Chemistry: A Molecular Approach, 2e (Tro)
Chapter 16 Aqueous Ionic Equilibrium
Multiple Choice Questions
1) Animals will lick up ethylene glycol (antifreeze) due to its sweet taste. The antidote for ethylene
glycol poisioning is the administration of ________.
A) ethyl alcohol ( alcoholic drinks)
B) isopropyl alcohol (rubbing alcohol)
C) mineral oil (laxative)
D) vinegar
E) sodium bicarbonate (baking soda)
2) An important buffer in the blood is a mixture of ________.
A) sodium chloride and hydrochloric acid
B) hydrochloric acid and sodium hydroxide
C) carbonic acid and bicarbonate ion
D) acetic acid and bicarbonate ion
E) acetic acid and carbonate ion
3) Identify a good buffer.
A) small amounts of both a weak acid and its conjugate base
B) significant amounts of both a strong acid and a strong base
C) small amounts of both a strong acid and a strong base
D) significant amounts of both a weak acid and a strong acid
E) significant amounts of both a weak acid and its conjugate base
4) If the pKa of HCHO2 is 3.74 and the pH of an HCHO2/NaCHO2 solution is 3.11, which of the
following is TRUE?
A) [HCHO2] < [NaCHO2]
B) [HCHO2] = [NaCHO2]
C) [HCHO2] << [NaCHO2]
D) [HCHO2] > [NaCHO2]
E) It is not possible to make a buffer of this pH from HCHO2 and NaCHO2.
5) If the pKa of HCHO2 is 3.74 and the pH of an HCHO2/NaCHO2 solution is 3.89, which of the
following is TRUE?
A) [HCHO2] < [NaCHO2]
B) [HCHO2] = [NaCHO2]
C) [HCHO2] > [NaCHO2]
D) [HCHO2] >> [NaCHO2]
E) It is not possible to make a buffer of this pH from HCHO2 and NaCHO2.
6) If the pKa of HCHO2 is 3.74 and the pH of an HCHO2/NaCHO2 solution is 3.74, which of the
following is TRUE?
A) [HCHO2] > [NaCHO2]
B) [HCHO2] = [NaCHO2]
C) [HCHO2] < [NaCHO2]
D) [HCHO2] < <[NaCHO2]
E) It is not possible to make a buffer of this pH from HCHO2 and NaCHO2.
7) Calculate the pH of a buffer that is 0.225 M HC2H3O2 and 0.162 M KC2H3O2. The Ka for
HC2H3O2 is 1.8 × 10-5.
A) 4.89
B) 9.11
C) 4.74
D) 9.26
E) 4.60
8) Calculate the pH of a solution formed by mixing 250.0 mL of 0.15 M NH4Cl with 100.0 mL of 0.20
M NH3. The Kb for NH3 is 1.8 x 10-5.
A) 9.13
B) 9.25
C) 9.53
D) 4.74
E) 8.98
9) Calculate the pH of a solution formed by mixing 250.0 mL of 0.15 M NH4Cl with 200.0 mL of 0.12
M NH3. The Kb for NH3 is 1.8 × 10-5.
A) 9.06
B) 9.45
C) 4.55
D) 4.74
E) 9.26
10) Calculate the pH of a solution formed by mixing 100.0 mL of 0.20 M HClO with 200.0 mL of 0.30
M KClO. The Ka for HClO is 2.9 × 10-8.
A) 5.99
B) 8.01
C) 7.54
D) 7.06
E) 6.46
11) Calculate the pH of a solution formed by mixing 200.0 mL of 0.30 M HClO with 100.0 mL of 0.20
M KClO. The Ka for HClO is 2.9 × 10-8.
A) 5.99
B) 8.01
C) 7.54
D) 7.06
E) 6.46
12) Calculate the pH of a solution formed by mixing 200.0 mL of 0.30 M HClO with 300.0 mL of 0.20
M KClO. The Ka for HClO is 2.9 × 10-8.
