________
16. Addition of NH3 increases the solubility of silver chloride in water because
a.
b.
c.
d.
e.
the solubility of many salts is affected by the pH of the solution.
the formation of complex ions displaces the solubility equilibrium to the right.
the solubility of most salts increases as temperature increases.
some insoluble compounds are amphoteric.
a common ion displaces the solubility equilibrium toward the undissolved solute.
17. Which group contains only solutes that would decrease the solubility of barium sulfate?
a.
b.
c.
d.
e.
HNO3, H2SO4, CH3COOH
SO2, CO2, NH3
Ba(NO3)2, Na2SO4, NaHSO4
Ba(OH)2, NaOH, NH4Cl
Na2SO4, NaOH, NaCl
18. 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 Gibbs free energy is a function of both enthalpy and entropy.
ΔS = – ΔH for any chemical reaction.
in a spontaneous process, the entropy of the universe increases.
White Exam – Page 8
Name_________________________
19. Which of the following linear chain alcohols is likely to have the highest standard entropy in the
liquid state?
a.
b.
c.
d.
e.
CH3OH
CH3CH2OH
CH3CH2CH2OH
CH3CH2CH2CH2OH
CH3CH2CH2CH2CH2OH
20. If a chemical reaction has a positive change in entropy then
a.
b.
c.
d.
e.
the change in enthalpy is negative.
the reaction is exothermic.
heat goes from the system into the surroundings.
the disorder of the system increases.
the reaction is spontaneous.
21. All of the following processes lead to an increase in entropy EXCEPT
a.
b.
c.
d.
e.
decreasing the temperature of a gas.
melting a solid
chemical reactions that increase the number of moles of gas.
forming mixtures from pure substances.
evaporating a liquid
White Exam – Page 9
Name_________________________
22. Predict the signs of ΔH and ΔS for the condensation of steam at 85 °C.
a.
b.
c.
d.
e.
ΔH < 0 and ΔS < 0
ΔH < 0 and ΔS > 0
ΔH > 0 and ΔS < 0
ΔH > 0 and ΔS > 0
ΔH = 0 and ΔS < 0
23. The dissolution of ammonium nitrate occurs spontaneously in water, as we saw using cold packs
in class. What are the signs of ΔH, ΔS, and ΔG for this process?
a.
b.
c.
d.
e.
ΔH < 0, ΔS < 0, ΔG < 0
ΔH < 0, ΔS > 0, ΔG < 0
ΔH < 0, ΔS > 0, ΔG > 0
ΔH > 0, ΔS > 0, ΔG < 0
ΔH > 0, ΔS < 0, ΔG > 0
24. At what temperatures will a reaction be spontaneous if ΔH = –76.0 kJ and ΔS = +231 J/K?
a.
b.
c.
d.
e.
All temperatures below 329 K
Temperatures between 0 K and 231 K
All temperatures above 329 K
The reaction will be spontaneous at any temperature.
The reaction will never be spontaneous.
White Exam – Page 10
Name_________________________
25. For a reaction, ΔH = +265 kJ and ΔS = +271.3 J/K. At what temperature will the system be in
equilibrium?
a.
b.
c.
d.
e.
6.30 K
102 K
359 K
719 K
977 K
26. The figure shows a plot of the free
energy versus course of reaction. Which
of the following statements are true?
1. The system is at equilibrium at point
A.
2. The double arrow marked by B gives
the free energy difference between pure
reactants and pure products.
3. The reaction is exothermic.
a.
b.
c.
d.
e.
1 only
2 only
3 only
1 and 2
2 and 3
27. Estimate the boiling point of carbon tetrachloride given the following thermodynamic
parameters. (Your calculation is only an estimate since you must assume that the thermodynamic
quantities are independent of temperature.)
º
ΔHf (kJ/mol)
S º (J/K·mol)
a.
b.
c.
d.
e.
–272 °C
25 °C
67 °C
69 °C
109 °C
CCl 4 (l)
–128.4
214.4
CCl 4 (g)
–96.0
309.7
White Exam – Page 11
Name_________________________
28. Five coins are tossed. Which combination of heads (H) and tails (T) is least likely?
a.
b.
c.
d.
e.
THTHT
TTTHT
HTTHT
HHTHT
THHTT
29. Which has the highest entropy at a given temperature?
a.
b.
c.
d.
e.
SO2(s)
S(s)
O2(g)
SO2(l)
SO2(g)
30. A reaction is exothermic and has a negative value of ΔS°. The value of ΔG° for this reaction is
therefore:
a.
b.
c.
d.
e.
negative at all temperatures.
positive at all temperatures.
positive above 0°C and negative below 0°C.
positive above a certain temperature and negative below it.
negative above a certain temperature and positive below it.
White Exam – Page 12
Name_________________________
31. If a chemical reaction is at equilibrium, it must be true that
a.
b.
c.
d.
e.
ΔG° = 1.
ΔG > 1.
ΔG° < 1.
ΔG = 1.
ΔG = 0.
32. The plot on the right shows the standard entropy of a substance versus temperature starting with
a perfect crystal at T=0 K and going past the boiling point of the substance. Which of the
following statements are true?
1. The entropy at A is positive.
2. The sloping line between B and C shows
the change in entropy with temperature for
the crystal.
3. The plot shows two phase transitions.
a.
b.
c.
d.
e.
1 only
2 only
3 only
1 and 2
2 and 3
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
[A]a [B]b
!
(c+d–a–b)
Kp = Kc(RT)
!
Kw = [H3O+][OH–] = 1.0 × 10–14 (at 25˚C)
!
Ka =
[H 3O + ][A" ]
[HA]
Kb =
[BH + ][OH" ]
[B]
!
!
pOH = –log[OH–]
pK a = pH - log
[A– ]
[HA]
pKa = –logKa
!
k2
E # 1 1&
=" a% " (
k1
R $ T2 T1 '
K(T) =
pH = –log[H3O+]
!
ln k 2 " ln k1 = ln
Ka Kb = Kw
ΔS = qrev/T
ΔSuniv = ΔSsys + ΔSsurr
ΔG° = ΔH° – TΔS°
ΔG = ΔG° + RT ln Q
ΔG° = –RT ln K
pKw = 14.00 = pH + pOH