Acids, Bases, Salts and Buffer Solutions
Definitions and strength
Brønsted-Lowry: an acid donates a proton, a base accepts one. Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) ionize completely; weak acids ionize partially and are described by Ka. For a conjugate pair, Ka x Kb = Kw = 1.0x10^-14 at 25 °C, so a stronger acid has a weaker conjugate base. pH = -log[H3O+] and pH + pOH = 14 at 25 °C.
Salts and buffers
Salt hydrolysis decides the pH of a salt solution: a salt of a strong acid and strong base is neutral; strong acid plus weak base is acidic (NH4Cl); weak acid plus strong base is basic (CH3COONa). A buffer is a weak acid with its conjugate base (or weak base with conjugate acid) in comparable amounts; it resists pH change because added acid is consumed by the base component and added base by the acid component. Henderson-Hasselbalch: pH = pKa + log([base]/[acid]). Buffer capacity is greatest when pH is near pKa.
Titration curves
Strong acid with strong base: equivalence at pH 7. Weak acid with strong base: the equivalence point is basic (about 8-9 for acetate), with buffering before it. The half-equivalence point of a weak acid has pH = pKa.
Worked example
Find the pH of 0.10 M CH3COOH, Ka = 1.8x10^-5. Let x = [H3O+]. Ka = x^2/(0.10 - x) which gives approximately x^2/0.10, so x^2 = 1.8x10^-6 and x = 1.34x10^-3 M. pH = -log(1.34x10^-3) = 2.87. The approximation holds because ionization is only about 1.3%.
Traps
Using pH = -log(concentration) for a weak acid. Forgetting that a salt solution need not be neutral. Writing the buffer ratio upside down. Assuming the equivalence point of a weak acid titration is at pH 7.