CHEM 1252
Practice Exam 4 (120 minutes)
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Substance
CH2ClCOOH
HF
HClO
HNO2
NH3
CH3NH2
Equilibrium Constant
Ka = 1.4 × 10−3
Ka = 7.1 × 10−4
Ka = 2.9 × 10−8
Ka = 4.5 × 10−4
Kb = 1.76 × 10−5
Kb = 4.4 × 10−4
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Multiple Choice Instructions: Bubble the letter that corresponds to the best answer for each
question on your Gradescope answer sheet. Any answers for questions 1–15 recorded on these
pages will not be graded. Each question is worth 3 points.
Assume a temperature of 25 °C throughout all questions on the exam unless otherwise
noted. Refer to the cover page of the exam for the relevant equilibrium constants.
1. Which of the following solutions will have the lowest freezing point?
(A) 0.010 m RbI
(D) 0.010 m K2SeO4
(B) 0.015 m SrBr2
(E) 0.040 m CH2O
(C) 0.035 m CH3OH
2. Aqueous solutions of 30.0% (by mass) H2O2 are used to oxidize metals or organic molecules in
chemical reactions. Calculate the molality of this solution.
(A) 12.6 m
(D) 0.974 m
(B) 6.78 m
(E) 30.0 m
(C) 9.79 m
3. For the following reaction, which answer choice is a conjugate acid-base pair?
HC2O4−(aq) + H2O(l) ⇌ H3O+(aq) + C2O42−(aq)
(A) HC2O4– and H2O
(D) HC2O4– and C2O42–
(B) HC2O4– and H3O+
(E) H3O+ and C2O42−
(C) H2O and C2O42–
4. What is the pH of an aqueous solution that is 0.0025 M LiOH.
(A) 2.60
(D) –11.40
(B) –2.60
(E) –2.25
(C) 11.40
5. At 60°C the equilibrium constant for the autoionization of water is 1.0 × 10−13. What is the pH of
water at this temperature? Is it acidic, basic or neutral?
(A) 7.00; Neutral
(D) 6.50; Acidic
(B) 13.00; Basic
(E) 6.50; Neutral
(C) 1.00; Acidic
6. Ammonia (NH3) has a base ionization constant (Kb) of 1.8 × 10–5. What is the initial concentration
of an aqueous ammonia solution with a pH of 11.68?
(A) 0.28 M
(D) 0.0090 M
(B) 3.6 M
(E) 0.0048
(C) 1.3 M
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7. Chlorous acid has a pKa of 1.96 and lactic acid has a pKa of 3.86. Which of the following
statements is TRUE?
(A) Chlorous acid and lactic acid are both strong acids.
(B) The Ka value for lactic acid is larger than the Ka value for chlorous acid.
(C) Chlorous acid is a stronger acid than lactic acid.
(D) In a 1.0 M aqueous solution of each of these acids, the lactic acid solution will have the highest
concentration of H3O+.
(E) None of the above are true.
8. How many liters of chlorine gas at 25 °C and 0.950 atm can be produced by the reaction of
500. mL of 1.104 M HCl reacting with excess MnO2?
MnO2(s) + 4HCl(aq) → MnCl2(aq) + 2H2O(l) + Cl2(g)
(A) 5.36 × 10−3 L
(D) 3.09 L
(B) 0.138 L
(E) 3.55 L
(C) 0.282 L
9. Which of the following statements is FALSE?
(A) Dissolving NH4I in water will produce an acidic solution.
(B) Dissolving KNO2 in water will produce a neutral solution.
(C) Dissolving AlBr3 in water will produce an acidic solution.
(D) Dissolving CaCl2 in water will produce a neutral solution.
(E) Dissolving CH3COONa in water will produce a basic solution.
10. Which of the following acid-base reactions favors the formation of the products (in other
words, Kc > 1)?
−
(A) HCN(aq) + NO−
2 (aq) ⇌ CN (aq) + HNO2 (aq)
+ (aq)
(B) NH4
+ F − (aq) ⇌ NH3 (aq) + HF(aq)
(C) HF(aq) + H2 O(l) ⇌ F − (aq) + H3 O+ (aq)
+
(D) H2 CO3 (aq) + NH3 (aq) ⇌ HCO−
3 (aq) + NH4 (aq)
(E) All of these reactions are product-favored (Kc > 1).
11. Which of the following statements about buffers is FALSE?
(A) A buffer lessens the impact of added strong acid or strong base on the solution’s pH.
(B) A buffer consists of any weak acid and any weak base.
(C) The higher a buffer’s capacity, the larger [H3O+] and [OH−] that can be added while resisting
large pH changes.
(D) The acid component’s Ka has a significant impact on the buffer solution’s pH.
(E) All of the above statements are true.
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12. What is the pH of a solution that consists of 0.50 M H2C6H6O6 (ascorbic acid) and 0.75 M
NaHC6H6O6 (sodium ascorbate)? For ascorbic acid, Ka = 6.8 × 10–5.
