Endothermic process

This article is about the physical effect. For self-maintained thermal homeostasis, see Endotherm.

An endothermic process is any process which requires or absorbs energy from its surroundings, usually in the form of heat. 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" (ἔνδον) meaning "within", and the root "therm" (θερμ-), 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, "gives out" energy in the form of heat. Thus in each term (endothermic & exothermic) the prefix refers to where heat goes as the reaction occurs, though in reality it only refers to where the energy goes, without necessarily being in the form of heat.

Details

All chemical reactions involve both the breaking of existing and the making of new chemical bonds. A reaction to break a bond always requires the input of energy and so such a process is always endothermic. When atoms come together to form new chemical bonds, the electrostatic forces bringing them together leave the bond with a large excess of energy (usually in the form of vibrations and rotations). If that energy is not dissipated, the new bond would quickly break apart again. Instead, the new bond can shed its excess energy - by radiation, by transfer to other motions in the molecule, or to other molecules through collisions - and then become a stable new bond. Shedding this excess energy is the exothermicity that leaves the molecular system. Whether a given overall reaction is exothermic or endothermic is determined by the relative contribution of these bond breaking endothermic steps and new bond stabilizing exothermic steps.

The concept is frequently applied in physical sciences to, for example, chemical reactions, where thermal energy (heat) is converted to chemical bond energy.

Endothermic (and exothermic) analysis only accounts for the enthalpy change (∆H) of a reaction. The full energy analysis of a reaction is the Gibbs free energy (∆G), which includes an entropy (∆S) and temperature term in addition to the enthalpy. A reaction will be a spontaneous process at a certain temperature if the products have a lower Gibbs free energy (an exergonic reaction) even if the enthalpy of the products is higher. Entropy and enthalpy are different terms, so the change in entropic energy can overcome an opposite change in enthalpic energy and make an endothermic reaction favorable.

Examples

• Photosynthesis
• Melting of ice
• Evaporating liquid water
• Sublimation of carbon dioxide (dry ice)
• Cracking of alkanes
• Thermal decomposition reactions
• Electrolytic decomposition of sodium chloride into sodium hydroxide and hydrogen chloride
• Dissolving ammonium chloride in water
• Nucleosynthesis of elements heavier than nickel in stellar cores
• High-energy neutrons can produce tritium from lithium-7 in an endothermic reaction, consuming 2.466 MeV. This was discovered when the 1954 Castle Bravo nuclear test produced an unexpectedly high yield.
• Nuclear fusion of elements heavier than iron in supernovae
• Dissolving together Barium hydroxide and ammonium chloride
• Dissolving together citric acid and baking soda

External links

• Endothermic Definition – MSDS Hyper-Glossary