Deck 17: Equilibrium in the Aqueous Phase

Question 1

When [H⁺] = 1.0  10^-7 M in water at 25°C, then __________

A) pH = 1.
B) pH = 10^-7.
C) [OH^-] = 1.0  10^-7 M.
D) [OH^-] = 1.0  10⁷ M.
E) [OH^-] = 0 M.

Answer 1

[OH^-] = 1.0  10^-7 M.

Question 2

Ammonia (NH₃) acts as a weak base in aqueous solution. What is the acid that reacts with this base when ammonia is dissolved in water?

A) none, there are no acids in pure water
B) H₂O
C) NH₄⁺
D) trick question, because no acids are present, ammonia cannot act as a base
E) oxygen that always is dissolved in water

Answer 2

H₂O

Question 3

Which of these is a strong acid that ionizes to make a weak acid?

A) H₂SO₃
B) H₂SO₄
C) H₃PO₄
D) HNO₃
E) HCl

Answer 3

H₂SO₄

Question 4

In the following reaction in aqueous solution, the acid reactant is __________, and its conjugate base product is __________. CH₃NH₂ + HSO₄^- CH₃NH₃⁺ + SO₄^2-

A) CH₃NH₂; CH₃NH₃⁺
B) CH₃NH₂; SO₄^2-
C) HSO₄^-; CH₃NH₃⁺
D) HSO₄^-; SO₄^2-
E) HSO₄^-; H₃O+

Question 5

In the Brønsted-Lowry definition of acids and bases, a base __________

A) is a proton donor.
B) is a proton acceptor.
C) forms stable hydrogen bonds.
D) breaks stable hydrogen bonds.
E) corrodes metals.

Question 6

In the following reaction in aqueous solution, the acid reactant is __________ and its conjugate base product is __________. CH₃COOH + NH₃ CH₃COO^- + NH₄⁺

A) CH₃COOH; CH₃COO^-
B) CH₃COOH; NH₄⁺
C) NH₃; CH₃COO^-
D) NH₃; NH₄⁺
E) CH₃COOH; H₃O⁺

Question 7

The degree of ionization of a weak acid __________
(I) varies with the concentration of the acid.
(II) depends on which weak acid it is.
(III) is 100%.
(IV) is greater than 50% but less than 100%.

A) I only
B) II only
C) III only
D) both I and II
E) IV only

Question 8

A solution with pH of 9.50 has a pOH of __________

A) 9.50.
B) 0.50.
C) 4.50.
D) 23.5.
E) 19.0.

Question 9

The base ionization constant K_b describes which of the following reactions for a weak base, B, in aqueous solution?

A) B + H⁺ BH⁺
B) B + H₃O⁺ BH⁺ + H₂O
C) B + H₂O BH⁺ + OH^-
D) B + OH^- BH^- + O^2-
E) BH⁺ + OH^- B + H₂O

Question 10

The degree of ionization of a strong acid is __________

A) dependent on the concentration of the acid.
B) between 1 and 10%.
C) between 10 and 100%.
D) 100%.
E) dependent on which strong acid it is.

Question 11

Which one, A-D, is not related to the water autoionization constant, K_w? If all are related, respond E.

A) [H₃O⁺] [OH^-].
B) 1.0  10^-14 at 25°C.
C) 2 H₂O H₃O⁺ + OH^-
D) pH = 7 at 25°C
E) A-D are all related to K_w.

Question 12

The acid ionization equilibrium constant, K_a, describes the reaction (where HA is a generic weak acid):

A) HA + OH^- H₂O + A^-
B) HA + H₂O H₃O⁺ + A^-
C) HA + H₃O⁺ H₂A⁺ + H₂O
D) HA + H₂A⁺ H₂A⁺ + HA
E) H₃O⁺ + A^- HA + H₂O

Question 13

Which sketch best represents the qualitative molecular view of an aqueous solution of nitric acid? (Water molecules are not shown explicitly.)

