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An endothermic process is any process with an increase in the enthalpy H (or internal energy U) of the system. In such a process, a closed system usually absorbs thermal energy from its surroundings, which is heat transfer into the system. It may be a chemical process, such as dissolving ammonium nitrate in water, or a physical process, such as the melting of ice cubes.
The term was coined by Marcellin Berthelot from the Greek roots endo-, derived from the word "endon" () meaning "within", and the root "therm" (?-), meaning "hot" or "warm" in the sense that a reaction depends on absorbing heat if it is to proceed. The opposite of an endothermic process is an exothermic process, one that releases or "gives out" energy, usually in the form of heat and sometimes as electrical energy. Thus in each term (endothermic and exothermic) the prefix refers to where heat (or electrical energy) goes as the process occurs.
Whether a reaction can occur spontaneously depends not only on the enthalpy change but also on the entropy change (?S) and absolute temperature T. If a reaction is a spontaneous process at a certain temperature, the products have a lower Gibbs free energy G = H - TS than the reactants (an exergonic reaction), even if the enthalpy of the products is higher. Thus, an endothermic process usually requires a favorable entropy increase (?S > 0) in the system that overcomes the unfavorable increase in enthalpy so that still ?G < 0. While endothermic phase transitions into more disordered states of higher entropy, e.g. melting and vaporization, are common, spontaneous chemical reactions at moderate temperatures are rarely endothermic. The enthalpy increase ?H >> 0 in a hypothetical strongly endothermic reaction usually results in ?G = ?H -T?S > 0, which means that the reaction will not occur (unless driven by electrical or photon energy). An example of an endothermic and exergonic reaction is
C6H12O6 + 6 H2O -> 12 H2 + 6 CO2, ?rH° = +627 kJ/mol, ?rG° = -31 kJ/mol