Direction:  For exercises that require problem solving, you need to show the calculation set up leading to the answer in order to receive credit. Make sure your answers contain the correct units. YOU WILL NEED TO HAVE A TABLE OF IONIZATION CONSTANTS FOR THIS HOMEWORK.

1. In each reaction, identify the Bronsted-Lowry acid and its conjugate base, and the Bronsted-Lowry base     and its conjugate acid.

1. C₅H₅N(aq) +  H₂O(l)    C₅H₅NH⁺(aq)  +  OH^‒(aq)

1. H₂O(l) +  CO₃^2‒(aq)    HCO₃^‒(aq)  +  OH^‒(aq)

1. NH₃(aq) +  H₂O(l)    NH₄⁺(aq)  +  OH^‒(aq)

1. CH₃CO₂H(aq) +  H₂O(l)    CH₃CO₂^‒(aq)  +  H₃O⁺(aq)

2. Classify each substance is a weak acid or weak base. Then (a) write a balanced equation for its reaction in      water (include all physical labels) and (b) an equilibrium expression (K_a or K_b) for the reaction.

1. HCN

1. CH₃-NH₂ (methylamine)

1. NH₄⁺

3. For each solution, determine the [H₃O⁺], [OH^‒], and pH. Show problem set ups and set the answers with          correct sig figs. Provide correct units where applicable. Assume K_w = 1.00 x 10^‒14

1. 35 M HNO₃

Answers: [H₃O⁺] = 0.35 M; [OH^‒] = 2.9 x 10^‒14 M; pH = 0.46

1. A solution that is 0.045 M HClO₄ and 0.068 M HCl

Answers: [H₃O⁺] = 0.113 M; [OH^‒] = 8.85 x 10^‒14 M; pH = 0.947

1. A solution that is 0.655% HNO₃ by mass (i.e. in a 100-g solution, HNO₃ makes up 0.655 g). The density             of the solution is 1.01 g/mL. [Hint: Consider a solution that weighs exactly 100. g (3 sig figs)]

Answers: [H₃O⁺] = 0.105 M; [OH^‒] = 9.52 x 10^‒14; pH = 0.979

4. Determine the [H₃O⁺] and pH of a 0.100 M benzoic acid. Benzoic acid is a monoprotic acid with a formula      HC₇H₅O₂ (acidic proton is indicated by bold H). Benzoic acid has a K_a = 6.5 x 10^‒5 at 25 

Condensed structure

of benzoic acid

Answers:  [H₃O⁺] = 2.5 x 10^‒3 M; pH = 2.59

5. Determine the [H₃O⁺] and percent ionization of a 0.125 M HCN solution at 25  At this temperature,            hydrocyanic acid has K_a = 4.9 x 10^‒10.

Answers:  [H₃O⁺] = 7.8 x 10^‒6 M; % ionization = 0.0063%

6. A 0.148 M solution of a generic weak monoprotic acid (HA) has a percent ionization of 1.55% at certain          Determine the ionization constant (K_a) for the acid.

Answer:  K_a = 3.61 x 10^‒5

7. Find the pH of each mixture of acids. Show problem set up and explain briefly your strategy for determining     the solution [H₃O⁺] and its pH.

1. 150 M HNO₂ and 0.085 M HNO₃

Answer:  pH = 1.07

1. 050 M CH₃CO₂H and 0.050 M HCN

Answer:  pH = 3.02

8. Consider a solution of 4.8 x 10^‒4 M Sr(OH)₂. Calculate the [OH^‒], [H₃O⁺], pH, and pOH.

Answers:  [OH^‒] = 9.6 x 10^‒4 M; [H₃O⁺] = 1.0 x 10^‒11 M; pOH = 3.02; pH = 10.98

9. What volume, in mL, of 0.855 M KOH solution is required to make 3.55 L of a solution with pH = 12.4 ?

Answer: 104 mL

10. Caffeine (C₈H₁₀N₄O₂) contains nitrogens and acts a weak base with a pK_b = 10.4. The molar mass of caffeine is 194.19 g/mol. Calculate the pH of a solution containing a caffeine concentration of 455 mg/L.

                                              C₈H₁₀N₄O₂(aq)  +  H₂O(l)    HC₈H₁₀N₄O₂⁺(aq)  +  OH^‒(aq)

Answer: pH = 7.48

11. Determine whether each ion acts as a weak acid, weak base, or pH-neutral substance in aqueous solution.     For each ion, explain its acid or base behavior and, in the case of a weak acid or weak base, write a balanced      equation that shows how the ion acts as an acid or a base in water.

1. ClO^‒

1. Co³⁺

1. NO₂^‒

1. C₂O₄^2‒

1. Mg²⁺

(Question #11 continued)

1. HSO₄^‒

1. SO₄^2‒

1. BrO₂^‒

1. Cr³⁺

1. g) NH₄⁺

12. Determine whether each of the following substances, when dissolved in water, will produce an acidic, basic,    or pH-neutral For each case, explain its acid or base behavior. In case where the substance produces       an acidic or basic solution, write a balanced equation that shows how it acts as an acid or a base. For problems       (c) and (d), you may need to provide K_a and K_b values in your explanation.