A)
B)
C)
D)

Question 14

In the Brønsted-Lowry definition of acids and bases, an acid __________

A) is a proton donor.
B) is a proton acceptor.
C) forms stable hydrogen bonds.
D) breaks stable hydrogen bonds.
E) corrodes metals.

Question 15

Which statement about nitrous acid and nitric acid is correct?

A) They are both weak acids.
B) They are both strong acids.
C) They both have one ionizable proton.
D) Nitrous acid has the formula HNO₃.
E) Nitric acid has the formula HNO₂.

Question 16

A solution with a pOH of 6.92 has an [OH^-] concentration of __________

A) 1.20  10^-7 M.
B) 9.2  10^-6 M.
C) 6.8  10^-6 M.
D) 7.08 M.
E) 6.92 M.

Question 17

When [H⁺] = 4.0  10^-9 M in water at 25°C, then __________

A) pH = 9.40.
B) pH = 7.00.
C) pH = -8.40.
D) pH = 8.40.
E) pH = -9.40

Question 18

Which sketch best represents the qualitative molecular view of an aqueous solution of nitrous acid? (Water molecules are not shown explicitly.)

A)
B)
C)
D)

Question 19

When pure water autoionizes, the following species are produced: __________

A) O^2-, OH^-, H₃O⁺, and H₂O⁺.
B) OH^- and H₃O⁺.
C) O^2- and H₄O²⁺.
D) H⁺ and OH^-.
E) 2H⁺ and O^2-.

Question 20

Carbon dioxide in the atmosphere dissolves in rainwater to make rain __________

A) neutral.
B) acidic.
C) basic.
D) gray.
E) wet.

Question 21

A solution with a pOH of 4.3 has a [H⁺] of __________

A) 6.8  10^-9 M.
B) 3.2  10^-4 M.
C) 4.8  10^-5 M.
D) 2.0  10^-10 M.
E) 4.3 M.

Question 22

A substance that can act as both an acid and base is __________

A) amphibious.
B) amphiprotic.
C) bacidic.
D) androgynous.
E) acibasic

Question 23

Bert and Ernie were determining the pH of their goldfish's water. Bert's pH meter was in the fix-it shop, so Ernie used his pOH meter instead. Ernie insists that he can use the reading to calculate the pH by simple addition or subtraction that he learned in second grade, but Bert (showing off as usual) claims that the calculation involves finding an antilogarithm, dividing, and then finding a logarithm. This would make use of all the math he learned in third grade. Who is right?

A) Ernie
B) Bert
C) both
D) neither

Question 24

In evaluating the pH of an aqueous weak acid solution, __________ usually can be ignored.

A) the concentration of the weak acid
B) the concentration of hydronium ion produced by the autoionization of water
C) the reaction of the weak acid with water
D) the concentration of the ionized hydronium ion
E) the concentration of the conjugate base

Question 25

Phosphoric acid is a triprotic acid, ionizing in the following sequential steps: H₃PO₄ + H₂O H₂PO₄^- + H₃O⁺ K_a

H₂PO₄^- + H₂O HPO₄^2- + H₃O⁺ K_a
HPO₄^2- + H₂O PO₄^3- + H₃O⁺ K_a
What is the K_b expression for the base, sodium phosphate?

A)
B)
C)
D)
E)

Question 26

What is the actual concentration of the molecular form of HF in a 1.0 M HF solution given that K_a of HF is 6.8  10^-4?

A) 2.6  10^-2 M
B) 0.97 M
C) 1.59 M
D) 6.8  10^-4 M
E) 1.0 M

Question 27

Pure water at any temperature has __________

A) a pH less than 7.
B) a pOH more than 7.
C) [H₃O⁺] = [OH^-].
D) pH = 7.
E) no hydronium ions in it.

Question 28

When values of K_a are small (e.g., 1  10^-5) and concentrations of weak acids [HA] are relatively large (e.g., 0.10 M), the hydronium ion concentration of the solution can be calculated using which expression?

