Chapter 17 Practice Test

Question 1

Calculate the equilibrium constant for the reaction below given that Ka for sulfurous acid is 1.5 × 10−2.

H2SO3(aq) + OH-(aq) H2O(l) + HSO3^-(aq)

Answer 1

1.5 × 10^12

Question 2

Suppose you begin with 100 mL each of 0.10 M acetic acid and 0.1 M sodium hydrogen sulfate and 0.10 M hydrochloric acid. If you add to each of these solutions 0.005 moles of sodium hydroxide, which one of the resulting solutions will have the lowest pH?

Answer 2

hydrochloric acid

Question 3

In the following reaction,

HCl(aq) + NH3(aq) NH4Cl(aq)

the first intermediate reaction is the result of the strong acid, HCl(aq), and the weak base, NH3, interacting with the aqueous solution (i.e., water). This is also the final reaction.

First intermediate reaction:

H3O^+(aq) + Cl^-(aq) + NH3(aq) NH4^+(aq) + Cl^-(aq) + H2O(l)

Second intermediate reaction:

Nh3(aq) + H2O(l) NH4^+(aq) + OH-(aq)

Third intermediate reaction:

H3O^+(aq) + OH^-(aq) 2H2O(l)

Final reaction:

H3O^+(aq) + NH3(aq) NH4^+(aq) + H2O(l)

If all the reactants and products in the two intermediate reactions are added, the result is the final reaction. If Kb is the equilibrium constant for the second intermediate reaction and (1 / Kw ) is the equilibrium constant for the third intermediate reaction, what is the equilibrium constant for the final reaction, KFinal?

Answer 3

KFinal = Kb (1 / Kw ) = Kb / Kw

Question 4

Look at the reaction of hydrochloric acid, a strong acid, and ammonia, a weak base, to (potentially) yield ammonium chloride.

HCl(aq) + NH3(aq) NH4Cl(aq)

One of the ions created by the dissociation of HCl in water is Cl −. What is the charge on the ammonium ion before it bonds with the Cl − ion?

Answer 4

+1

Question 5

If 1.0 M solution of HNO2 is added to a solution that was already 0.2 M in NO2−, what is the pH of the solution for HNO2Ka = 5.6 x 10–4?

Answer 5

2.6

Question 6

Look at the following reaction:

PCl5(g) PCl3(g) + Cl2(g)

Which of the following events would not cause the equilibrium to shift to the left?

Answer 6

Remove some Cl2(g).

Question 7

A conjugate pair consists of two substances that are exclusively related by the gain or loss of a single hydrogen ion (H + ). Which of the following substance pairs is not a conjugate pair?

Answer 7

H3O and OH −

Question 8

Suppose that you have an acidic buffer (pH < 7) that is formed by adding 1.12 moles of formic acid (HCOOH, Ka = 1.77 × 10−4 ) and 0.89 moles of sodium formate (NaHCOO) to water to make 1 L of solution. The weak acid in this reaction is formic acid, HCOOH. Its conjugate base, derived by removing a H + ion, is HCOO−. What is the pH of the acidic buffer solution?

Answer 8

3.65

Question 9

Suppose that you add 0.35 M HCl to a buffer solution that consists of 0.95 M HCOOH and 0.53 M HCOO−. You have 1 L of the final solution. In this particular event, what is the acid buffer capacity of the solution?

Answer 9

0.18 M

Question 10

1 L of an acidic buffer solution is made up of 1.53 M HCOOH and 1.53 M HCOO− is mixed with 0.03 mol HCl. What is the pH of the buffer solution (Ka = 1.77 × 10−4 )?

Answer 10

3.74

Question 11

Suppose that we do a partial neutralization of 0.48 moles of NH3 by adding 0.18 moles of HCl to make 1.0 L of a basic buffer solution. What is the equilibrium pH of the buffer solution? Ka = 5.6 × 10−10 and Kw = 1.0 × 10−14.

Answer 11

9.48

Question 12

Suppose that we do a partial neutralization of 0.25 moles of NH3 by adding 0.15 moles of HCl to make 1.0 L of a basic buffer solution. If x = [OH − ], which of the following gives the correct theoretical equilibrium concentrations for OH −, NH3, and NH4+?

Answer 12

[OH − ] = x;
[ NH3 ] = 0.10 − x;
[ NH4+ ] = 0.15 + x

Question 13

Suppose that you want to prepare a buffer solution with a pH of 4.58. What is the correct value of [A− ] / [HA] for the correct buffer solution for this event?

