chemical energy

(physical chemistry) Energy of a chemical compound which, by the law of conservation of energy, must undergo a change equal and opposite to the change of heat energy in a reaction; the rearrangement of the atoms in reacting compounds to produce new compounds causes a change in chemical energy.

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A useful but obsolescent term for the energy available from elements and compounds when they react, as in a combustion reaction. In precise terminology, there is no such thing as chemical energy, since all energy is stored in matter as either kinetic energy or potential energy. See also Combustion; Energy.

When a chemical reaction takes place, the atoms of the reactants change their bonding pattern and become products. The breaking of bonds in the reactants requires energy, and the formation of bonds in the products releases energy. The net change in energy is commonly referred to as chemical energy. To be more precise, when a reaction takes place, there is an overall change in the enthalpy H of the system as bonds are broken and new bonds are formed. This change in enthalpy is denoted ΔH. Under standard conditions [a pressure of 1 bar (100 kilopascals) and all substances pure], the change is noted ΔH° and called the standard enthalpy of reaction. Provided the pressure is constant, the standard enthalpy can be identified with the energy released as heat (when ΔH° < 0) or gained as heat (ΔH° > 0) when the reaction takes place. Reactions for which ΔH° < 0 are classified as exothermic; those for which ΔH° > 0 are classified as endothermic. All combustions are exothermic, the released heat being used either to provide warmth or to raise the temperature of a working fluid in an engine of some kind. There are very few common endothermic reactions; one example is the dissolution of ammonium nitrate in water (a process utilized in medical cold packs). See also Chemical equilibrium; Chemical thermodynamics; Enthalpy.

The “chemical energy” available from a typical fuel (that is, the enthalpy change accompanying the combustion of the fuel, when carbon-hydrogen bonds are replaced by stronger carbon-oxygen and hydrogen-oxygen bonds) is commonly reported as either the specific enthalpy or the enthalpy density. The specific enthalpy is the standard enthalpy of combustion divided by the mass of the reactant. The enthalpy density is the standard enthalpy of combustion divided by the volume of the reactant. The former is of primary concern when mass is an important consideration, as in raising a rocket into orbit. The latter is of primary concern when storage space is a limitation. The specific enthalpy of hydrogen gas is relatively high (142 kilojoules/g), but its enthalpy density is low (13 kJ/L). The values for octane, a compound representative of gasoline, are 48 kJ/g and 38 MJ/L, respectively (note the change in units). The high enthalpy density of octane means that a gasoline tank need not be large to store a lot of “chemical energy.”

Energy liberated by a chemical reaction or absorbed in the formation of a chemical compound.