Acids, Bases, and Salt Solutions
This domain tests Brønsted–Lowry theory, the distinction between strong and weak acids, pH and pKa calculations, buffer behaviour, titration curves and the hydrolysis that decides whether a salt solution is acidic, neutral or alkaline.
Theory
- An acid donates a proton, a base accepts one; a conjugate pair differs by a single H+.
- Strong acids (HCl, HNO3, H2SO4, HClO4) dissociate essentially completely; weak acids such as ethanoic acid (Ka ≈ 1.8 × 10−5) dissociate only partially.
- Kw = [H+][OH−] = 1.0 × 10−14 at 298 K, so pH + pOH = 14 at that temperature.
Calculations
- Strong acid: [H+] = c, so pH = −log c. A tenfold dilution raises the pH by 1.
- Weak acid: Ka = [H+]²/(c − [H+]), simplified to [H+]²/c when dissociation is under about 5%.
- Buffer: pH = pKa + log([A−]/[HA]); at half neutralisation pH = pKa.
- pKa = −log Ka; a smaller pKa means a stronger acid.
Salts and indicators
- Strong acid + strong base gives a neutral salt; weak acid + strong base gives an alkaline solution because the anion hydrolyses; strong acid + weak base gives an acidic solution.
- The indicator's pH range must bracket the equivalence pH: phenolphthalein 8.2–10.0 for weak acid–strong base titrations, methyl orange 3.1–4.4 for strong acid–weak base.
Worked example
0.100 M ethanoic acid with Ka = 1.8 × 10−5: x²/0.100 = 1.8 × 10−5, so x = [H+] = 1.34 × 10−3 M and pH = −log(1.34 × 10−3) = 2.87. Only about 1.3% is dissociated, confirming weak behaviour and justifying the approximation.
Traps that cost marks
- Using pH + pOH = 14 at temperatures other than 298 K.
- Assuming the equivalence point of a weak acid–strong base titration is pH 7; it is above 7.
- Inverting the buffer ratio, or forgetting that dilution hardly changes the pH of a buffer.
- Ignoring the contribution of water autoionisation for very dilute or very weak acids.
- Confusing strength (degree of dissociation) with concentration (amount per dm3).