A) [H⁺] = K_a
B) [H⁺] = K_a[HA]
C)
D) [H⁺] = K_aK_b[HA]
E) [H⁺] = K_a[HA]/[A^-]

Question 29

What is the pOH of a 0.20 M solution of ammonia? The K_b value for ammonia is 1.8  10^-5

A) 4.44
B) 4.74
C) 0.70
D) 2.72
E) 3.38

Question 30

The degree of ionization __________

A) increases with increasing concentration of a weak acid.
B) decreases with increasing concentration of a weak acid.
C) does not change with changing concentration of a weak acid.
D) is not related to the concentration of a weak acid.
E) is independent of the composition of the weak acid.

Question 31

What is the actual concentration of molecular NH₃ in a 0.200 M solution of ammonia? The K_b value for ammonia is 1.80  10^-5.

A) 0.200 M
B) 0.198 M
C) 1.80  10^-5 M
D) 1.90  10^-3 M
E) 3.6  10^-6 M

Question 32

Sometimes liquid ammonia, NH₃, is used as a solvent rather than water. Which expression defines the ammonia autoionization counterpart of K_w?

A) [H₃O⁺][OH^-]
B) [NH₃][NH₄⁺]
C) [NH₂^-][NH₄⁺]
D) [H₃O⁺][NH₂^-]
E) [NH₄⁺][OH^-]

Question 33

What is the hydronium ion concentration of a 0.20 M solution of ammonia? The K_b value for ammonia is 1.8  10^-5

A) 2.8  10^-10
B) 5.5  10^-10
C) 1.8  10^-5
D) 5.2  10^-12
E) 1.9  10^-3

Question 34

Which statement, A-D, is not correct? If all are correct, respond E. Pure water at 25°C has __________

A) K_w = 1.0  10^-14.
B) pOH = 7.
C) [H₃O⁺] = [OH^-].
D) pH = 7.
E) A-D are all correct.

Question 35

Which expression defines the autoionization constant for water, K_w?

A) [H₃O⁺][OH^-]
B) [H₂O][H₃O⁺]
C) [OH^-][H₂O]
D) [H₄O²⁺][O^2-]
E) [H₂O][H₂O]

Question 36

What is the pH of a 0.20 M solution of ammonia? The K_b value for ammonia is 1.8  10^-5

A) 9.56
B) 9.26
C) 4.74
D) 11.28
E) 2.72

Question 37

What is the concentration of [OH^-] in a 0.20 M solution of ammonia? The K_b value for ammonia is 1.8  10^-5.

A) 3.6  10^-6 M
B) 1.8  10^-5 M
C) 0.20 M
D) 1.9  10^-3 M
E) 4.2  10^-4 M

Question 38

What is the pH of a 0.010 M solution of acetic acid? K_a for acetic acid is 1.8  10^-5

A) 2.74
B) 4.74
C) 2.00
D) 3.37
E) 6.74

Question 39

Phosphoric acid is a triprotic acid, ionizing in the following sequential steps: H₃PO₄ + H₂O H₂PO₄^- + H₃O⁺ H₂PO₄^- + H₂O HPO₄^2- + H₃O⁺ HPO₄^2- + H₂O PO₄^3- + H₃O⁺ What is the K_b expression for the base, dihydrogen phosphate?

A)
B)
C)
D)
E)

Question 40

What is the hydronium ion concentration of a 0.010 M solution of acetic acid? K_a for acetic acid is 1.8  10^-5

A) 1.8  10^-3
B) 1.8  10^-5
C) 1.0  10^-2
D) 1.8  10^-7
E) 4.2  10^-4

Question 41

Phosphoric acid is a triprotic acid, ionizing in the following sequential steps: H₃PO₄ + H₂O H₂PO₄^- + H₃O⁺ H₂PO₄^- + H₂O HPO₄^2- + H₃O⁺ HPO₄^2- + H₂O PO₄^3- + H₃O⁺ Which equilibrium is most important in determining the pH of a solution of sodium phosphate?

A) HPO₄^2- + H₂O PO₄^3- + H₃O⁺
B) H₃PO₄ + H₂O H₂PO₄^- + H₃O⁺
C) PO₄^3- + H₂O HPO₄^2- + OH^-
D) H₂PO₄^- + H₂O H₃PO₄ + OH^-
E) 2H₂O H₃O⁺ + OH^-

Question 42

What is the solubility of barium sulfate in a solution containing 0.050 M sodium sulfate? The K_sp value for barium sulfate is 1.1  10^-10.

