Chemical thermodynamics compares measured states with agreed standard states so tabulated quantities from different substances and reactions use the same reference convention.

A superscript degree symbol, as in

$$\Delta H^\circ,\qquad \Delta G^\circ,\qquad E^\circ,$$

indicates that the relevant species are in their standard states.

For pressure-based standard states, the standard pressure is

$$\boxed{p^\circ=1,\mathrm{bar}}.$$

For a pure solid or liquid, the standard state is the pure substance at the stated temperature and standard pressure. For a gas, the standard state is referenced to ideal-gas behavior at $p^\circ$.

For species in solution, a standard state also requires a composition convention. In introductory concentration-based calculations this is commonly represented using the standard concentration

$$c^\circ=1,\mathrm{mol/L}.$$

More rigorous solution thermodynamics uses activity, a dimensionless effective composition that accounts for nonideal behavior, rather than raw concentration alone.

The temperature must still be specified. A superscript $\circ$ does not by itself mean $25,^\circ\mathrm C$ or $298.15,\mathrm K$; those temperatures are common tabulation choices, not part of the degree symbol's definition.

Standard state is also different from STP, a convention for quoting a particular temperature and pressure for gases. A standard thermodynamic quantity is a property relative to defined reference states, not a claim that every experiment actually occurs under those conditions.

Standard states let formation enthalpies, Gibbs energies, equilibrium data and electrode potentials be combined consistently. Changing the reference convention would change tabulated standard values, but properly calculated physical predictions remain unchanged when the convention is used consistently.