Acids, Bases and Salts
Acids, bases and salts tests Brønsted–Lowry theory, pH and Ka calculations, titration curves, indicator choice and the hydrolysis of salts.
Core rules
- A Brønsted–Lowry acid is a proton donor; a base is a proton acceptor. Strong acids dissociate almost completely in water; weak acids dissociate only partially and set up an equilibrium.
- pH = −log[H+]; [H+] = 10−pH. At 298 K, Kw = [H+][OH−] = 1.0 × 10−14.
- For a weak acid HA, Ka = [H+][A−]/[HA]. When dissociation is small, [H+] ≈ √(Ka × [HA]).
- In titration calculations, use the balanced equation to convert moles of titrant to moles of analyte, then divide by the analyte volume in dm3.
- Indicator choice depends on the vertical section of the titration curve: phenolphthalein for strong acid–strong base and weak acid–strong base; methyl orange for strong acid–weak base.
- Salts of weak acids with strong bases hydrolyse to give alkaline solutions; salts of weak bases with strong acids give acidic solutions.
Worked example
Calculate the pH of 0.10 mol dm−3 ethanoic acid (Ka = 1.8 × 10−5). [H+] ≈ √(1.8 × 10−5 × 0.10) = √(1.8 × 10−6) = 1.34 × 10−3 mol dm−3. pH = −log(1.34 × 10−3) = 2.87. The acid is weak, so the pH is higher than the pH of 0.10 mol dm−3 HCl, which would be 1.00.
Traps
Using pH = −log(acid concentration) for a weak acid; forgetting to convert cm3 to dm3; choosing an indicator whose range lies outside the steep section; assuming every salt dissolves to give a neutral solution; ignoring the temperature dependence of Kw when it is not 298 K.