Problems -- Chapter 16 1.

Write balanced chemical equations for the important equilibrium that is occurring in an aqueous solution of the following. (a) (b) (c) (d) (e) (f) (g)

NaNO2 and HNO2 NaHCO3 and H2CO3 HClO2 and KClO2 RbF and HF KCl and HCl H3PO4 and KH2PO4 (CH3)2NH2NO3 and (CH3)2NH

answers: see end of problem set

2.

Write the appropriate equilibrium constant expressions (Ka, etc.) for each of the the solutions above.

3.

Determine the pH of a mixture of 0.100 M solutions of the two components in part (a) above (Ka = 4.5 x 10-4) Answer: pH = 3.35

4.

Calculate the pH of the following: (a) 0.35 M NaIO4/0.20M HIO4 (Ka = 5.6x10-9) (b) 0.48 M NaIO/0.78 M HIO (Ka = 2.0x10-11) (c) 0.40 M NH3 (Kb = 1.8x10-5)/0.55 M NH4Cl.

Answers: (a) 8.49 (b) 10.49 (c) 9.12

5.

HNO2 is a weak acid whose Ka= 4.5x10-4. How many grams of NaNO2 would have to be added to 250 mL of a 0.44M HNO2 acid solution to give a buffer of pH = 4.00. (34.2 g)

6.

A 0.625-gram sample of an unknown weak acid (call it HA for short) is dissolved in enough water to make 25.0 mL of solution. This weak acid solution is then titrated with 0.100 M NaOH and 45.0 mL of the NaOH solution is required to reach the equivalence point. Using a pH meter, the pH of the solution at the equivalence point is found to be 8.25. (a) (b)

Determine the molecular mass of the unknown acid. Determine the pKa value of the unknown acid.

138.9 g/mole 3.69

7.

The pKb value for ammonia (NH3) is 4.74. Calculate the mass (in grams) of solid ammonium chloride (NH4Cl, formula mass = 53.5 g/mole) that must be added to 300 mL of 0.25 M NH3 solution to make a buffer solution with a pH equal to 10.00. 0.73 g

8.

A 50 mL portion of a 0.10 M solution of a weak acid, HA, whose Ka equals 1.0 x 10-4, is diluted to 250 mL and then titrated with 0.10 M NaOH. Calculate pH after the following mL of the NaOH solution have been added. (a) 0 (b) 5 (c) 25 (d) 50 (e) 60 (a) 2.85 (b) 3.05

9.

(c) 4.00

(d) 8.11

(e) 11.51

Fifty mL of 0.1 M NH3 is titrated with 0.10 M HCl. The Kb of the weak base NH3 is 1.8 x 10-5. Calculate the pH after the following mL of the HCl solution have been added. (a) 0 (b) 10 (c) 25 (d) 50 (e) 55 (a) 11.13 ( b) 9.86 (c) 9.26 (d) 5.28 (e) 2.32

2

10. Shown on the next page is the titration curve for the titration of 50.00 mL of a solution of the monoprotic acid, HA, with 0.01 M NaOH. What is the molarity of the acid? Is the acid a weak acid or a strong one? If it is a weak acid, what is the value of Ka? (0.02 M, weak, Ka=2.8 x 10-4) Titration of 50.00 mL of HA with 0.01M NaOH

12 10

pH

8 6 4 2 0

20

40

60

80

100

120

140

160

mL NaOH

11. Shown below is the titration curve for the titration of 50.0 mL of a solution of the weak ammoniatype base, methyl amine, CH3NH2, with 0.01M HCl. What is the molarity of the base? What is the value of Kb for 12 methyl amine? (0.02 M answer given; Kb = 3.2x10-4)

10

pH

8 6 4 2 0

50 100 mL 0.01M HCl

150

12. Consider reactions involving the following acids and bases. State in each case whether the pH at the equivalence point would be below 7, at 7, or above 7.

3 (a) acid = HClO4 base = KOH (b) acid = HNO3 base = NH3 (c) acid = HF base = LiOH (a) 7

(b) 7

13. A 36.0 mL sample of a 0.40 M HNO3 solution was titrated with 0.48 M KOH. Calculate the pH after the following volumes of KOH have been added: (a) 0 mL (b) 10.0 mL (c) 20.0 mL (30.0 mL (e) 40.0 mL ( f) 50.0 mL. (a) 0.40

(b) 0.68

(c) 1.07

d) 7.00

(e) 12.80

(f) 13.05

14. Calculate the pH’s of the following: (a) the solution formed by mixing 20.0 mL of a 0.30M HCl solution with 12 mL of a 0.25 M NaOH. (b) the solution formed by mixing 18.0 mL of a 0.46M HC6H7O6 solution with 10.0 mL of a –5 0.35M NaOH (Ka for HC6H7O6 = 8.0x10 ) (c) the solution formed by mixing 15.0 mL of a 0.47M HF with 20.0 mL of a 0.20M KOH solution (Ka for HF = 7.1x10–4) (d) the solution formed by mixing 30.0 mL of a 0.20 M HCHO2 solution with 22.0 mL of a 0.40 M KOH solution (Ka for HCHO2 = 1.7x10–4) (e) the solution formed by dissolving 2.85 g of KF in 120 mL of a 0.30 M HF solution (Ka for HF = 7.1x10–4) (f) the end point of the titration of 58.0 mL of a 0.48 M HCHO2 with 0.40 M KOH. (a) 1.03

