Chemical Thermodynamics — AI Study Guide

Mastering Chemical Thermodynamics

Thermodynamics is the study of energy and its transformations in chemical and physical processes. The first law of thermodynamics states that energy cannot be created or destroyed — only converted between forms. In chemistry, this is expressed as ΔE = q + w, where ΔE is the change in internal energy, q is heat, and w is work. At constant pressure (typical laboratory conditions), heat change equals enthalpy change (ΔH).

Enthalpy (H) is the heat content of a system at constant pressure. Enthalpy changes (ΔH) are negative for exothermic reactions (release heat) and positive for endothermic reactions (absorb heat). Hess's law states that ΔH for a reaction is independent of the pathway — allowing calculation of ΔH for any reaction from standard enthalpies of formation. Standard enthalpies of formation are defined relative to elemental forms in standard states.

Entropy (S) measures the disorder or dispersal of energy in a system. The second law of thermodynamics states that the entropy of the universe always increases in spontaneous processes. Entropy increases with: temperature, volume, number of moles of gas, mixing, and structural complexity. Calculating entropy changes (ΔS) requires considering both system and surroundings; the universe's total entropy must increase for a process to be spontaneous.

Gibbs free energy (G) combines enthalpy and entropy into a single criterion for spontaneity at constant temperature and pressure: ΔG = ΔH - TΔS. A process is spontaneous when ΔG < 0, non-spontaneous when ΔG > 0, and at equilibrium when ΔG = 0. Understanding how temperature affects spontaneity (by changing the -TΔS term) explains why some reactions are spontaneous only at high or low temperatures. ΔG is also related to the equilibrium constant K.

Frequently Asked Questions: Chemical Thermodynamics

What is the difference between exothermic and endothermic reactions?

Exothermic reactions release heat to the surroundings (ΔH < 0, negative) — the products have lower enthalpy than the reactants. Examples: combustion, neutralization, most oxidation reactions. Endothermic reactions absorb heat from the surroundings (ΔH > 0, positive) — the products have higher enthalpy than the reactants. Examples: dissolving ammonium nitrate, photosynthesis, cooking. Exothermic reactions feel warm; endothermic reactions feel cool.

When is a reaction spontaneous?

Spontaneity is determined by ΔG = ΔH - TΔS: the reaction is spontaneous when ΔG < 0. Four combinations: (1) ΔH < 0 and ΔS > 0: always spontaneous (favored by both); (2) ΔH > 0 and ΔS < 0: never spontaneous; (3) ΔH < 0 and ΔS < 0: spontaneous at low temperature (enthalpy dominates); (4) ΔH > 0 and ΔS > 0: spontaneous at high temperature (entropy dominates). Spontaneous does not mean fast — kinetics determines rate, thermodynamics determines direction.

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