A) 5.99
B) 8.01
C) 7.54
D) 7.06
E) 6.46
13) Calculate the pH of a solution formed by mixing 150.0 mL of 0.10 M HC7H5O2 with 100.0 mL of
0.30 M NaC7H5O2. The Ka for HC7H5O2 is 6.5 × 10-5.
A) 4.19
B) 9.69
C) 4.49
D) 4.31
E) 10.51
14) Calculate the pH of a solution formed by mixing 250.0 mL of 0.15 M HCHO2 with 100.0 mL of
0.20 M LiCHO2. The Ka for HCHO2 is 1.8 × 10-4.
A) 3.87
B) 3.74
C) 10.53
D) 3.47
E) 10.13
15) A 1.00 L buffer solution is 0.150 M in HC7H5O2 and 0.250 M in LiC7H5O2. Calculate the pH of
the solution after the addition of 100.0 mL of 1.00 M HCl. The Ka for HC7H5O2 is 6.5 × 10-5.
A) 4.19
B) 5.03
C) 4.41
D) 3.34
E) 3.97
16) A 1.00 L buffer solution is 0.250 M in HF and 0.250 M in NaF. Calculate the pH of the solution
after the addition of 100.0 mL of 1.00 M HCl. The Ka for HF is 3.5 × 10-4.
A) 3.09
B) 4.11
C) 3.82
D) 3.46
E) 2.78
17) A 1.00 L buffer solution is 0.250 M in HF and 0.250 M in LiF. Calculate the pH of the solution
after the addition of 0.150 moles of solid LiOH. Assume no volume change upon the addition of base.
The Ka for HF is 3.5 × 10-4.
A) 3.46
B) 4.06
C) 2.85
D) 3.63
E) 4.24
18) A 1.50 L buffer solution is 0.250 M in HF and 0.250 M in NaF. Calculate the pH of the solution
after the addition of 0.0500 moles of solid NaOH. Assume no volume change upon the addition of base.
The Ka for HF is 3.5 × 10-4.
A) 3.34
B) 3.46
C) 3.57
D) 3.63
E) 2.89
19) A 1.50 L buffer solution is 0.250 M in HF and 0.250 M in NaF. Calculate the pH of the solution
after the addition of 0.100 moles of solid NaOH. Assume no volume change upon the addition of base.
The Ka for HF is 3.5 × 10-4.
A) 3.22
B) 3.82
C) 3.69
D) 3.09
E) 4.46
20) Which of the following is TRUE?
A) An effective buffer has a [base]/[acid] ratio in the range of 10 – 100.
B) A buffer is most resistant to pH change when [acid] = [conjugate base]
C) An effective buffer has very small absolute concentrations of acid and conjugate base.
D) A buffer can not be destroyed by adding too much strong base. It can only be destroyed by adding
too much strong acid.
E) None of the above are true.
21) Define buffer capacity.
A) Buffer capacity is the amount of acid or base that can be added to a buffer without destroying its
effectiveness.
B) Buffer capacity is the amount of acid that can be added until all of the base is used up.
C) Buffer capacity is the amount of base that can be added until all of the acid is used up.
D) Buffer capacity is the amount of acid that can be added until all of the acid is used up.
E) Buffer capacity is the amount of base that can be added until all of the base is used up.
22) Which of the following is TRUE?
A) The equivalence point is where the amount of acid equals the amount of base during any acid-base
titration.
B) At the equivalence point, the pH is always 7.
C) An indicator is not pH sensitive.
D) A titration curve is a plot of pH vs. the [base]/[acid] ratio.
E) None of the above are true.
23) When titrating a strong monoprotic acid and KOH at 25°C, the
A) pH will be less than 7 at the equivalence point.
B) pH will be greater than 7 at the equivalence point.
C) titration will require more moles of base than acid to reach the equivalence point.
D) pH will be equal to 7 at the equivalence point.
E) titration will require more moles of acid than base to reach the equivalence point.
24) When titrating a weak monoprotic acid with NaOH at 25°C, the
A) pH will be less than 7 at the equivalence point.
B) pH will be equal to 7 at the equivalence point.
C) pH will be greater than 7 at the equivalence point.
D) titration will require more moles of base than acid to reach the equivalence point.
E) titration will require more moles of acid than base to reach the equivalence point.
25) When titrating a monoprotic strong acid with a weak base at 25°C, the
A) pH will be 7 at the equivalence point.
B) pH will be greater than 7 at the equivalence point.
C) titration will require more moles of the base than acid to reach the equivalence point.
D) titration will require more moles of acid than base to reach the equivalence point.
E) pH will be less than 7 at the equivalence point.
26) A 100.0 mL sample of 0.18 M HClO4 is titrated with 0.27 M LiOH. Determine the pH of the
solution before the addition of any LiOH.