A) 7.4  10^-6 M
B) 5.5  10^-11 M
C) 1.0  10^-5 M
D) 2.2  10^-9 M
E) 1.1  10^-10 M

Question 43

Suppression of the solubility of one ion by the addition of the counter-ion in its insoluble salt is known as the __________

A) ionic suppression effect.
B) counter ion effect.
C) common ion effect.
D) excession effect.
E) supersaturation effect.

Question 44

Which one of the following salts forms aqueous solutions with pH = 7?

A) Na₂S
B) NaBr
C) NaClO₂
D) NaNO₂
E) Na₂CO₃

Question 45

Stalactites-the long, icicle-like formations that hang from the ceilings of caves-are formed from recrystallizing minerals like calcite (calcium carbonate). The K_sp of calcium carbonate is 4.5  10^-9. What is the concentration of a saturated calcium carbonate solution?

A) 0.00104 M
B) 4.5  10^-9 M
C) 6.7  10^-5 M
D) 2.25  10^-9 M
E) 4.5  10^-5 M

Question 46

Aqueous solutions of __________ are basic.

A) NaF
B) NaCl
C) NaBr
D) NaI
E) KI

Question 47

The solubility product for an insoluble salt with the formula M₃X is written as __________, where s is the molar solubility.

A) K_sp = s²
B) K_sp = 4s³
C) K_sp = 27s⁴
D) K_sp = 3s³
E) K_sp = 3s⁴

Question 48

Which one of the following salts forms aqueous solutions with pH = 7?

A) Na₂S
B) NaCl
C) NaClO₂
D) NaNO₂
E) Na₂CO₃

Question 49

The solubility product for an insoluble salt with the formula M₂X is written as __________, where s is the molar solubility.

A) K_sp = s²
B) K_sp = 4s³
C) K_sp = 4s²
D) K_sp = 2s³
E) K_sp = 2s²

Question 50

The solubility product for an insoluble salt with the formula MX₂ is written as __________, where s is the molar solubility.

A) K_sp = s²
B) K_sp = 4s³
C) K_sp = 4s²
D) K_sp = 2s³
E) K_sp = 2s²

Question 51

Which one of the following salts does not produce a basic solution when dissolved in water?

A) Na₂S
B) NaBr
C) NaClO₂
D) NaNO₂
E) Na₂CO₃

Question 52

The solubility product for an insoluble salt with the formula MX₃ is written as __________, where s is the molar solubility.

A) K_sp = s²
B) K_sp = 4s³
C) K_sp = 27s⁴
D) K_sp = 3s³
E) K_sp = 3s⁴

Question 53

Magnesium sulfate can be obtained at a drugstore as Epsom salts. The monohydrate is found as the mineral kieserite. What would be the pH of a 500 mL aqueous solution containing 1.62 g kieserite? The value for sulfuric acid is 1.2  10^-2.

A) nearly 14
B) slightly above 7.00
C) 7.00
D) slightly below 7.00
E) nearly 1.00

Question 54

The pH of an aqueous sodium fluoride (NaF) solution is __________ because __________

A) 7; sodium fluoride is a simple salt.
B) above 7; fluoride is a weak base.
C) below 7; fluoride reacts with water to make hydrofluoric acid.
D) about 7; fluoride is a weak base, but produces hydrofluoric acid, and these two neutralize one another.
E) 0; sodium fluoride is a salt not an acid or a base.

Question 55

The stronger the acid, __________

A) the stronger its conjugate base.
B) the weaker its conjugate base.
C) the more concentrated the acid.
D) the less concentrated the conjugate base.
E) the more concentrated the conjugate base.

Question 56

If 500 mL of a solution containing 1.00 g of barium nitrate (199.33 g/mol) is combined with 500 mL of a solution containing 1.00 g of sodium sulfate (142.03 g/mol), will there be a precipitate? Barium sulfate has a K_sp value of 9.1  10^-11.