Answer 13

0.692

Question 14

The Henderson-Hasselbalch equation is based on the following equation:

Ka = [H3O+ ] [A− ] / [HA].

Which of the following correctly identifies the components of this equation?

Answer 14

Ka is the acid dissociation constant. [H3O+ ] is the hydronium ion concentration. [A− ] is the concentration of the conjugate base. [HA] is the concentration of the conjugate acid.

Question 15

Suppose you have a 59 mL solution of 0.14 M HCl that is being titrated with 0.19 M NaOH. Before the equivalence point, you stop the titration. At this point, 25 mL of NaOH have been added. What is the change in pH at this point?

Answer 15

0.59

Question 16

The equivalence point in a titration reaction is __________________

Answer 16

the point at which the stoichiometric amount of the standardized solution has been added to the unknown solution.

Question 17

Look at the equivalence point of the titration reaction of acetic acid (with NaOH). After all of the acetic acid has been consumed, the CH3COO− interacts with water to form acetic acid and hydroxide.

CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq),

where Kb = 5.6 × 10−10

Suppose that 150 mL of 0.12 M acetic acid has been titrated with 150 mL of 0.12 M NaOH. What is the pH of the solution at this point?

Answer 17

8.8

Question 18

In this titration reaction, we were told that the initial pH of the solution was 2.9. What were the initial concentrations of H + and OH −?

CH3COOH(aq) + OH-(aq) CH3COOO-(aq) + H2O(l)

Answer 18

[H + ] = 1.26 × 10−3 M;
[OH − ] = 7.94 × 10−12 M

Question 19

Look at the plot (pH vs. volume of NaOH) for the titration of diprotic oxalic acid. The two deprotonation reactions that occur in this titration event are:

(Reaction 1) H2(Ox)(aq) + H2O(l) H(Ox)-(aq) + H3O+(aq)

Ka= 5.9 x 10^-2

(Reaction 2) H(Ox)-(aq) + H2O(l) (Ox)^2-(aq) + H3O^+(aq)

Ka = 6.4 x 10^-4

Which statement correctly describes point D in the plot?

Answer 19

Point D is the equivalence point for the first deprotonation event. At this point, it is assumed that all of the H2(Ox) has been consumed via the first deprotonation.

Question 20

Look again at Point I of this reaction. Oxalate anion (Ox)2− is in solution. It is a weak base. Two of the reactions can occur at this point.

(Reaction 3) (Ox)^2-(aq) + H2O(l) H(Ox)^-(aq) + OH-(aq)

Kb= 1.6 x 10^-10

(Reaction 4) H(Ox)-(aq) + H2O(l) H2(Ox)(aq) + OH-(aq)

Kb = 1.7 x 10^-13

Assume that Reaction 3 is the dominant equilibrium to consider when determining the solution’s pH. What is the pH of the solution if we start with 100 mL of 0.12 M H2(Ox) and titrate it with 0.12 M NaOH? Use the plot for the amounts of NaOH added during the reaction.

Answer 20

8.4

Question 21

Use the plot to determine the change in pH between Points G and H. Assume that we started with 100.0 mL of 0.140 M NH3 titrated with 0.140 M HCl.

Answer 21

0.477

Question 22

Look at the plot for the titration of a weak base, NH3, with a strong acid, HCl. Suppose that we have 100.0 mL of 0.100 M NH3 titrated with 0.100 M HCl. At Point A, we have only NH3 in aqueous solution.

NH3(aq) + H2O(l) NH4^+(aq) + Oh-(aq)

where Kb = 1.8 x 10^-5

What is the pH of the solution at the initialization point

Kw= 1.0 x 10^-14?

Answer 22

11.13

Question 23

Which of the following best defines a pH indicator?

Answer 23

a substance whose acid is a different color than its conjugate base

Question 24

If we rearrange the Ka expression for the reaction below we get [H3O+ ] / (Ka ) = [HIn] / [ In− ].

Which statement about this rearranged Ka expression is not correct?

HIn(aq) + H2O(l) H3O^+(aq) + In-(aq)
(colorless) (color)

Answer 24

If [H3O+ ] > Ka, you have an indicator in the deprotonated form. If [H3O+ ] < Ka, you have an indicator in the protonated form.

Question 25

How many moles of AgBr would dissolve in 1 L of water if Ksp for AgBr is 5.0 × 10−13?

AgBr(s) Ag^+(aq) + Br-(aq)

Answer 25

7.1 × 10−7 M