A) 1.74
B) 1.05
C) 0.74
D) 0.57
E) 1.57
27) A 100.0 mL sample of 0.18 M HClO4 is titrated with 0.27 M LiOH. Determine the pH of the
solution after the addition of 30.0 mL of LiOH.
A) 0.86
B) 1.21
C) 2.00
D) 1.12
E) 2.86
28) A 100.0 mL sample of 0.18 M HClO4 is titrated with 0.27 M LiOH. Determine the pH of the
solution after the addition of 50.0 mL of LiOH.
A) 12.48
B) 0.68
C) 2.35
D) 1.52
E) 3.22
29) A 100.0 mL sample of 0.18 M HClO4 is titrated with 0.27 M LiOH. Determine the pH of the
solution after the addition of 66.67 mL of LiOH (this is the equivalence point).
A) 0.97
B) 13.03
C) 2.76
D) 11.24
E) 7.00
30) A 100.0 mL sample of 0.180 M HClO4 is titrated with 0.270 M LiOH. Determine the pH of the
solution after the addition of 75.0 mL of LiOH.
A) 12.1
B) 2.65
C) 11.35
D) 1.89
E) 13.06
31) A 100.0 mL sample of 0.18 M HClO4 is titrated with 0.27 M LiOH. Determine the pH of the
solution after the addition of 100.0 mL of LiOH.
A) 13.13
B) 12.65
C) 1.35
D) 0.87
E) 12.95
32) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. Determine the pH of the solution
before the addition of any KOH. The Ka of HF is 3.5 x 10-4.
A) 4.15
B) 0.70
C) 2.08
D) 3.46
E) 1.00
33) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. Determine the pH of the solution
after the addition of 100.0 mL of KOH. The Ka of HF is 3.5 × 10-4.
A) 2.08
B) 3.15
C) 4.33
D) 3.46
E) 4.15
34) A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution
before the addition of any HNO3. The Kb of NH3 is 1.8 × 10-5.
A) 4.74
B) 9.26
C) 11.13
D) 13.00
E) 12.55
35) A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution
after the addition of 50.0 mL of HNO3. The Kb of NH3 is 1.8 × 10-5.
A) 4.74
B) 7.78
C) 7.05
D) 9.26
E) 10.34
36) A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution
after the addition of 150.0 mL of HNO3. The Kb of NH3 is 1.8 × 10-5.
A) 1.70
B) 6.44
C) 7.56
D) 12.30
E) 2.30
37) A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution
after the addition of 200.0 mL of HNO3. The Kb of NH3 is 1.8 × 10-5.
A) 6.44
B) 1.48
C) 2.00
D) 12.52
E) 12.00
38) A 100.0 mL sample of 0.10 M Ca(OH)2 is titrated with 0.10 M HBr. Determine the pH of the
solution before the addition of any HBr.
A) 12.86
B) 13.00
C) 13.30
D) 0.70
E) 1.00
39) A 100.0 mL sample of 0.10 M Ca(OH)2 is titrated with 0.10 M HBr. Determine the pH of the
solution after the addition of 100.0 mL HBr.
A) 2.00
B) 12.00
C) 1.30
D) 12.70
E) 7.00
40) A 100.0 mL sample of 0.10 M Ca(OH)2 is titrated with 0.10 M HBr. Determine the pH of the
solution after the addition of 200.0 mL HBr.
A) 2.62
B) 2.00
C) 1.48
D) 12.52
E) 7.00
41) A 100.0 mL sample of 0.10 M Ca(OH)2 is titrated with 0.10 M HBr. Determine the pH of the
solution after the addition of 300.0 mL HBr.
A) 1.60
B) 1.30
C) 1.00
D) 12.40
E) 1.12
42) A 100.0 mL sample of 0.10 M Ca(OH)2 is titrated with 0.10 M HBr. Determine the pH of the
solution after the addition of 400.0 mL HBr.
A) 1.00
B) 1.40
C) 1.22
D) 1.30
E) 2.11
43) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. Determine the pH of the solution
after the addition of 300.0 mL of KOH. The Ka of HF is 3.5 × 10-4.
A) 12.40
B) 9.33
C) 5.06
D) 8.94
E) 12.00
44) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. Determine the pH of the solution
after the addition of 400.0 mL of KOH. The Ka of HF is 3.5 × 10-4.
A) 13.08
B) 12.60
C) 13.85
D) 12.30
E) 12.78
45) A 100.0 mL sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the pH of the solution
after the addition of 100.0 mL of HNO3. The Kb of NH3 is 1.8 × 10-5.
