Chapter 17

1. (a) Consider the equilibrium: HA (aq) « H+ (aq) + A- (aq) In terms of LeChatelier’s principle, explain the effect of the presence of a salt of A- on the ionization of HA. (b) Give an example of a salt that can decrease the ionization of NH3 in solution. (c) Explain why the ionization of a weak acid is suppressed by the presence of its conjugate base.

2. Describe the effect on pH (increase, decrease, or no change) that results from each of the following additions: (a) ammonia to a solution of HCl; (b) ammonium chloride to a solution of HCl; (c) sodium cyanide, NaCN, to a solution of HBr; (d) pyridinium nitrate, C5H5NHNO3, to a solution of pyridine, C5H5N.

3. Calculate the pH of the following solutions: (a) 0.100 M in sodium formate, NaCHO2, and 0.180 M in formic acid, HCHO2; (b) 0.070 M in pyridine, C5H5N, and 0.0500 M in pyridinium chloride, C5H5NHCl.

4. How many milliliters of 0.0750 M HCl are needed to titrate each of the following solutions to the equivalence point: (a) 50.0 mL of 1.50 M NaOH; (b) 43.5 mL of 0.0320 M KOH; (c) 27.0 mL of 0.0600 M Ca(OH)2?

5. Three 40.0 mL samples of different KOH solutions are titrated with 0.0750 M HNO3 solution. The volumes of acid needed to reach the equivalence point for each sample are (a) 39.1 mL; (b) 20.7 mL; (c) 48.5 mL. What was the original pH of each of the KOH solutions?

6. A 30.0 mL sample of 0.200 M KOH is titrated with 0.150 M HClO4 solution. Calculate the pH after the following volumes of acid have been added: (a) 30.0 mL; (b) 39.5 mL; (c) 39.9 mL; (d) 40.0 mL; (e) 40.1 mL.

7. Assume that 30.0 mL of a 0.10 M solution of a weak base B that accepts one proton is titrated with a 0.10 M solution of a monoprotic strong acid HX. (a) How many moles of HX have been added at the equivalence point? (b) What is the predominant form of B at the equivalence point? (c) What factor determines the pH at the equivalence point? (d) Which indicator, phenolphthalein or methyl red, is likely to be the better choice for this titration?

8. Consider the titration of 60.0 mL of 0.100 M NH3 with 0.150 M HCl. Calculate the ph after the following volumes of titrant have been added: (a) 0.0 mL; (b) 20.0 mL; (c) 39.5 mL; (d) 40.0 mL; (e) 40.5 mL; (f) 60.0 mL.

9. Calculate the pH at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 M NaOH: (a) hydrobromic acid, HBr; (b) lactic acid, HC3H5O3; (c) sodium hydrogen chromate, NaHCrO4.

10. What factors determine (a) the pH and (b) the capacity of a buffer solution?

11. A buffer is prepared by adding 20.0 g of acetic acid and 20.0 g of sodium acetate to enough water to form 2.00 L of solution. (a) Determine the pH of the buffer. (b) Write the complete ionic equation for the reaction that occurs when a few drops of hydrochloric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of sodium hydroxide solution are added to the buffer.

12. (a) Write the net ionic equation for the reaction that occurs when a solution of NaOH is mixed with a solution of methylammonium bromide, CH3NH3Br. (b) Calculate the equilibrium constant for this reaction. (c) Calculate the concentrations of Na+, Br-, H+, CH3NH3+, and CH3NH2 when 0.250 L of 0.300 M NaOH is mixed with 0.400 L of 0.100 M CH3NH3Br.

13. How many grams of sodium lactate, NaC3H5O3, should be added to 1.00 L of 0.0.0150 M lactic acid, HC3H5O3, to form a buffer solution with a pH of 3.90? Assume that no volume change occurs when the NaC3H5O3is added.