(A) 3.76
(D) 4.57
(B) 3.99
(E) 5.66
(C) 4.34
13. Which of the following combinations can be used to make a buffer? Assume equal volumes are
used.
(A) 0.20 M NH3 and 0.10 M NH4Cl
(B) 0.10 M NH4Cl and 0.10 M NaF
(C) 0.20 M NH3 and 0.20 M HCl
(D) 0.20 M NH3 and 0.10 M HF
(E) All of these combinations can be used to make a buffer.
14. A 0.200 L solution of 0.20 M HCl is titrated with 0.80 M NaOH. What volume of the NaOH
solution is needed to reach the equivalence point?
(A) 0.200 L
(D) 0.050 L
(B) 0.800 L
(E) 1.25 L
(C) 0.032 L
15. A 0.050 L solution of 0.64 M CH3COOH is titrated with 0.32 M NaOH. Which species remains in
solution with the highest concentration (i.e. the primary species) after 0.100 L of the NaOH solution
has been added?
(A) CH3COOH(aq)
(D) OH−(aq)
(B) H3O+(aq)
(E) CH3COOH(aq) and CH3COO−(aq)
(C) CH3COO−(aq)
16. Which of the following is the correct solubility product constant expression for magnesium
fluoride?
(A) 𝐾𝑠𝑝 = [Mg 2+ ][F − ]
(B) 𝐾𝑠𝑝 = [Mg 2+ ][F − ]2
(C) 𝐾𝑠𝑝 =
[Mg2+ ][F− ]2
[MgF2 ]
1
(D) 𝐾𝑠𝑝 = [Mg2+][F−]2
1
(E) 𝐾𝑠𝑝 = [Mg2+][F−]
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17. What is the molar solubility of PbF₂ in a solution that contains 0.0500 M F⁻ ions? The Ksp of
PbF₂ is 3.60×10⁻⁸.
(A) 2.1 × 10⁻3 M
(D) 9.0 × 10⁻11 M
(B) 1.4 × 10⁻5 M
(E) 7.2 × 10⁻7 M
(C) 3.6 × 10⁻6 M
Free Response Instructions: Answer the following problems in the box provided. Be sure to show
all your work, including units, round your answers to the appropriate significant figures, and
box your final answer. Check that the assumption is valid if you use the 5% rule. Provide the
quadratic equation if appropriate, even if you solved it using your calculator.
18. The beakers below contain an aqueous solution of three different acids, HX, HY, and HZ,
respectively, and represent the solutions at equilibrium. The initial concentrations of all three
solutions were the same. Assume the volume of each solution is the same.
HX =
HY =
HZ =
H3O+ =
H3O+ =
H3O+ =
Answer the following questions. Briefly explain your answers.
(A) The conjugate base particles are left out of each figure. Draw the appropriate number of
conjugate base particles in each beaker above, matching the legend provided for the acid in each
beaker.
Explanation
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(B) Which acid is the strongest acid? Briefly explain.
Circle one
Explanation
HX
HY
HZ
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(C) What acid will have the lowest Ka value? Briefly explain.
Circle one
Explanation
HX
HY
HZ
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(D) What conjugate base will be the weakest base? __________ Briefly explain.
Explanation
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(E) Each particle represents 0.030 M. Calculate the hydroxide ion concentration in the beaker that
contains the acid HY (the middle beaker).
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19. The molar solubility of nickel(II) hydroxide is 5.2 x 10–6 M. What is the solubility product
constant (Ksp) for nickel(II) hydroxide?
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20. For each of the following species, indicate whether the resulting solution would be acidic (A),
basic (B), or neutral (N), if 1.00 mole of each were dissolved in 1.0 liter of water. Provide the
predominant acid-base reaction (including physical states) that occurs to explain your choice.
Substance
Acidic/Basic/Neutral and Predominant Reaction that occurs
(A) sodium fluoride
(B) KClO4
(C) NH4NO3
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21. A 0.0450 L sample of 0.350 M weak acid (Ka = 1.8×10–4) is titrated with 0.200 M KOH. What is
the pH of the solution after the following volumes of KOH have been added?