A) 6.58
B) 10.56
C) 8.72
D) 3.44
E) 5.28
46) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH. Determine the pH of the solution
after the addition of 200.0 mL of KOH. The Ka of HF is 3.5 × 10-4.
A) 9.62
B) 7.00
C) 3.46
D) 10.54
E) 8.14
47) Determine the molar solubility of AgBr in a solution containing 0.150 M NaBr. Ksp (AgBr) = 7.7 ×
10-13.
A) 8.8 × 10-7 M
B) 3.9 × 10-13 M
C) 5.8 × 10-5 M
D) 5.1 × 10-12 M
E) 0.150 M
48) Determine the molar solubility of AgBr in a solution containing 0.200 M NaBr. Ksp (AgBr) = 7.7 ×
10-13.
A) 8.8 × 10-7 M
B) 1.54 × 10-13 M
C) 5.8 × 10-5 M
D) 3.8 × 10-12 M
E) 0.200 M
49) Determine the molar solubility of PbSO4 in pure water. Ksp (PbSO4) = 1.82 × 10-8.
A) 1.82 × 10-8 M
B) 1.35 × 10-4 M
C) 9.1 × 10-9 M
D) 3.31 × 10-16 M
E) 4.48 × 10-4 M
50) Determine the molar solubility of MgCO3 in pure water. Ksp (MgCO3) = 6.82 × 10-6.
A) 6.82 × 10-6 M
B) 3.41 × 10-6 M
C) 4.65 × 10-3 M
D) 2.61 × 10-3 M
E) 3.25 × 10-4 M
51) Determine the molar solubility of AgI in pure water. Ksp (AgI) = 8.51 × 10-17.
A) 9.22 × 10-9 M
B) 4.26 × 10-17 M
C) 8.51 × 10-17 M
D) 2.77 × 10-6 M
E) 4.40 × 10-6 M
52) The molar solubility of ZnS is 1.6 × 10-12 M in pure water. Calculate the Ksp for ZnS.
A) 8.0 × 10-13
B) 3.2 × 10-12
C) 1.6 × 10-35
D) 2.6 × 10-24
E) 6.80 × 10-5
53) The molar solubility of CuI is 2.26 × 10-6 M in pure water. Calculate the Ksp for CuI.
A) 5.11 × 10-12
B) 4.52 × 10-6
C) 1.50 × 10-3
D) 4.62 × 10-17
E) 1.02 × 10-11
54) Determine the molar solubility of Fe(OH)2 in pure water. Ksp for Fe(OH)2= 4.87 × 10-17.
A) 2.44 × 10-17 M
B) 1.62 × 10-17 M
C) 4.03 × 10-9 M
D) 3.65 × 10-6 M
E) 2.30 × 10-6 M
55) Determine the molar solubility of BaF2 in pure water. Ksp for BaF2 = 2.45 × 10-5.
A) 1.83 × 10-2 M
B) 1.23 × 10-5 M
C) 2.90 × 10-2 M
D) 4.95 × 10-3 M
E) 6.13 × 10-6 M
56) Determine the molar solubility for Al(OH)3 in pure water. Ksp for Al(OH)3 = 1.3 x 10-33.
A) 3.6 x 10-12 M
B) 2.2 x 10-10 M
C) 4.8 x 10-35 M
D) 2.6 x 10-9 M
E) 6.0 x 10-19 M
57) Determine the molar solubility for Pb3(PO4)2 in pure water. Ksp for Pb3(PO4)2 is 1.0 x 10-54.
A) 4.1 x 10-28 M
B) 5.8 x 10-10 M
C) 1.1 x 10-11 M
D) 6.2 x 10-12 M
E) 1.0 x 10-54 M
58) Determine the molar solubility for Zn(OH)2 in pure water. Ksp for Zn(OH)2 is 3.00 x 10-17.
A) 3.0 x 10-17 M
B) 7.5 x 10-18 M
C) 3.11 x 10-6 M
D) 1.96 x 10-6 M
E) 5.5 x 10-19 M
59) The molar solubility of Ba3(PO4)2 is 8.89 x 10-9 M in pure water. Calculate the Ksp for
Ba3(PO4)2.
A) 4.94 x 10-49
B) 5.55 x 10-41
C) 5.33 x 10-37
D) 8.16 x 10-31
E) 6.00 x 10-39
60) The molar solubility of Ag2S is 1.26 × 10-16 M in pure water. Calculate the Ksp for Ag2S.
