Chemical Equilibrium and Acid-Base Chemistry
What you must understand
Write equilibrium expressions from balanced equations, remembering that pure solids and liquids are omitted. Use the reaction quotient Q to predict the direction of shift and apply Le Chatelier's principle to changes in concentration, pressure, volume, and temperature. Know that catalysts do not shift equilibrium. For acids and bases, use Bronsted-Lowry and Lewis definitions, conjugate pairs, and the relationships Ka times Kb = Kw and pH + pOH = 14 at 25 degrees Celsius. Handle weak-acid calculations with ICE tables and approximations, buffers with the Henderson-Hasselbalch equation, and titration curves including equivalence-point pH.
Traps to watch for
- Including solids or pure liquids in K expressions.
- Using strong-acid assumptions for weak acids, or forgetting the square-root step in Ka problems.
- Confusing the equivalence point with the neutral point; weak acid-strong base titrations end basic.
- Applying Le Chatelier to catalysts or to changes that do not affect the equilibrium composition.
- Using pH + pOH = 14 at temperatures other than 25 degrees Celsius.
Worked example
Find the pH of 0.10 M acetic acid with Ka = 1.8 x 10^-5. Set up the ICE table: x squared divided by (0.10 - x) equals Ka. Assume x is small relative to 0.10, so x squared equals 1.8 x 10^-6 and x is about 1.3 x 10^-3 M. Then pH is -log(1.3 x 10^-3), about 2.87. Check that x is under five percent of 0.10 M, which it is, so the approximation holds.
Rule of thumb
For buffers, pH = pKa + log([base]/[acid]); the buffer works best when the ratio is near one.