(A) 0.0450 L KOH
(B) 0.100 L KOH
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CHEM 1252 Exam Information Sheet
Fall 2023
Useful Constants
Avogadro’s Number (N)
Molar gas constant
Faraday constant
6.022 x 1023 /mol
R = 8.314 J·mol−1·K−1
R = 0.08206 L·atm·mol−1·K−1
F = 96,485 C·mol−1
Conversion Factors and definitions
1 atm=
T [K] =
1 cal =
1 V=
1 amp=
1J=
760 mmHg
760 Torr
101,325 Pa
1.01325 bar
T [°C] + 273.15
4.184 J
1 J·C−1
1 C/s
1N·m
1 kg · m2
s2
1 N·m−2
1 kg·m−1·s−2
1N =
1 kg · m ·s−2
giga (G) =
109
mega (M) =
106
nano (n) =
10–9
pico (p) =
10–12
femto (f) =
10–15
Example: 400 nm to m:
1 Pa=
400 nm *
10−9 𝑚
=4.00x10-9
1 𝑛𝑚
m
Equations:
𝑃𝑖 = 𝑋𝑖 𝑃tot
0
𝑃𝑣𝑎𝑝,𝑠𝑜𝑙𝑛 = 𝑋𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑃𝑠𝑜𝑙𝑣𝑒𝑛𝑡
𝑚=
𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒
𝑘𝑔 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
𝛥𝑇b/f = 𝑖 ⋅ 𝐾b/f 𝑚
[𝐴]𝑡 = −𝑘𝑡 + [𝐴]0
𝑡½ =
[𝐴]0
2𝑘
𝑙𝑛[ 𝐴]𝑡 = −𝑘𝑡 + 𝑙𝑛[ 𝐴]0
𝑡½ =
𝑙𝑛 2
𝑘
1
1
= 𝑘𝑡 +
[𝐴]𝑡
[𝐴]0
𝑙𝑛 𝑘 = −
𝐸𝑎
+ 𝑙𝑛 𝐴
RT
𝑡½ =
𝑙𝑛
1
𝑘[𝐴]𝑜
𝑘1 𝐸𝑎 1
1
=
( − )
𝑘2
𝑅 𝑇2 𝑇1
𝑎𝑥 2 + 𝑏𝑥 + 𝑐 = 0:
−𝑏 ± √𝑏 2 − 4𝑎𝑐
𝑥=
2𝑎
[𝐴− ]
𝑝 𝐻 = 𝑝 𝐾𝑎 + 𝑙𝑜𝑔
[HA]
𝛥𝐺 (𝑜) = −𝑛𝐹𝐸 (𝑜)
ΔG = ΔGo + RT ln Q
0
𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑐𝑒𝑙𝑙
−
0.05916𝑉
log 𝑄
𝑛
°
𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑐𝑒𝑙𝑙
−
𝑅𝑇
ln 𝑄
𝑛𝐹
𝛥𝐺° = −𝑅𝑇 𝑙𝑛 𝐾
𝐸° =
𝑅𝑇
𝑙𝑛𝐾
𝑛𝐹
𝑜
𝑜
Eo=𝐸𝑟𝑒𝑑,𝑐𝑎𝑡ℎ𝑜𝑑𝑒
− 𝐸𝑟𝑒𝑑,𝑎𝑛𝑜𝑑𝑒
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Properties of water:
Specific heat of H2O(s)
Heat of fusion
Specific heat of H2O(l)
Heat of vaporization
s = 2.09 J·g−1·K−1
ΔHfus = 6.01 kJ·mol−1
s = 4.184 J·g−1·K−1
ΔHvap = 40.67 kJ·mol−1
Specific heat of H2O(g)
Molal b.p. elevation
Molal f.p. depression
Ion product (at 25 °C)
s = 1.84 J·g−1·K−1
Kb = 0.51 °C·m−1
Kf = 1.86 °C·m−1
Kw ≈ 1.00  10−14
Acid
“HX”
H2SO4
H3O+
HSO4−
H3PO4
HNO2
HF
HCO2H (formic acid)
CH3CO2H (HC2H3O2)
H2CO3
H2S
H2PO4−
HCN
NH4+
HCO3−
HPO42−
H2O
CH3CH2OH (ethanol)
NH3
Conjugate Base
“X−”
HSO4−
H2O
SO42−
H2PO4−
NO2−
F−
HCO2− (formate)
CH3CO2− (C2H3O2−)
HCO3−
HS−
HPO42−
CN−
NH3
CO32−
PO43−
OH−
CH3CH2O− (ethoxide)
NH2− (amide)
← increasing base strength
increasing acid strength →
Relative Strengths of Conjugate Acid-Base Pairs in Water
Solubility Rules:
1. Most nitrate, chlorate, and acetate salts are soluble.
2. Most salts of alkali metals and ammonium cations are soluble.
3. Most chloride, bromide, and iodide salts are soluble.
Exceptions: Salts containing Ag+, Pb2+, and Hg22+ ions are insoluble.
4. Most sulfate salts are soluble.
Exceptions: Sulfates containing Ag+, Ca2+, Sr2+, Ba2+, Pb2+ and Hg22+ ions are insoluble.
5. Most hydroxide and sulfide salts are insoluble.
Exceptions: Salts containing Ca2+, Sr2+, Ba2+, alkali metals, and ammonium cations are soluble.
6. Most carbonate, chromate, and phosphate salts are insoluble.
Exceptions: Salts of alkali metal and ammonium cations are soluble.
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