A) 1.59 × 10-32
B) 1.12 × 10-8
C) 6.81 × 10-63
D) 3.78 × 10-12
E) 8.00 × 10-48
61) Determine the molar solubility of BaF2 in a solution containing 0.0750 M LiF. Ksp (BaF2) = 1.7 ×
10-6.
A) 2.3 × 10-5 M
B) 8.5 × 10-7 M
C) 1.2 × 10-2 M
D) 0.0750 M
E) 3.0 × 10-4 M
15
62) Determine the molar solubility of Al(OH)3 in a solution containing 0.0500 M AlCl3. Ksp
(Al(OH)3) = 1.3 × 10-33.
A) 2.6 x 10-9M
B) 5.2 x 10-31 M
C) 2.6 x 10-32 M
D) 1.04 x 10-29 M
E) 6.5 x 10-35 M
63) Stalactites and stalagmites form as ________ precipitates out of the water evaporating in
underground caves.
A) hydrochloric acid
B) sodium hydroxide
C) sodium chloride
D) calcium carbonate
E) sodium bicarbonate
64) Give the equation for an unsaturated solution in comparing Q with Ksp.
A) Q > Ksp
B) Q < Ksp
C) Q = Ksp
D) Q ≠ Ksp
E) none of the above
65) Give the equation for a saturated solution in comparing Q with Ksp.
A) Q > Ksp
B) Q < Ksp
C) Q = Ksp
D) Q ≠ Ksp
E) none of the above
66) Give the equation for a supersaturated solution in comparing Q with Ksp.
A) Q > Ksp
B) Q < Ksp
C) Q = Ksp
D) Q ≠ Ksp
E) none of the above
67) A solution containing AgNO3 is mixed with a solution of NaCl to form a solution that is 0.10 M in
AgNO3 and 0.075 M in NaCl. What will happen once these solutions are mixed? Ksp (AgCl) = 1.77 ×
10-10.
A) Nothing will happen since the molar solubility of AgCl is higher than the solution concentrations.
B) Silver chloride will precipitate out of solution, leaving an unsaturated solution of AgCl.
C) Silver chloride will precipitate out of solution, leaving a saturated AgCl solution.
D) Nothing will happen since NaCl and AgNO3 are both soluble compounds.
E) There is not enough information to say anything about this solution.
68) A solution containing CaCl2 is mixed with a solution of Li2C2O4 to form a solution that is 2.1 × 10–
5 M in calcium ion and 4.75 × 10-5 M in oxalate ion. What will happen once these solutions are mixed?
Ksp (CaC2O4) = 2.3 × 10-9.
A) A precipitate will form since Q > Ksp for calcium oxalate.
B) Nothing will happen since both calcium chloride and lithium oxalate are soluble compounds.
C) Nothing will happen since calcium oxalate is extremely soluble.
D) Nothing will happen since Ksp > Q for all possible precipitants.
E) There is not enough information to determine.
69) A solution containing CaCl2 is mixed with a solution of Li2C2O4 to form a solution that is 3.5 × 10–
4 M in calcium ion and 2.33 × 10-4 M in oxalate ion. What will happen once these solutions are mixed?
Ksp (CaC2O4) = 2.3 × 10-9.
A) Nothing will happen since both calcium chloride and lithium oxalate are soluble compounds.
B) Nothing will happen Ksp > Q for all possible precipitants.
C) A precipitate will form as calcium oxalate is not soluble to any extent.
D) A precipitate will form since Q > Ksp for calcium oxalate.
E) There is not enough information to determine.
70) A solution contains 0.021 M Cl⁻ and 0.017 M I⁻. A solution containing copper (I) ions is added to
selectively precipitate one of the ions. At what concentration of copper (I) ion will a precipitate begin to
form? What is the identity of the precipitate? Ksp(CuCl) = 1.0 × 10-6, Ksp(CuI) = 5.1 × 10-12.
A) 3.0 × 10-10 M, CuI
B) 3.0 × 10-10 M, CuCl
C) 4.8 × 10-5 M, CuCl
D) 4.8 × 10-5 M, CuI
E) No precipitate will form at any concentration of copper (I).
