Honors Chemistry
Take Home Quiz: Ch 16 & 17
Name __________________________
Period ____ Date ___________
1. 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 HF and 0.10 M NaC2H3O2
C) A solution that is 0.10 M HCl and 0.10 M NH 4+
D) A solution that is 0.10 M NaOH and 0.10 M KOH
E) None of the above are buffer systems.
2. 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 acid or base.
C) A buffer resists pH change by neutralizing added acids and bases.
D) A buffer does not change pH when strong acid or base is added.
E) None of the above are true.
3. 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]
E) It is not possible to make a
buffer of this pH from HCHO2
B) [HCHO2] = [NaCHO2]
and NaCHO2.
C) [HCHO2] << [NaCHO2]
D) [HCHO2] > [NaCHO2]
4. You wish to prepare an HC2H3O2 buffer with a pH of 4.24. If the pKa of is 4.74, what ratio of
C2H3O2⁻/HC2H3O2 must you use?
A) 0.10
D) 2.0
B) 0.50
E) 2.8
C) 0.32
5. A buffer solution is 0.100 M in both HC7H5O2 and LiC7H5O2 and has a pH of 4.19. Which of the
following pH values would you expect from the addition of a small amount of a dilute solution of a
strong base?
A) 3.89
D) 4.49
B) 3.69
E) There is not enough information
C) 5.69
to determine.
6. Calculate the pH of a buffer that is 0.105 M HC 2H3O2 and 0.146 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
7. 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
8. 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.
9. 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.
10. The highest pH for an effective buffer occurs when the base is how many times as concentrated
as the acid?
A) 2
D) 20
B) 5
E) 15
C) 10
11. Which of the following acids (listed with pKa values) and their conjugate base would form a buffer
with a pH of 8.10?
A) HC7H5O2, pKa = 4.19
D) HCN, pKa = 9.31
B) HF, pKa = 3.46
E) HClO2, pKa = 1.96
C) HClO, pKa = 7.54
12. A 1.0 L buffer solution is 0.250 M HC2H3O2 and 0.050 M LiC2H3O2. Which of the following
actions will destroy the buffer?
D) adding 0.050 moles of NaOH
A) adding 0.050 moles of LiC2H3O2
E) None of the above will destroy the
adding
0.050
moles
of
HC
H
O
B)
2 3 2
buffer.
C) adding 0.050 moles of HCl
13. 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.
14. 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
D) 11.24
B) 13.03
E) 7.00
C) 2.76
15. 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
D) 3.46
B) 3.15
E) 4.15
C) 4.33
16. 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
17. Which of the following compounds will have the highest molar solubility in pure water?
A) PbSO4, Ksp = 1.82 × 10-8
D) PbS, Ksp = 9.04 × 10-29
B) MgCO3, Ksp = 6.82 × 10-6
C) AgI, Ksp = 8.51 × 10-17
E) FeS, Ksp = 3.72 × 10-19
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18. Which of the following compounds will be more soluble in acidic solution than in pure water?
D) CuI
A) PbCl2
E) None of the above will be more
B) FeS
soluble in acidic solution.
Ca(ClO
)
C)
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19. Identify the salts that are in hard water.
A) CaCO3 and MgCO3
B) MgSO4 and CaSO4
C) NaCl and KBr
D) NaHSO4 and Na2SO4
E) NaOCl and HOCl
20. Determine the molar solubility of CuCl in a solution containing 0.050 M KCl.
Ksp (CuCl) = 1.0 × 10-6.
A) 1.0 × 10-12 M
D) 1.0 × 10-3 M
-7
E) 0.050 M
B) 5.0 × 10 M
C) 2.0 × 10-5 M
21. Give the expression for the solubility product constant for PbCl2.
[Pb 2  ][Cl  ]2
[PbCl 2 ]
[PbCl 2 ]
B)
[Pb 2  ][Cl  ]2
A)
D)
[Pb 2  ]2 [Cl  ]
[PbCl 2 ]
E) [Pb2+]2[Cl⁻]
C) [Pb2+][Cl⁻]2
22. Determine the molar solubility of PbSO4 in pure water. Ksp (PbSO4) = 1.82 × 10-8.
A) 1.82 × 10-8 M
D) 3.31 × 10-16 M
B) 1.35 × 10-4 M
E) 4.48 × 10-4 M
-9
C) 9.1 × 10 M
23. Which of the following compounds will have the highest molar solubility in pure water?
A) PbS, Ksp = 9.04 × 10-29
D) ZnS, Ksp = 1.6 × 10-24
B) CuS, Ksp = 1.27 × 10-36
C) Al(OH)3, Ksp = 3 × 10-34
E) Ag2S, Ksp = 8 × 10-48
24. The molar solubility of ZnS is 1.6 × 10-12 M in pure water. Calculate the Ksp for ZnS.
A) 8.0 × 10-13
D) 2.6 × 10-24
B) 3.2 × 10-12
E) 6.80 × 10-5
C) 1.6 × 10-35
25. 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
26. Give the equation for a saturated solution in comparing Q with Ksp.
A) Q > Ksp
D) Q ≠ Ksp
B) Q < Ksp
E) none of the above
C) Q = Ksp
3
27. 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.
28. 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
D) 4.8 × 10-5 M, CuI
-10
E) No precipitate will form at any
B) 3.0 × 10
M, CuCl
concentration of copper (I).
C) 4.8 × 10-5 M, CuCl
29. Identify the compound that is acid-insoluble.
A) PbCl2
B) As2S
C) CoS
D) Ca3(PO4)2
E) NaCl
30. Describe the solubility of Al(OH)3 with respect to pH.
A) soluble at low pH, insoluble in pH-neutral solution, and soluble at high pH
B) soluble at low pH, insoluble in pH-neutral solution, and insoluble at high pH
C) insoluble at low pH, insoluble in pH-neutral solution, and soluble at high pH
D) soluble at low pH, in pH-neutral solution, and at high pH
E) insoluble at low pH, in pH-neutral solution, and at high pH
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