Deck 16: Additional Aqueous Equilibria

Question 1

In how many different pairs can these substances be mixed to produce buffer solutions?

A) 2
B) 3
C) 4
D) 5
E) 6

Answer 1

3

Question 2

A buffer solution is 0.500 M in ascorbic acid and 0.500 M in sodium ascorbate. Its pH is 4.10. After addition of 10 mL of 1 M NaOH to 1.00 L of this buffer, the most likely value of the pH is

A) 4.08.
B) 4.10.
C) 4.12.
D) 5.95.
E) 10.15.

Answer 2

4.12.

Question 3

An alkaline buffer would best be made using

A) a weak acid with a pK_a > 7.00 or K_a < 1.00 × 10^-7 and its conjugate base
B) a weak acid with a pK_a < 7.00 or K_a > 1.00 × 10^-7 and its conjugate base
C) a strong acid and a strong base
D) a weak acid with a K_a > 1.00 × 10¹⁴
E) a strong base and its conjugate acid

Answer 3

a weak acid with a pK_a > 7.00 or K_a < 1.00 × 10^-7 and its conjugate base

Question 4

In a buffer solution, if [A^-] = [HA] then

A) pH < pK_a
B) pH = pK_a
C) pH > pK_a
D) pH < 7.00
E) pH > 7.00

Question 5

A buffer solution is 0.500 M in acetic acid and 0.500 M in sodium acetate. Its pH is 4.74. What is its pH after dilution by a factor of 2?

A) 4.44
B) 4.49
C) 4.74
D) 4.99
E) 5.04

Question 6

A buffer solution

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

Question 7

Consider a buffer solution made up of H₂PO₄^- and HPO₄^2- For H₂PO₄^-, K_a = 6.2 × 10^-8. What mole ratio of HPO₄^2- to H₂PO₄^- will give a pH of 7.35?

A) 0.14 to 1
B) 0.72 to 1
C) 1 to 1
D) 1.4 to 1
E) More information is needed to answer this question.

Question 8

When the concentration of the acid and conjugate base components of a buffer are present in exactly equimolar proportions

A) the buffer is very limited in capacity to absorb either acid or base
B) the buffer has no capacity to absorb any added acid
C) the buffer has no capacity to absorb any added base
D) the buffer has a useful range of pH 6.5 - pH 7.5
E) the pH of the buffer solution is equal to the pK_a of the acid component of the buffer

Question 9

In a buffer solution, if [A^-] < [HA] then

A) pH < pK_a
B) pH = pK_a
C) pH > pK_a
D) pH < 7.00
E) pH > 7.00

Question 10

Which statement about buffer solutions is not true?

A) they resist changes in pH upon addition of small quantities of acid or base
B) they consist of approximately equimolar quantities of a weak acid and its conjugate base
C) they consist of approximately equimolar quantities of a weak base and its conjugate acid
D) their acid and base components must react with each other to form new products
E) their pH changes only slightly if the solution is diluted by a factor of 10 with pure water

Question 11

A buffer solution is 0.080 M in lactic acid (K_a = 1.8 × 10^-4) and 0.070 M in sodium lactate. The pH of the solution is

A) 2.86.
B) 3.68.
C) 3.80.
D) 4.18.
E) 4.62.

Question 12

In a buffer solution, if [A^-] > [HA], which of the following must be true?

A) pH < pK_a
B) pH > pK_a
C) pH = pK_a
D) pH < 7.00
E) pH > 7.00

Question 13

A buffer solution will result when ____ is added to a solution of K₃PO₄.

A) NH₃
B) K₂HPO₄
C) NaOH
D) Na₃PO₄
E) KCl

Question 14

Which formula represents the Henderson-Hasselbalch equation for the generic acid HA?

A)
B)
C)
D)
E)

Question 15

To make a buffer using acetic acid one could add

A) carbonic acid
B) sodium acetate
C) sodium chloride
D) ammonium chloride
E) ammonium phosphate

Question 16

If a buffer is made up using 1.00 mole of a weak acid (pK_a = 5.00) and 0.90 mole of its conjugate base, which of the following must be true?

A) pH < 5.00
B) pH = 5.00
C) pH > 5.00
D) pH = 7.00
E) pH > 7.00

Question 17

Which combination of solutions is the best choice for making a buffer solution?