17
71) A solution contains 0.036 M Cu2+ and 0.044 M Fe2+. A solution containing sulfide ions is added
to selectively precipitate one of the metal ions from solution. At what concentration of sulfide ion will a
precipitate begin to form? What is the identity of the precipitate? Ksp(CuS) = 1.3 × 10-36, Ksp(FeS) =
6.3 × 10-18.
A) 1.4 × 10-16 M, FeS
B) 3.6 × 10-35 M, CuS
C) 3.6 × 10-35 M, FeS
D) 1.4 × 10-16 M, CuS
E) No precipitate will form at any concentration of sulfide ion.
72) A solution contains 2.2 × 10-3 M in Cu2+ and 0.33 M in LiCN. If the Kf for Cu(CN)42- is 1.0 ×
1025, how much copper ion remains at equilibrium?
A) 3.8 × 10-24 M
B) 1.9 x 10-26 M
C) 6.7 × 10-28 M
D) 2.9 × 10-27 M
E) 4.6 × 10-25 M
73) A solution contains 3.8 × 10-2 M in Al3+ and 0.29 M in NaF. If the Kf for AlF63- is 7 × 1019, how
much aluminum ion remains at equilibrium?
A) 1.1 × 10-19 M
B) 3.1 × 10-22 M
C) 9.1 × 10-19 M
D) 1.9 × 10-21 M
E) 4.4 × 10-20 M
Algorithmic Questions
1) Which of the following solutions is a good buffer system?
A) A solution that is 0.10 M NaCl and 0.10 M HCl
B) A solution that is 0.10 M HCN and 0.10 M LiCN
C) A solution that is 0.10 M NaOH and 0.10 M HNO3
D) A solution that is 0.10 M HNO3 and 0.10 M NaNO3
E) A solution that is 0.10 M HCN and 0.10 M KI
2) Which of the following solutions is a good buffer system?
A) A solution that is 0.10 M HC2H3O2 and 0.10 M LiC2H3O2
B) A solution that is 0.10 M HBr and 0.10 M KC2H3O2
C) A solution that is 0.10 M HI and 0.10 M NH4+
D) A solution that is 0.10 M NaOH and 0.10 M KOH
E) None of the above are buffer systems.
3) Which one of the following statements is TRUE?
A) A buffer is an aqueous solution composed of two weak acids.
B) A buffer can absorb an unlimited amount of base.
C) A buffer resists pH change by neutralizing added acids and bases.
D) A buffer does not change pH when strong acid is added.
E) None of the above are true.
4) A solution is prepared by dissolving 0.23 mol of hydrazoic acid and 0.27 mol of sodium azide in
water sufficient to yield 1.00 L of solution.The addition of 0.05 mol of NaOH to this buffer solution
causes the pH to increase slightly. The pH does not increase drastically because the NaOH reacts with
the ________ present in the buffer solution. The Ka of hydrazoic acid is 1.9 × 10-5.
A) H2O
B) H3O+
C) azide
D) hydrazoic acid
E) This is a buffer solution: the pH does not change upon addition of acid or base.
5) A solution is prepared by dissolving 0.23 mol of chloroacetic acid and 0.27 mol of sodium
chloroacetate in water sufficient to yield 1.00 L of solution. The addition of 0.05 mol of HCl to this
buffer solution causes the pH to drop slightly. The pH does not decrease drastically because the HCl
reacts with the __________ present in the buffer solution. The Ka of chloroacetic acid is 1.36 × 10-3.
A) H2O
B) H3O+
C) chloroacetate ion
D) chloroacetic acid
E) This is a buffer solution: the pH does not change upon addition of acid or base.
6) Calculate the pH of a solution that is 0.295 M in sodium formate (NaHCO2) and 0.205 M in formic
acid (HCO2H). The Ka of formic acid is 1.77 × 10-4.
A) 3.910
B) 3.587
C) 13.84
D) 10.10
E) 4.963
7) Calculate the pH of a solution that is 0.210 M in nitrous acid (HNO2) and 0.290 M in potassium
nitrite (KNO2). The acid dissociation constant of nitrous acid is 4.50 × 10-4.
A) 3.487
B) 3.210
C) 13.86
D) 10.51
E) 4.562
8) Calculate the percent ionization of nitrous acid in a solution that is 0.222 M in nitrous acid (HMO2)
and 0.278 M in potassium nitrite (KNO2). The acid dissociation constant of nitrous acid is 4.50 × 10-4.
A) 55.6
B) 0.162
C) 15.5
D) 2.78 × 10-3
E) 3.448