Thermochemistry Study Guide
Thermochemistry is energy accounting: how much heat a reaction absorbs or releases, and whether it can happen on its own. The tools form a chain — calorimetry measures heat, formation enthalpies and Hess's law compute it, bond enthalpies estimate it, and entropy plus Gibbs free energy decide spontaneity. Exam traps cluster around three things: sign conventions, the kJ/J unit trap, and forgetting stoichiometric coefficients. This guide makes each tool explicit and each trap visible.
Key formulas
q = mcΔT, where q is heat in J, m is mass in g, c is specific heat in J/(g·°C), ΔT in °C or K
ΔH°rxn = Σ n·ΔHf°(products) − Σ m·ΔHf°(reactants), with n and m the stoichiometric coefficients
Hess's law: ΔH_reaction = Σ ΔH_steps; reversing a step changes the sign of its ΔH
ΔH ≈ Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed)
ΔG° = ΔH° − TΔS°, with T in Kelvin; ΔG° < 0 means spontaneous
ΔS > 0 when disorder increases (e.g., solid → gas, or more gas moles formed)
Topic checklist
Heat, temperature, and enthalpy
Heat (q) is energy in transit; temperature measures average kinetic energy of particles.
Endothermic processes absorb heat (surroundings cool); exothermic processes release heat (surroundings warm).
Enthalpy (H) is the heat content at constant pressure; ΔH is negative for exothermic, positive for endothermic.
Calorimetry and specific heat
q = mcΔT links heat to mass, specific heat capacity, and temperature change.
ΔT is always T_final − T_initial; watch the sign of a cooling process.
For water, c = 4.184 J/(g·°C) — memorize it, it appears constantly.
Enthalpy of formation
Standard enthalpy of formation (ΔHf°) is the enthalpy change for forming 1 mol of a compound from its elements in standard states.
Elements in their standard states have ΔHf° = 0 by definition.
ΔH°rxn = Σ n·ΔHf°(products) − Σ m·ΔHf°(reactants), with each term multiplied by its coefficient.
Hess's law
Hess's law: enthalpy is a state function, so the ΔH of a reaction equals the sum of ΔH values of its steps.
Reverse a step → flip the sign of its ΔH; scale a step → scale its ΔH the same way.
Add the manipulated steps so intermediates cancel, leaving the target reaction.
Bond enthalpies
Bond enthalpy estimates ΔH: break bonds (energy in, positive) minus form bonds (energy out, negative).
ΔH ≈ Σ(bond enthalpies broken) − Σ(bond enthalpies formed).
This is an estimate — average bond enthalpies differ slightly from real molecular values.
Entropy and free energy
Entropy (S) measures disorder; ΔS is positive when gas moles increase or solids become gases.
Gibbs free energy: ΔG° = ΔH° − TΔS°; negative ΔG° means spontaneous under standard conditions.
A reaction can be spontaneous at high T only (ΔH+, ΔS+), low T only (ΔH−, ΔS−), always, or never.
Common mistakes
Mistake: Getting the sign of ΔH backwards for exothermic vs endothermic
Fix:
Exothermic means heat leaves the system, so ΔH is negative; endothermic absorbs heat, so ΔH is positive. Say the words 'released = negative' before writing the sign.
Mistake: Mixing kilojoules and joules in ΔG = ΔH − TΔS
Fix:
Convert one to the other before adding: 1 kJ = 1000 J. If ΔH is in kJ and ΔS is in J/K, convert ΔS to kJ/K before computing ΔG = ΔH − TΔS.
Mistake: Forgetting that tabulated ΔH values are per mole of reaction as written
Fix:
ΔH°rxn from formation data is for the reaction exactly as written — including its coefficients. If the problem asks per mole of a specific reactant, divide by that coefficient.
Mistake: Dropping the stoichiometric coefficients in the formation-enthalpy sum
Fix:
Multiply every ΔHf° by its stoichiometric coefficient n before subtracting. Products minus reactants, coefficients included.
Exam tips
Run thermochemistry-practice until q = mcΔT and ΔG = ΔH − TΔS are reflexes — they are the two most-tested equations.
On Hess's law problems, write each manipulated step with its new ΔH before adding; neat bookkeeping beats cleverness.
For spontaneity questions, first write the signs of ΔH and ΔS, then decide the temperature dependence — the four cases are a classic multiple-choice set.
Always convert J to kJ (or vice versa) as your very first step in any ΔG calculation.
Pair with equilibrium-practice for ΔG° = −RT ln K crossover questions.