A) equal volumes of 1 M acetic acid and 0.005 M sodium acetate
B) equal volumes of 0.5 M nitric acid and 0.5 M sodium hydroxide
C) equal volumes of 0.1 M formic acid and 0.1 M sodium formate
D) equal volumes of 0.1 M sulfuric acid and 0.001 M sodium sulfate
E) equal volumes of 0.05 M hydrochloric acid and 0.075 ammonium chloride

Question 18

In order to make up a buffer solution with a neutral pH, it would be best to make use of

A) a weak acid with a K_a at or near 1.00 × 10⁷ and its conjugate base
B) a weak acid with a pK_a near 7.00 or a K_a near 1.00 × 10^-7 and its conjugate base
C) an equimolar mixture of a strong acid and a strong base
D) an equimolar mixture of an extremely weak acid and an extremely strong base
E) an equimolar mixture of NaCl and HCl

Question 19

Which acid, in combination with its conjugate base, would be the best choice to make a buffer of pH = 4.20?

A) acetic acid (K_a = 1.8 × 10^-5)
B) benzoic acid (K_a = 6.3 × 10^-5)
C) formic acid (K_a = 1.8 × 10^-4)
D) hydrofluoric (K_a = 7.2 × 10^-4)
E) nitrous acid (K_a = 4.5 × 10^-4)

Question 20

What is the smallest amount of solid NaOH shown that will exceed the buffer capacity of 500 mL of a buffer that is 0.40 M in acetic acid and 0.15 M in sodium acetate. (K_a for acetic acid = 1.8 × 10^-5.)

A) 8.00 g
B) 5.16 g
C) 4.74 g
D) 4.31 g
E) 3.00 g

Question 21

Isotonic saline solution is 0.154 M NaCl(aq). What is the solubility of AgCl (K_sp = 1.8 × 10^-10) in such a solution?

A) 2.8 × 10^-11 M
B) 1.2 × 10^-9 M
C) 5.3 × 10^-6 M
D) 3.4 × 10^-5 M
E) 0.077 M

Question 22

What is the minimum volume of 0.55 M NaOH required to completely convert 1.00 L of 0.45 M acetic acid (pK_a = 4.74) to its conjugate base?

A) 0.82 L
B) 1.0 L
C) 1.2 L
D) 1.8 L
E) 2.2 L

Question 23

Exhibit 16-1 The following question(s) refer to the same titration: a 50.00 mL sample of 0.0950 M acetic acid (K_a = 1.8 × 10^-5) is being titrated with 0.106 M NaOH.
Refer to Exhibit 16-1. What is the pH at the midpoint of the titration?

A) 3.06
B) 3.18
C) 4.44
D) 4.74
E) 5.04

Question 24

Which of the following indicator(s) would be most suitable for the titration of acetic acid (pK_a = 4.74) with NaOH? Your choices are:
I. bromothymol blue (transition range pH 6 to 8)
II. methyl red (transition range pH 4 to 6.3)
III. phenolphthalein (transition range pH 8.3 to 11)

A) I only
B) II only
C) III only
D) I or II only
E) I, II, or III

Question 25

The K_sp expression for silver phosphate, Ag₃PO₄, is

A)
B)
C)
D)
E)

Question 26

Which acid-base titration would yield a titration curve of the general form shown?

A) H₂CO₃ titrated with NaOH
B) NaOH titrated with H₃PO₄
C) Na₃PO₄ titrated with HCl
D) H₂SO₄ titrated with NaOH
E) H₃PO₄ titrated with NaOH

Question 27

Exhibit 16-1 The following question(s) refer to the same titration: a 50.00 mL sample of 0.0950 M acetic acid (K_a = 1.8 × 10^-5) is being titrated with 0.106 M NaOH.
Refer to Exhibit 16-1. What is the pH at the equivalence point of the titration?

A) 5.28
B) 7.00
C) 8.72
D) 9.26
E) Need more information to answer

Question 28

When conducting a titration of an acid by adding base to the equivalence point, it is appropriate to use an indicator which has

A) a pK_a = 7.00
B) a color change at a pH close to the pH at the equivalence point
C) a color change at a pH considerably lower than the pH at the equivalence point
D) a color change at a pH considerably higher than the pH at the equivalence point
E) an easily detected color change regardless of the pH where the change occurs

Question 29

What volume of 0.1060 M NaOH is needed to neutralize a 50.00 mL sample of 0.09500 M HNO₃?

