Unit 4 Acid-Base and Solubility Equilibria

4.4 Exercises

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Section 4.4 Exercises

  1. Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the acid H3PO4 and a salt of its conjugate base NaH2PO4.
  2. What is [H3O+] in a solution of 0.25 M CH3CO2H and 0.030 M CH3CO2Na?
    [latex]\text{CH}_3\text{CO}_2\text{H}(aq)\;+\;\text{H}_2\text{O}(l)\;{\rightleftharpoons}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{CH}_3\text{CO}_2^{\;\;-}(aq)\;\;\;\;\;\;\;K_{\text{a}} = 1.8\;\times\;10^{-5}[/latex]
  3. What is [OH−] in a solution of 0.125 M CH3NH2 and 0.130 M CH3NH3Cl?
    [latex]\text{CH}_3\text{NH}_2(aq)\;+\;\text{H}_2\text{O}(l)\;{\rightleftharpoons}\;\text{CH}_3\text{NH}_3^{\;\;+}(aq)\;+\;\text{OH}^{-}(aq)\;\;\;\;\;\;\;K_{\text{b}} = 4.4\;\times\;10^{-4}[/latex]
  4. What concentration of NH4NO3 is required to make [OH−] = 1.0 × 10−5 M in a 0.200-M solution of NH3?
  5. What is the effect on the concentration of acetic acid, hydronium ion, and acetate ion when the following are added to an acidic buffer solution of equal concentrations of acetic acid and sodium acetate:

    (a) HCl
    (b) CH3CO2K
    (c) NaCl
    (d) KOH
    (e) CH3CO2H

  6. What will be the pH of a buffer solution prepared from 0.20 mol NH3, 0.40 mol NH4NO3, and just enough water to give 1.00 L of solution?  Hint: Use the Kb for NH3.
  7. How much solid CH3CO2Na•3H2O must be added to 0.300 L of a 0.50-M acetic acid solution to give a buffer with a pH of 5.00?  Hint: Assume a negligible change in volume as the solid is added.
  8. A buffer solution is prepared from equal volumes of 0.2000 M acetic acid and 0.6000 M sodium acetate. Use 1.800 × 10−5 as Ka for acetic acid.

    (a) What is the pH of the solution?
    (b) Is the solution acidic or basic?
    (c) What is the pH of a solution that results when 3.000 mL of 0.03400 M HCl is added to 0.2000 L of the original buffer?

Solutions

  1. Excess H3O+ is removed primarily by the reaction:
    [latex]\text{H}_3\text{O}^{+}(aq)\;+\;\text{H}_2\text{PO}_4^{\;\;-}(aq)\;{\longrightarrow}\;\text{H}_3\text{PO}_4(aq)\;+\;\text{H}_2\text{O}(l)[/latex]
    Excess base is removed by the reaction:
    [latex]\text{OH}^{-}(aq)\;+\;\text{H}_3\text{PO}_4(aq)\;{\longrightarrow}\;\text{H}_2\text{PO}_4^{\;\;-}(aq)\;+\;\text{H}_2\text{O}(l)[/latex]
  2. [H3O+] = 1.5 × 10−4 M
  3. [OH−] = 4.2 × 10−4 M
  4. [NH4NO3] = 0.36 M
  5. (a) The added HCl will increase the concentration of H3O+ slightly, which will react with CH3CO2− and produce CH3CO2H in the process. Thus, [CH3CO2−] decreases and [CH3CO2H] increases.
    (b) The added CH3CO2K will increase the concentration of [CH3CO2−] which will react with H3O+ and produce CH3CO2 H in the process. Thus, [H3O+] decreases slightly and [CH3CO2H] increases.
    (c) The added NaCl will have no effect on the concentration of the ions.
    (d) The added KOH will produce OH− ions, which will react with the H3O+, thus reducing [H3O+]. Some additional CH3CO2H will dissociate, producing [CH3CO2−] ions in the process. Thus, [CH3CO2H] decreases slightly and [CH3CO2−] increases.
    (e) The added CH3CO2H will increase its concentration, causing more of it to dissociate and producing more [CH3CO2−] and H3O+ in the process. Thus, [H3O+] increases slightly and [CH3CO2−] increases.
  6. pH = 8.95
  7. 14.37 g (0.27 mol)
  8. (a) pH = 5.222
    (b) The solution is acidic
    (c) pH = 5.220

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