Equilibrium and Redox Reactions
Domain 5: Equilibrium and Redox Reactions
This domain tests dynamic equilibrium, acid–base calculations, solubility, and the electron-balance method for redox equations.
Core rules
- At equilibrium, forward and reverse rates are equal and concentrations are constant, not equal. Kc uses equilibrium concentrations; Kp uses partial pressures; Kp=Kc(RT)^Δn_g for gases.
- Le Chatelier: changing concentration, pressure, or temperature shifts the position; only temperature changes K. A catalyst changes neither position nor K.
- Weak acid: Ka=[H+][A−]/[HA]; for a weak monoprotic acid, [H+]≈√(KaC) when dissociation is small; pH=−log[H+]. Buffers use Henderson–Hasselbalch: pH=pKa+log([A−]/[HA]).
- Ksp is an equilibrium constant; common-ion addition lowers solubility. Precipitation occurs when the ion product exceeds Ksp.
- Redox: oxidation number tracks electron ownership; balance atoms, then oxygen with H2O, hydrogen with H+, and charge with electrons. In basic medium, add OH− to neutralise H+.
Worked example
For 0.100 M acetic acid, Ka=1.8 x 10^−5. Let x=[H+]. x^2/(0.100−x)=1.8 x 10^−5; x≈√(1.8 x 10^−6)=1.34 x 10^−3 M. pH=−log(1.34 x 10^−3)=2.87. The approximation is valid because x is about 1.3% of C.
Traps
- Treating strong and weak acids identically; HCl is fully ionised, acetic acid is not.
- Forgetting that Kp is unitless in the thermodynamic convention but has units in some expressions.
- Thinking pressure changes shift equilibria with equal gas moles.
- Mis-assigning oxidation numbers in peroxides (O −1) and hydrides (H −1).
- Balancing atoms but not charge in redox.