(b) 3.96

(c) 3.27

(d) 12.73

(e) 3.28

(f) 8.55

15. A 40.0 mL sample of a 0.40 M HNO2 solution was titrated with 0.80 M KOH. Given that HNO2 is a weak acid whose Ka = 4.5x10-4, calculate the pH after the following volumes of KOH have been added: (a) 0 mL; (b) 5.0 mL; (c) 10.0 mL; (d) 20.0 mL; (e) 30.0 mL; (f) 40.0 mL (a) 1.87

(b) 2.87

(c) 3.35

(d) 8.39

(e) 13.06

(f) 13.30

16. Drawn below is a plot of pH vs mL of 0.40 M NaOH for the titration of 50.0 mL of a weak diprotic acid. What is the concentration of the acid? What are the values of Ka1 and Ka2 for the acid? (Answer: 0.80M, Ka1 = 5.0x10-3, Ka2 = 6.3x10-8) 14

12

pH

10 8 6 4 2 0 0

50

100 150 200 250 300 350 400 mL 0.4M NaOH

17. The Ksp of Pb(IO3)2 at 25 °C = 3.2 x 10-13. What is the solubility of Pb(IO3)2 in moles per liter and in grams per liter? (4.3 x 10-5 mol/L ; g/L = 0.024)

4

18. Calculate the Ksp of the following compounds. The solubilities are given in moles per liter. (a) Mg(OH)2 = 1.3 x 10-4 M (8.8 x 10-12) (b) Ag2C2O4 = 1.4 x 10-4 M (1.1 x 10-11) 19. Calculate the Ksp of the following salts. Solubilities are given in grams/liter. (a) BaCrO4 = 2.3 x 10--3 g/l (b) CaF2 = 2.7 x 10-2 g/l

(8.3 x 10-11) (1.7 x 10-10)

20. What [SO42-] must be exceeded to produce a RaSO4 precipitate in 500 mL of a solution containing 0.00010 moles of Ra2+? (Ksp = 4 x 10-11). (2 x 10-7M) 21. A solution contains an Mg2+ concentration of 0.001 mole/liter. Will Mg(OH)2 (Ksp = 8.9 x 10-12) precipitate if the OH- concentration of the solution is (a) 10-5 mole/liter? (b) 10-4 mole/liter? (a) Q = 1x10-13, no (b) Q = 1x10-11, yes] 22. How many grams of NaOH are required to start the precipitation of Mg(OH)2 in 100 mL of a solution which contains 0.1 g of MgCl2? (1.16 x 10-4 g) 23. The value of Ksp for PbCl2 is 1.6 x 10-5. Will a precipitate of PbCl2 form when the following solutions are mixed? (a) 100 mL of 0.01 M Pb(NO3)2 and 100 mL of 0.002 M NaCl. (b) 10 mL of 0.01 M Pb(NO3)2 and 30 mL of 0.2 M NaCl. (c) 10 mL of 0.01 M Pb(NO3)2 and 20 mL of 0.06 M NaCl. (d) 10 mL of 0.01 M Pb(NO3)2 and 20 mL of 0.06 M CaCl2. (a) no, Q = 5x10–9 (b) yes, Q= 5.6x10–5 (c) no, Q= 5.3x10–6 (d) yes, Q=2.1x10–5 24. Silver chromate, Ag2CrO4, is an "insoluble" substance with a Ksp value of 1.2 x 10-12. Silver ion forms a stable complex ion with cyanide ion that has the formula Ag(CN)2- and a formation constant (Kf) of 5.3 x 1018. Calculate the molar solubility of Ag2CrO4 in each of the following solutions. Write balanced chemical equations for any important equilibrium reactions that are occurring. (a) in water 6.69 x 10-5 M (b) in 2.00 M Na2CrO4 3.87 x 10-7 M (c) 2.00 M NaCN 0.50 M 25. A saturated solution of Mg3(PO4)2 has a concentration of 0.939 mg per liter. Calculate the solubility product constant for Mg3(PO4)2. 6.27 x 10-26 26. Cadmium carbonate, CdCO3, is a sparingly soluble salt whose Ksp= 2.5x10-14. It is also known that Cd2+ forms the [Cd(NH3)4]2+ complex ion that has a formation constant, Kf = 1.0x107. (a) Calculate the molar solubility of CdCO3 in 6.0M NH3. (0.018 mol/L) (b) What NH3 concentration would be required to dissolve 4.00 mmol of CdCO3 in one liter of the solution? (2.83 M)

5 Answers to #1 (a) (b) (c) (d) (e) (f) (g)

HNO2 H+ + NO2H2CO3 HCO3- + H+ HClO2 H+ + ClO2F- + H+ HF not a buffer: the only equilibrium in this neutral solutions is dissociation of water, i.e., H+ + OHH2O H3PO4 H+ + H2PO4(CH3)2NH2+ H+ + (CH3)2NH