Chemistry · Ch 3 — Thermodynamics

Summary

System, surroundings, and boundary define the thermodynamic system. The universe = system + surroundings.
Internal energy ( $U$ ) is a state function; its change ( $\Delta U$ ) depends only on the initial and final states, not the path.
First law of thermodynamics: $\Delta U = q + w$ , where $q$ is heat added to the system and $w$ is work done on the system.
Work done in a single step against a constant external pressure (irreversible): $w = -P_{\text{ext}} \Delta V$ ; for an ideal gas undergoing isothermal reversible expansion: $w = -nRT \ln\frac{V_2}{V_1}$ .
Enthalpy ( $H$ ) is defined as $H = U + PV$ ; at constant pressure, $\Delta H = q_p$ .
Heat capacity: $C = \frac{q}{\Delta T}$ ; molar heat capacity at constant volume ( $C_V$ ) and constant pressure ( $C_P$ ). For an ideal gas: $C_P - C_V = R$ .
Enthalpy change of a reaction ( $\Delta_r H$ ) is the heat absorbed or released at constant pressure. Standard enthalpy change ( $\Delta_r H^\ominus$ ) refers to 1 bar pressure and specified states.
Hess’s law: $\Delta_r H$ for a reaction is the same whether it occurs in one step or several steps — enthalpy is a state function.
Bond enthalpy: average energy required to break one mole of a specific bond in gaseous molecules. $\Delta_r H = \sum \text{(bond enthalpies of reactants)} - \sum \text{(bond enthalpies of products)}$ .
Spontaneity is determined by the second law: total entropy of the universe increases for a spontaneous process. $\Delta S_{\text{universe}} > 0$ .
Entropy ( $S$ ) is a measure of disorder. For a reversible isothermal process: $\Delta S = \frac{q_{\text{rev}}}{T}$ . …

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