A) 55.79 mL
B) 55.19 mL
C) 50.00 mL
D) 44.81 mL
E) 5.19 mL

Question 30

The greater the K_sp of a salt

A) the lower its solubility in water
B) the greater its solubility in water
C) the more easily it is dissolved in all solvents
D) the more easily it dissolves in non-hydrogen bonding solvents
E) the smaller the mass of solute required to prepare a saturated solution

Question 31

Which statement about the titration of 0.10 M HNO₃ with 0.10 M KOH is not correct?

A) The endpoint always comes before the equivalence point.
B) The pH at the equivalence point is 7.00.
C) The initial pH is 1.00.
D) At the equivalence point the pH increases sharply.
E) The net ionic equation for the titration reaction is H₃O⁺ + OH^- 2 H₂O.

Question 32

Silver chloride has K_sp = 1.8 × 10^-10. What is its molar solubility in water?

A) 9.0 × 10^-11 M
B) 3.6 × 10^-10 M
C) 6.7 × 10^-6 M
D) 9.5 × 10^-6 M
E) 1.3 × 10^-5 M

Question 33

The common ion effect predicts that the solubility of a salt in a solution already containing one of the ions in the salt will be ____ than the solubility in water. This is predictable, because ____.

A) more; Le Chatelier's Principle predicts that the common ion will suppress precipitation
B) more; the presence of the common ion will increase that concentration, hence K_sp will fall
C) less; Le Chatelier's Principle predicts that the common ion will suppress dissolving
D) less; the presence of the common ion will increase that concentration, hence K_sp will rise
E) less; at equilibrium, Q will be less than K_sp

Question 34

When 1.00 L of 0.45 M acetic acid (pK_a = 4.74) is mixed with the exact volume of 0.55 M NaOH required to convert the acid to its conjugate base, at the endpoint the solution will have a

A) pH < 7.00
B) pH > 7.00
C) pH = 7.00
D) pH < pK_a
E) pH = pK_a

Question 35

Which statement is incorrect?

A) water-insoluble salts containing anions that are Bronsted-Lowry bases generally dissolve in solutions of strong bases
B) the solubility of an ionic compound is decreased by the presence of a second solute that provides a common ion
C) amphoteric metal hydroxides are soluble in both highly acidic and highly basic solutions
D) complex ion formation allows the solubilization of many otherwise insoluble metal salts
E) the solubility of water-insoluble salts cannot be affected by the addition of acids or bases to the solution.

Question 36

Silver phosphate is less soluble in Na₃PO₄ than in water because

A) a common ion displaces the solubility equilibrium towards the undissolved solute.
B) some insoluble compounds are amphoteric.
C) the solubility of most salts increases as temperature increases.
D) the formation of complex ions displaces the solubility equilibrium to the right.
E) the solubility of many salts is affected by the pH of the solution.

Question 37

A 50.00 mL sample of 0.0950 M acetic acid (K_a = 1.8 × 10^-5) is being titrated with 0.0848 M NaOH. What is the pH after 28.00 mL of NaOH has been added?

A) 5.04
B) 4.74
C) 4.44
D) 3.18
E) 3.06

Question 38

In which liquid or solution would calcium phosphate have the greatest solubility?

A) 0.1 M aqueous sodium phosphate
B) distilled water
C) 0.1 M aqueous calcium nitrate
D) saturated aqueous calcium hydroxide
E) 0.1 M aqueous phosphoric acid

Question 39

The solubility of silver chloride in water is increased by the addition of NH₃. Why?

A) The solubility of many salts is affected by the pH of the solution.
B) The formation of complex ions displaces the solubility equilibrium to the right.
C) The solubility of most salts increases as temperature increases.
D) Some insoluble compounds are amphoteric.
E) A common ion displaces the solubility equilibrium toward the undissolved solute.

Question 40

Exhibit 16-1 The following question(s) refer to the same titration: a 50.00 mL sample of 0.0950 M acetic acid (K_a = 1.8 × 10^-5) is being titrated with 0.106 M NaOH.
Refer to Exhibit 16-1. What is the pH at the endpoint of the titration?