14. A buffer solution contains 0.120 mol of propionic acid, HC3H5O2, and 0.105 mol of sodium propionate, NaC3H5O2, and has a total volume of 1.00 L. (a) What is the pH of this buffer? (b) What is he pH after the addition of 10.0 mL of 1.50 M HCl solution? (c) What is the pH after the addition of 50.0 mL of 0.400 M NaOH solution?

15.A phosphate buffer, consisting of H2PO4- and HPO42-, helps control the pH of physiological fluids. Many carbonated soft drinks also use this buffer system. What is the pH of a soft drink in which the major buffer ingredients are 6.5 g of NaH2PO4 and 8.0 g of Na2HPO4 per 355 mL of solution?

16. (a) Explain the difference between solubility and solubility product constant. (b) Write the expression for the solubility product constant for each of the following ionic compounds: CuS, NiC2O4, Ag2SO4, Co(OH)3, and Fe3(AsO4)2.

17. (a) The molar solubility of PbBr2 at 250C is 1.0 x 10-2 mol/L. Calculate Ksp. (b) If 0.490 g of AgIO3 dissolves per liter of solution, calculate the solubility product. (c) Using the appropriate Ksp value from appendix D, calculate the solubility of Cu(OH)2 in grams per liter of solution.

18. Calculate the solubility of CaF2in grams per liter in (a) pure water, (b) 0.15 M KF solution, (c) 0.080 M Ca(NO3)2 solution.

19. Calculate the solubility of Mn(OH)2 in grams per liter (a) at pH = 7.0, (b) at pH = 9.5, (c) at pH = 11.8.

20. What is the ratio of [Ca2+] to [Fe2+] in a lake in which the water is in equilibrium with deposits of both CaCO3 and FeCO3?

21. (a) The Ksp for cerium iodate, Ce(IO3)3, is 3.2 x 10 -10. Calculate the molar solubility of Ce(IO3)3 in pure water. (b) What concentration of NaIO3 in solution would be necessary to reduce the Ce3+ concentration in a saturated solution of Ce(IO3)3 by a factor of 10 below that calculated in part (a)?

22. (a) Will Co(OH)2 precipitate from solution if the pH of a 0.020 M solution of Co(NO3)2 is adjusted to 8.5? (b) Will AgIO3precipitate when 100 mL of 0.010 M AgNO3 is mixed with 10 mL of 0.015 M NaIO3? (Ksp of AgIO3is 3.0 x 10-8.)

23. Suppose a 50 mL sample of a solution is to be tested for Cl- ion by addition of 1 drop (0.2 mL) of 0.10 M AgNO3. What is the minimum number of grams of Cl- that must be present in order for AgCl (s) to form?

24. Which of the following salts will be substantially more soluble in acidic solution than in pure water: (a) ZnCO3; (b) LaF3; (c) BiI3; (d) AgCN; (e) Ba3(PO4)2?

25. A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2+ and 0.010 M in Sr2+. (a) What concentration of SO42- is necessary to begin precipitation? (Neglect volume changes; BaSO4: Ksp = 1.1 x 10-10; SrSO4: Ksp = 2.8 x 10-7). (b) Which cation precipitates first? (c) What is the concentration of SO42- when the second cation begins to precipitate?

26. Calculate the solubility of ZnS (in grams per liter) in a 0.10 M solution of H2S in which the pH is 2.40.

27. Complete and balance each of the following reactions. In each case indicate which species in the reaction can be considered a Lewis base and which a Lewis acid:

(a) Cu2+(aq) + CN- (aq) ®

(b) AgCl (s) + S2O32- (aq) ®

(c) Cu(OH)2 (s) + NH3 (aq) ®

28. To what final concentration of NH3 must a solution be adjusted to just dissolve 0.020 mol of NiC2O4 in 1 L of solution? (Hint: You can neglect the hydrolysis of C2O42-, because the solution will be quite basic.)

29. Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2S, and Kf 1.1 x 105 for AgCl2-, calculate the equilibrium constant for the following reaction:

Ag2S (s) + 4 Cl- (aq) + 2 H+ (aq) « 2 AgCl2 - (aq) + H2S (aq)