Unit 4 Acid-Base and Solubility Equilibria
4.5 Exercises
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Section 4.5 Exercises
- Explain how to choose the appropriate acid-base indicator for the titration of a weak base with a strong acid.
- Why can we ignore the contribution of water to the concentrations of H3O+ in the solutions of following acids:
0.0092 M HClO, a weak acid
0.0810 M HCN, a weak acid
but not the contribution of water to the concentration of OH−?
- Calculate the pH at the following points in a titration of 40 mL (0.040 L) of 0.100 M barbituric acid (Ka = 9.8 × 10−5) with 0.100 M KOH.
(a) no KOH added
(b) 20 mL of KOH solution added
(c) 39 mL of KOH solution added
(d) 40 mL of KOH solution added
(e) 41 mL of KOH solution added
Solutions
- At the equivalence point in the titration of a weak base with a strong acid, the resulting solution is slightly acidic due to the presence of the conjugate acid. Thus, pick an indicator that changes color in the acidic range and brackets the pH at the equivalence point. Methyl orange is a good example.
- In an acid solution, the only source of OH− ions is water. We use Kw to calculate the concentration. If the contribution from water was neglected, the concentration of OH− would be zero.
- (a) pH = 2.50
(b) pH = 4.01
(c) pH = 5.60
(d) pH = 8.35
(e) pH = 11.08
