Thermochemistry

Heat (q)

energy transferred between objects due to a temperature difference; flows hot → cold.

Temperature

average kinetic energy of the particles in a substance.

System / surroundings

the system is the reaction being studied; the surroundings are everything else.

Exothermic

process that releases heat; q and ΔH are negative.

Endothermic

process that absorbs heat; q and ΔH are positive.

First Law of Thermodynamics

energy is conserved; ΔE = q + w.

Specific heat (c)

energy needed to raise 1 g of a substance by 1 °C; water = 4.18 J/(g·°C).

Calorimetry

measuring heat flow, based on heat lost = heat gained (q = mcΔT).

Enthalpy (ΔH)

heat content of a system at constant pressure; ΔH = H_products − H_reactants.

Bond enthalpy

energy to break a bond; ΔH ≈ (bonds broken) − (bonds formed).

Hess's Law

ΔH of an overall reaction equals the sum of the ΔH of its steps (enthalpy is a state function).

Standard enthalpy of formation (ΔH°f)

heat to form 1 mol of a compound from its elements in standard states; = 0 for a pure element.

State function

a property that depends only on the current state, not the path taken (e.g., H, S, G).

Entropy (S)

measure of the dispersal of energy and matter; increases with gases, dissolving, and higher temperature.

Gibbs free energy (ΔG)

predicts spontaneity; ΔG = ΔH − TΔS. Negative ΔG = spontaneous.

Spontaneous (thermodynamically favored)

a reaction that proceeds forward on its own (ΔG < 0); note this says nothing about speed.