Thermochemistry and Thermodynamics
Enthalpy and calorimetry
Thermochemistry tracks heat flow. At constant pressure, q = ΔH. Exothermic reactions release heat (ΔH < 0); endothermic reactions absorb heat (ΔH > 0). Calorimetry uses q = mcΔT, where c is specific heat capacity. A coffee-cup calorimeter measures solution reactions; a bomb calorimeter measures combustion at constant volume. Hess's law adds enthalpy changes for steps; formation enthalpies give ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants).
Entropy and free energy
Entropy (S) measures dispersal of energy; ΔS is positive when disorder increases, especially gas production. Gibbs free energy: ΔG = ΔH − TΔS. A reaction is spontaneous at constant T and P when ΔG < 0. ΔG° = −RT ln K connects thermodynamics and equilibrium. Bond breaking absorbs energy; bond forming releases energy.
Worked example
How much heat raises 50.0 g of water from 20.0 °C to 30.0 °C? q = mcΔT = (50.0 g)(4.184 J/g·°C)(10.0 °C) = 2.09 × 10^3 J = 2.09 kJ. The water absorbs 2.09 kJ; the source released that heat. If this occurred in a calorimeter, the reaction's q would be −2.09 kJ under ideal insulation.
Traps
Dropping the sign on q or ΔH; using Celsius temperature differences incorrectly (differences are fine, but T in ΔG must be Kelvin); assuming spontaneous means fast; and confusing ΔH with ΔG.