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The chemical potential assigned to a specified standard state. With dimensionless thermodynamic activity $a$, the per-particle convention is $\mu ={\mu }^{\circ }+k_{B}T\ln a$; a molar convention uses $RT$. The standard state changes both ${\mu }^{\circ }$ and the activity definition without changing observable equilibrium.

The standard enthalpy of formation is the enthalpy change for forming one mole of a specified chemical substance from its chemical elements in their reference states at the stated temperature. For sodium chloride at $298.15\mathrm{K}$, those reactant states are solid sodium and diatomic gaseous chlorine. Standard state specifies conditions; it does not by itself select a temperature.

The Gibbs energy of reaction for specified standard states. It determines the dimensionless equilibrium constant through ${\Delta }_r{G}^{\circ }=-RT\ln K$.

Molar entropy in a specified standard state. Absolute entropies of elemental reference substances are generally not zero, unlike their conventional standard formation enthalpies.

A standard state is a specified reference state used to define a species' standard chemical potential ${\mu }_i^{\circ }$. The reference convention, including temperature, must be consistent when calculating thermodynamic activities, reaction quotients, and equilibrium constants.

Thermochemical equilibrium requires every chemical reaction to have zero Gibbs free-energy change. For ideal reactants, the equilibrium constant is $K=\exp [-{\Delta }_r{G}^{\circ }/(RT)]$ after standard-state activity factors are included.

Water activity measures the chemical potential of water relative to a chosen pure-water standard state. Under ideal-vapour conditions above a flat solution, it is the equilibrium vapour pressure divided by that above pure liquid water at the same temperature. Dissolved solute usually lowers activity. Raoult's law gives the ideal dilute-solution limit; nonideal interactions require an activity model.

A weak acid establishes a chemical equilibrium between its proton-donating form and its conjugate base in water. For a neutral monoprotic acid $\mathrm{HA}$, the aqueous chemical equation isWater accepts a proton to become hydronium; ${\mathrm{A}}^{-}$ is the conjugate base. The double arrow denotes reversible exchange at chemical equilibrium, not equal populations. The acid dissociation constant and composition determine the hydronium concentration; substituting total chemical acid concentration into the pH formula is generally incorrect.

“Weak” describes the tendency to transfer a proton in a specified solvent, rather than how much chemical acid was dissolved. A weak acid can be concentrated, and a strong acid can be dilute. Even the same weak acid can have a large weak acid dissociation fraction after sufficient dilution, while its acid dissociation constant stays fixed at the same temperature and standard state conventions.^[ref]

A weak electrolyte yields only a fraction of ions from its dissolved substance under the specified conditions, as illustrated by acetic acid in water.

For a neutral weak acid, unconverted molecules coexist with hydronium and the conjugate base at chemical equilibrium. The weak acid dissociation fraction describes the partition of the dissolved inventory; the acid dissociation constant describes the chemical equilibrium at fixed solvent, temperature and standard states. A single fraction is not a permanent label for every concentration: dilution can increase that fraction while the acid dissociation constant stays fixed. Calling a substance a weak electrolyte does not mean its chemical solution must be dilute.^[ref]

The following ideal weak electrolyte dilution law makes the distinction quantitative for a monoprotic acid. Other weak electrolytes, mixed chemical solutions and additional proton-transfer steps require their own chemical equilibrium and inventory balances.