A) 5.28
B) 7.00
C) 8.72
D) 9.26
E) Need more information to answer

Question 41

Will a precipitate form when 10.0 mL of 0.500 M NaCl is added to 10.0 mL of 0.0500 M AgNO₃? The K_sp for AgCl is 1.8 × 10^-10.

A) Yes, because Q < K_sp
B) Yes, because Q = K_sp
C) Yes, because Q > K_sp
D) No, because Q < K_sp
E) No, because Q > K_sp

Question 42

Answer the following questions:
a. Explain why the solubility of carbonates and phosphates is increased by lowering the pH, but the solubility of chlorides is unaffected by lowering the pH.
b. The solubility of copper(II) carbonate is dramatically increased by the addition of NH₃, but the solubility of copper(II) phosphate is not. Referring to the data below, explain why.

K_sp for copper(II) carbonate = 2.5 ´ 10^-10
K_sp for copper(II) phosphate = 1.4 ´ 10^-37
Cu²⁺(aq) + 4 NH₃(aq)
(Kf = 1.1 ´ 10¹³)

(aq)

Question 43

A student attempts to dissolve a small amount of CuCO₃ in 0.0100 M NaOH. Despite vigorous stirring, the material does not apparently dissolve. However, after a while the student notices that the solid has changed appearance. When the pH of the solution is measured, it is found to have gone down. Explain these observations. Was the original goal of the experiment achieved?

Question 44

In which aqueous solution would the smallest amount of barium sulfate dissolve?

A) 0.1 M Mg(OH)₂
B) 0.1 M Na₂SO₄
C) 0.1 M Al₂(SO₄)₃
D) 0.1 M Ba(NO₃)₂
E) 0.1 M HCl

Question 45

A weak acid-strong base titration is being performed using a visual indicator. By titrating a standard acid sample, it is found that the end point is occurring before the equivalence point.
a. Define endpoint and equivalence point.
b. Explain the probable origin of the error, and how you would modify the titration procedure to eliminate it.

Question 46

In which liquid or solution would CuI(s) have the greatest solubility?

A) 0.1 M aqueous ammonia
B) distilled water
C) 0.1 M aqueous copper(I) nitrate
D) saturated aqueous calcium hydroxide
E) 0.1 M aqueous sodium iodide

Question 47

What is the largest mass of solid NaCl that can be dissolved in 1.000 L of 5.00 × 10^-3 M AgNO₃ without a precipitate of AgCl forming? K_sp for AgCl is 1.8 × 10^-10.

A) 0.78 mg
B) 86 mg
C) 13 mg
D) 2.1 mg
E) 36 ng

Question 48

If solid sodium chloride is added to a saturated aqueous solution of sodium chloride

A) the total dissolved solids concentration in the solution will increase
B) the anion concentration in the solution will increase
C) the cation concentration in the solution will increase
D) both the anion and anion concentrations in the solution will decrease
E) there will be no change in the anion or cation concentrations in the solution

Question 49

A solution contains 0.100 M Pb(NO₃)₂ and 1.00 × 10^-5 M AgNO₃. It is intended to separate out the silver by selective precipitation of AgI. What is the maximum percentage of the total silver that can be recovered free of contamination by PbI₂?

Question 50

A mixture of two metal ions in solution is to be separated by selective precipitation. As the concentration of the precipitating ion is increased, which metal ion precipitates first?

A) The one with the lower value of K_sp for its precipitate
B) The one with the higher value of K_sp for its precipitate
C) The one present at higher concentration
D) The one present at lower concentration
E) Cannot answer the question without more information

Question 51

Which group contains only solutes that would decrease the solubility of barium sulfate?

A) HNO₃, H₂SO₄, HCH₃COO
B) SO₂, CO₂, NH₃
C) Ba(NO₃)₂, Na₂SO₄, H₂SO₄
D) Ba(OH)₂, NaOH, NH₄OH
E) Na₂SO₄, NaOH, NaCH₃COO

Question 52

For a solution containing a dissolved salt, if Q = K_sp

A) the solution is unsaturated and no precipitate forms
B) the solution is unsaturated but is at the point of precipitation
C) the solution is supersaturated and a precipitate forms
D) a precipitate will form from an unsaturated solution
E) the solution is saturated and a precipitate will form if more anions and/or cations are added