Ionic Equilibria
Ionic Equilibria covers acid-base strength, pH, buffers, hydrolysis and solubility product. The exam expects numerical pH and Ksp problems plus qualitative reasoning about common-ion effects.
Core rules
- Kw = [H+][OH-] = 1.0 x 10^-14 at 25 °C; pH + pOH = 14 only at this temperature.
- Weak acid: Ka = x^2/(C - x) ≈ x^2/C, so [H+] = sqrt(Ka C).
- Buffer: pH = pKa + log([salt]/[acid]) (Henderson-Hasselbalch).
- Ksp for AgCl = [Ag+][Cl-]; a common ion lowers solubility.
- Salt hydrolysis: cation of weak base gives acidic solution; anion of weak acid gives basic solution.
Worked example
Find pH of 0.010 M acetic acid, Ka = 1.8 x 10^-5. [H+] = sqrt(1.8 x 10^-5 x 0.010) = sqrt(1.8 x 10^-7) = 4.24 x 10^-4 M. pH = -log(4.24 x 10^-4) = 3.37. Reasoning: the approximation C - x ≈ C is valid because x is under 5% of C; check this before using it.
Traps
Using 14 for pH + pOH at temperatures other than 25 °C, ignoring the common-ion effect in Ksp, and writing the buffer ratio upside down. Strong acids are assumed fully dissociated, but concentrated weak acids may need the quadratic form.