A) Barium sulfate will precipitate.
B) There will be no precipitate.
C) It is impossible to tell.
D) Sodium nitrate will precipitate.
E) Barium nitrate will precipitate.

Question 57

What is the pH of a 0.15 M solution of ammonium bromide? The K_b value for ammonia is 1.8  10^-5.

A) 11.22
B) 7.00
C) 2.78
D) 5.04
E) 10.08

Question 58

Magnesium sulfate can be obtained at a drugstore as Epsom salts. The monohydrate is found as the mineral kieserite. What would be the pH of a 500 mL aqueous solution containing 1.62 g kieserite? The value for sulfuric acid is 1.2  10^-2.

A) 7.00
B) 6.85
C) 7.15
D) 3.55
E) 12.22

Question 59

Stalactites-the long, icicle-like formations that hang from the ceilings of caves-are formed from recrystallizing minerals like calcite (calcium carbonate). The K_sp of calcium carbonate is 4.5  10^-9. What is the minimum volume of dripping water saturated with calcium carbonate that would be required to form a small stalactite that had a mass of 1.000 kg?

A) 1.4  10⁶ L
B) 1.5  10⁵ L
C) 4.5  10⁹ L
D) 4.5  10⁴ L
E) 1.5  10⁴ L

Question 60

When sodium chloride is added to a saturated solution of lead(II) chloride, some of the lead(II) chloride precipitates. This phenomenon is called __________

A) the common ion effect.
B) selective precipitation.
C) supersaturation.
D) a solubility anomaly.
E) deionization.

Question 61

Which of the following compounds would not have a pH dependent solubility?

A) Mg(OH)₂
B) PbS
C) AgI
D) Na₂O
E) pbS

Question 62

Acid-base indicators need to have very intense colors so that a very low concentration is visible. Could there be a problem in using too much indicator?

A) No, the colors would just be darker.
B) No, indicators are inert.
C) Yes, more indicator requires more extreme pH values to change color.
D) Yes, the indicator could affect the acid-base chemistry being measured.
E) No, the colors would just be sharper.

Question 63

A phosphate buffer solution (25.00 mL sample) used for a growth medium was titrated with 0.1000 M hydrochloric acid. The components of the buffer were sodium monohydrogenphosphate and sodium dihydrogenphosphate. The first endpoint occurred at a volume of 10.32 mL; the second occurred after an additional 18.62 mL was added (or total volume = 28.94 mL). What was the approximate ratio of HPO₄^2- to H₂PO₄^- in the buffer?

A) 1:1
B) 4:5
C) 5:3
D) 5:4
E) 3:5

Question 64

As the pH decreases, the solubility of __________ would increase.

A) lead(II) chloride
B) silver(I) iodate
C) calcium carbonate
D) mercury(I) bromide
E) silver(I) chloride

Question 65

Which of the following is not a buffer system? A solution containing roughly equal concentrations of __________

A) fluoride ion and hydrofluoric acid.
B) bromide ion and hydrobromic acid.
C) phosphate ion and hydrogen phosphate ion.
D) carbonate ion and hydrogen carbonate ion.
E) phosphoric acid and dihydrogen phosphate ion.

Question 66

A solution that contains a weak acid and its conjugate base in roughly equal concentrations is __________

A) neither acidic or basic.
B) a half-acid solution.
C) a buffer.
D) a heterogeneous mixture.
E) neutral.

Question 67

The following titration curve is most likely to be associated with

A) the titration of a strong acid with a strong base titrant.
B) the titration of a weak acid with a strong base titrant.
C) the titration of a strong base with a strong acid titrant.
D) the titration of a weak base with a strong acid titrant.

Question 68

The most suitable acid-base indicator for a titration of acetic acid with NaOH has __________

A) a pK_in = pK_a of the acid.
B) a pK_in = pK_b of the base.
C) pK_in = pH of a sodium acetate solution.
D) pK_in = pH of an acetic acid solution.
E) pK_in = 7.0

Question 69

A 0.500 g sample of an unknown substance was titrated with a 0.1 M HCl solution. Another 0.500 g sample of it was titrated with a 0.1 M NaOH solution. The resulting titration curves are illustrated here. Given the following possibilities, what is the sample?