1. K₂CO₃

1. Na₂SO₃

1. CaHPO₄

1. NH₄CN

13. Calculate the equilibrium concentrations of H₂SO₃, HSO₃^‒, SO₃^‒2, and the pH of a 0.500 M H₂SO₃

H₂SO₃(aq)  +  H₂O(l)    HSO₃^‒(aq)  +  H₃O⁺(aq)          K_a1 = 1.6 x 10^‒2

HSO₃^‒(aq)  +  H₂O(l)    SO₃^2‒(aq)  +  H₃O⁺(aq)            K_a2 = 6.4 x 10^‒8

Answers:  [H₂SO₃] =  0.418 M; [HSO₃^‒] = 0.082 M; [SO₃^‒2] = 6.4 x 10^‒8 M; pH = 1.09

14. If each of the salts listed below were dissolved in water to give a 0.10 M solution, which solution would have the highest pH ?  Which would have the lowest pH ?  Explain your choices by providing K_a or K_b values for

1. Na₂S c) NaH₂PO₄     e) CH₃CO₂Na
2. Na₃PO₄ d) NaF f) AlCl₃

15. For each of the following reactions, predict whether the equilibrium lies predominantly to the left or to the       Explain your predictions using K_a or K_b values.

1. NH₄⁺(aq) +  Br^‒(aq)    NH₃(aq)  +  HBr(aq)

1. HPO₄^2‒(aq) +  CH₃CO₂^‒(aq)    PO₄^3‒(aq)  +  CH₃CO₂H(aq)

1. [Fe(H₂O)₆]³⁺(aq) +  HCO₃^‒(aq)    [Fe(H₂O)₅(OH)]²⁺(aq)  +  H₂CO₃(aq)

16. Equal molar quantities of ammonia and sodium dihydrogen phosphate (NaH₂PO₄) are mixed. Write a   balanced, net ionic equation for the acid-base reaction that can, in principle, occur. Use K_a and K_b values to           determine if the equilibrium lies to the right or the left.

17. Identify the Lewis acid and Lewis base from among the reactants in each equation, then briefly explain your

1. Ag⁺(aq) +  2 NH₃(aq)    [Ag(NH₃)₂]⁺(aq)

1. F^‒(aq) +  BF₃(aq)    BF₄^‒(aq)

(Question 17 continued)

1. SO₂(aq) +  H₂O(l)    H₂SO₃(aq)

1. Ni²⁺(aq) +  6 H₂O(l)    [Ni(H₂O)₆]⁺²(aq)

1. Cr(OH)₃(s) +  3 H₃O⁺(aq)    Cr³⁺(aq)  +  6 H₂O(l)

1. Zn(OH)₂(s) +  2 OH^–(aq)    [Zn(OH)₄]^2–(aq

18. Of each pair of acids, circle the stronger acid and briefly explain your choice. [Hint: use concepts such as bond            energy, electronegativity, bond polarity, and inductive effect in your explanation.]

1. HF or  HBr

1. H₂S or  H₂Se

1. HClO or  HIO

1. H₂O or  HF

(Question 18 continued)

1. HClO₃ or  HClO

1. CH₃CO₂H or  CF₃CO₂H

1. HOCN or  HCN

19. Calculate the pH of the solution that results from mixing 25.0 mL of 0.14 M formic acid (HCO₂H) and 50.0     mL of 0.070 M NaOH. [Note: Formic acid is a weak monoprotic acid that reacts with a base to produce             formate ion as follows:  HCO₂H(aq)  +  OH^‒(aq)    HCO₂^‒(aq) +  H₂O(l); the K_a of HCO₂H is 1.8 x 10^‒4 (at

25 C)]

Answer: pH = 8.21

20. A solution of hydrofluoric acid (HF) is prepared by dissolving enough HF in water to produce a solution with pH = 2.30. The K_a of hydrofluoric acid is 7.2 x 10^‒4.

1. Calculate the equilibrium concentrations of HF and F^‒.

Answers: [HF]_equil = 0.035 M; [F^‒] = 0.0050 M

1. Calculate the mass, in g, of HF initially dissolved per liter solution.

Answers:  0.80 g per 1 L

21. Calculate the pH for a 0.025 M solution of sodium oxalate (Na₂C₂O₄). Oxalic acid (H₂C₂O₄) is a diprotic acid   with the following dissociation constants:

H₂C₂O₄  +  H₂O    HC₂O₄^‒  +  H₃O⁺       K_a1 = 5.9 x 10^‒2

HC₂O₄^‒  +  H₂O    C₂O₄^2‒  +  H₃O⁺        K_a2 = 6.4 x 10^‒5

Answer:  pH = 8.30

22. The acid form of pyridine, which contains a single acidic proton, can be synthesized with a different     substituent X, where X is an atom such as Cl or a group such as CH₃. The table below gives the K_a values for a        variety of substituted pyridines.

Atom X or Group X                              K_a of substituted pyridines

NO₂                                                         5.9 x 10^‒2

Cl                                                            1.5 x 10^‒4

H                                                             6.8 x 10^‒6

                                           CH₃                                   1.0 x 10^‒6

1. Suppose each substituted pyridine is dissolved in sufficient water to give a 0.050 M solution. Which                  solution would have the highest pH ?  The lowest pH ?

1. Which of the substituted pyridines is the strongest Bronsted base? Which is the weakest Bronsted base?