Subarea V - Thermodynamics, Kinetics, and Equilibrium
Subarea V covers energy changes, reaction rates, and equilibrium, including calorimetry, Hess’s law, Gibbs free energy, collision theory, Le Chatelier’s principle, and equilibrium constants.
Thermodynamics
Enthalpy change ΔH is heat at constant pressure. Exothermic reactions have ΔH < 0; endothermic have ΔH > 0. Calorimetry: q = mcΔT, where c is specific heat capacity (water = 4.18 J/g·°C). Hess’s law: ΔH for a reaction is the sum of stepwise ΔH values, reversing sign when a step is reversed. Gibbs free energy: ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG < 0.
Kinetics
Rates increase with concentration, surface area, and temperature, and with a catalyst. Collision theory requires sufficient energy (activation energy) and proper orientation. A catalyst lowers activation energy but does not change ΔH or the equilibrium position. Rate laws are determined experimentally, not from coefficients alone.
Equilibrium
At equilibrium, forward and reverse rates are equal and concentrations are constant. For aA + bB ⇌ cC + dD, K = [C]^c[D]^d / ([A]^a[B]^b). K changes with temperature. Le Chatelier’s principle: a stress shifts the system to counteract it. Increasing pressure favors the side with fewer gas moles; adding a catalyst changes neither K nor position.
Worked example
How much heat raises 50.0 g of water from 25.0 °C to 35.0 °C? q = mcΔT = (50.0 g)(4.18 J/g·°C)(10.0 °C) = 2,090 J = 2.09 kJ. Reasoning: ΔT is final minus initial, the mass is in grams, and the specific heat units cancel to joules. If the water absorbed the heat, the reaction releasing it would have ΔH negative in the calorimeter.
Traps
Do not use Celsius in gas-law calculations; do use Celsius differences in ΔT. Do not reverse the sign of ΔH incorrectly in Hess’s law. Do not assume a catalyst shifts equilibrium. Do not calculate K from initial concentrations. Do not ignore phases when writing K expressions.