A) Na₂CO₃
B) NaHCO₃
C) H₂CO₃
D) CO₂
E) There is no way to tell.

Question 70

Halfway to the equivalence point in a titration curve of a weak acid with a strong base, __________

A) nothing is happening yet.
B) the pH = pK_a of the weak acid.
C) pH = 3.5 exactly.
D) pH = pK_a of the indicator.
E) the pH has not yet changed.

Question 71

What are the characteristics of a pH indicator? A pH indicator __________
(I) is a weak acid.
(II) has different characteristic colors in protonated and unprotonated forms.
(III) changes color when the pH is near its pK_a.

A) I only
B) II only
C) both II and III
D) I, II, and III
E) I and III

Question 72

When an acetic acid solution is titrated with sodium hydroxide, the slope of the titration curve (pH vs volume of NaOH added) increases when sodium hydroxide is first added. This change shows that __________.

A) nothing is happening during this part of the titration.
B) the reaction is very slow during this part of the titration.
C) a more concentrated solution of NaOH needs to be present to initiate the reaction.
D) acetic acid is being converted to sodium acetate.
E) the pH is not affected until all the acetic acid is consumed.

Question 73

The following titration curve is most likely to be associated with __________

A) the titration of a strong acid with a strong base titrant.
B) the titration of a weak acid with a strong base titrant.
C) the titration of a strong base with a strong acid titrant.
D) the titration of a weak base with a strong acid titrant.

Question 74

A solution of sulfuric acid (H₂SO₄, 25.00 mL) was titrated to completion with 34.55 mL of 0.1020 M sodium hydroxide. What was the concentration of the sulfuric acid?

A) 0.07048 M
B) 0.1410 M
C) 0.2819 M
D) 0.0353 M
E) 0.0533 M

Question 75

In a titration of monoprotic acids and bases, there is a large change in pH __________

A) at the point where pH = pK_a of the acid.
B) when the volume of acid is exactly equal to the volume of base.
C) when the concentration of acid is exactly equal to the concentration of base.
D) when the number of moles of acid is exactly equal to the number of moles of base.
E) at the point where pH = pK_b of the base.

Question 76

What is indicated by the shape of the titration curve?

A) A diprotic acid was titrated with a strong base.
B) A triprotic acid was titrated with a strong base.
C) A diprotic base was titrated with a strong acid.
D) A triprotic base was titrated with a strong acid.
E) A strong acid was titrated with a strong base.

Question 77

A solution of hydrochloric acid (HCl, 25.00 mL) was titrated to completion with 34.55 mL of 0.1020 M sodium hydroxide. What was the concentration of the hydrochloric acid?

A) 0.07048 M
B) 0.1410 M
C) 0.2819 M
D) 0.0353 M
E) 0.0533 M

Question 78

A phosphate buffer solution (25.00 mL sample) used for a growth medium was titrated with 0.1000 M hydrochloric acid. The components of the buffer were sodium monohydrogenphosphate and sodium dihydrogenphosphate. The first endpoint occurred at a volume of 10.32 mL, and the second occurred after an additional 18.62 mL was added, for a total volume of 28.94 mL. What was the total concentration of phosphate (in any form) in the buffer?

A) 0.03992 M
B) 0.1198 M
C) 0.04243 M
D) 0.07448 M
E) 0.08382 M

Question 79

Bromocresol green is yellow in its acidic form and blue in its basic form. When is it green?

A) at the equivalence point in a titration
B) in its neutral form
C) when the solution pH equals its pK_a
D) when the solution pH equals 7
E) at the midpoint in a titration

Question 80

Acid-base indicators change color __________

A) exactly when pH = pK_a of the indicator.
B) generally over a range of 1 or 2 pH units.
C) at pH = 7.
D) always between a pH of 6 and 8.
E) at the midpoint of a titration.