Core Chemistry 2: Thermodynamics 1st Year Recap

A comprehensive set of vocabulary flashcards covering the fundamental laws, state functions, and spontaneity criteria from the Core Chemistry 2 Thermodynamics lecture.

Zeroth Law

States that if two systems are each in thermal equilibrium with a third system, they are in thermal equilibrium with each other; systems with different temperatures exchange heat, $q$, until they reach thermal equilibrium.

1st Law of Thermodynamics

The Conservation of Energy principle which states that internal energy, $U$, is the sum of a system’s potential and kinetic energies, and an increase in internal energy ($\Delta U$) for a closed system is achieved by heating it (q > 0) and doing work on it (w > 0).

Internal Energy ($U$)

The sum of a system’s potential and kinetic energies.

Entropy ($S$)

A state function and measure of the dispersal of energy, thermodynamically defined by $dS = \frac{\delta q_{rev}}{T}$. Entropy increases with heat input and accessible volume.

2nd Law of Thermodynamics

States that the entropy of an isolated system increases in a spontaneous process (\Delta S > 0).

Microstate ($\Omega$)

A specific way of distributing the available energy amongst translational, rotational, vibrational, and electronic energy levels.

Clausius Inequality

The relationship defined by $dS \geq \frac{\delta q}{T}$.

3rd Law of Thermodynamics

States that the entropy of all perfect crystalline substances is $0$ at $T = 0\,K$, consistent with the idea that there is no energy available to disperse.

Zero point energy

Energy that the Uncertainty Principle shows is irremovable and therefore cannot be dispersed.

Enthalpy ($H$)

Defined as $H = U + pV$; it is more useful than internal energy at constant pressure because the change in enthalpy is equal to the flow of heat ($dH = dq_p$).

Heat Capacity at Constant Volume ($C_V$)

The partial differential of internal energy with respect to temperature at constant volume: $C_V = \left( \frac{\partial U}{\partial T} \right)_V$.

Heat Capacity at Constant Pressure ($C_p$)

The partial differential of enthalpy with respect to temperature at constant pressure: $C_p = \left( \frac{\partial H}{\partial T} \right)_p$.

Gibbs Free Energy ($G$)

Defined as $G = H - TS$; it represents the amount of energy ‘free’ to do non-expansion work at constant $T$ and $p$.

Fundamental Equation of Thermodynamics

The equation $dG = V dp - S dT$, which shows how Gibbs free energy varies with temperature and pressure for a closed system.

Helmholtz Free Energy ($A$)

Defined by the relation $A = U - TS$; it is useful in statistical thermodynamics and must decrease ($dA \leq 0$) for a spontaneous process at constant $T$ and $V$.

Isolated System

A system that has no exchange of energy or matter with the surroundings.

Closed System

A system that exchanges energy with the surroundings but not matter.

Open System

A system that exchanges both energy and matter with the surroundings.

Spontaneity Criteria (Constant $T, p$)

For a process to occur spontaneously under these conditions, the change in Gibbs Free Energy must be less than or equal to zero ($\Delta G = \Delta H - T \Delta S \leq 0$).