Enthalpy (H) is used to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the sum of a system's internal energy (U) and the mathematical product of its pressure (P) and volume (V):
H = U + PV
Enthalpy is also a state function. Enthalpy values for specific substances cannot be measured directly; only enthalpy changes for chemical or physical processes can be determined. For processes that take place at constant pressure (a common condition for many chemical and physical changes), the enthalpy change (ΔH) is:
ΔH = ΔU + PΔV
The heat given off when you operate a Bunsen burner is equal to the enthalpy change of the methane combustion reaction that takes place, since it occurs at the essentially constant pressure of the atmosphere. A negative value of an enthalpy change, ΔH < 0, indicates an exothermic reaction; a positive value, ΔH > 0, indicates an endothermic reaction.
The ΔH and the change in free energy, called delta G (∆G), is related by the following equation, which is known as Gibbs Helmholtz equation;
ΔG = ΔH − TΔS
We generally calculate standard pH, temperature, and pressure conditions at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 bar), respectively. Note that cellular conditions vary considerably from these standard conditions, and so standard calculated ∆G values for biological reactions will be different inside the cell.
This text is adapted from Openstax, Biology 2e, Section 6.3: The Laws of Thermodynamics and Openstax, Chemistry 2e, Section 5.3: Enthalpy.
