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Exothermic reaction

An exothermic reaction releases energy, typically as heat, into its surroundings. This article covers definitions, thermodynamic basis, examples, measurement, history, applications, and safety notes.

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Overview

An exothermic reaction is a process in which a chemical transformation releases energy to its surroundings, most commonly as heat. In thermodynamic terms the system loses energy and the enthalpy change (ΔH) for the reaction is negative. Exothermic events are contrasted with an endothermic reaction, which absorbs heat from the environment.

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Characteristics and thermodynamic basis

At the molecular level, exothermic behaviour arises when the total energy of products is lower than that of reactants; the excess energy is liberated as kinetic energy of molecules, light, or thermal energy transferred into the surroundings. Calorimetry measures this released heat and provides values such as standard enthalpy of reaction. Observationally, a mixture undergoing an exothermic reaction typically warms and may release fumes or light.

Common examples and uses

Many familiar processes are exothermic. Examples include:

  • Combustion of fuels (wood, coal, gasoline) where chemical bonds react with oxygen to release heat and often light.
  • Neutralization of acids and bases, which releases heat when protons and hydroxide ions form water.
  • Oxidation reactions such as metal rusting (slow) or thermite reactions (rapid and intensely exothermic).
  • Phase changes like condensation and freezing, which give off latent heat to the surroundings.
  • Biological respiration, where glucose oxidation provides energy to organisms.

Measurement, history and context

Quantifying heat release has been central to chemistry since the development of calorimeters in the 18th and 19th centuries and the formulation of thermochemistry principles, such as Hess's law. Modern techniques determine enthalpy changes and reaction kinetics to predict whether a reaction will be self-heating, explosive, or manageable in industrial settings. For fundamental definitions, see entries on chemical reaction and energy transfer.

Importance and safety considerations

Exothermic reactions are exploited for heating applications (hand warmers, combustion engines, industrial furnaces) and energy generation. They also pose hazards: uncontrolled heat release can cause thermal runaway, fires, or explosions. Proper temperature control, heat dissipation, and safe reactor design are essential when scaling or storing reactive materials.

Distinctions and notable facts

Not every reaction that feels warm is strongly exothermic; the rate of heat release and specific heat capacity of the surroundings determine perceived temperature change. Some processes combine exothermic and endothermic steps; the net effect and practical consequences depend on both heat magnitude and reaction rate.

Examples

Typical exothermic reactions are:

  • Fire (combustion)
  • Setting (= hardening) of concrete.
  • After brief heating, iron and sulphur react to form iron sulphide under the generation of light and heat.

The mixing of substances (heat of mixing) or the adsorption and absorption of substances, for example on activated carbon or zeolites, is also often exothermic, although to a much lesser extent.

Exothermic and exergonic reactions

At first, it seems obvious to assume that exothermic reactions are precisely those reactions that take place voluntarily, and that the more heat is released, the more violent they become. In many cases, chemical reactions do indeed behave in this way. This experience led to the formulation of the principle of Thomsen and Berthelot in the early years of thermochemistry. This empirical - but not strictly valid - rule states that if reactants are brought together under isobaric and isothermal conditions so that a chemical reaction can proceed, then the resulting new equilibrium state is characterized by the fact that the process leading to it releases more heat than any other possible process. In other words, of all possible processes, the most exothermic is realized. The principle is also equivalent to saying that the realized process should make the enthalpy difference {\displaystyle H_{\mathrm {Anfang} }-H_{\mathrm {Ende} }}as large as possible and thus the resulting enthalpy {\displaystyle H_{\mathrm {Ende} }}small as possible.

The existence of voluntarily occurring endothermic reactions (for example, an evaporating liquid) shows, of course, that this principle cannot claim general validity. The actual criterion is: Those reactions take place voluntarily which lead to an increase in the total entropy of the system and its environment. Under isobaric and isothermal conditions, this criterion of total entropy maximization is equivalent to minimizing the Gibbs energy of the system. A reaction that reduces the Gibbs energy of the system is called an exergonic reaction. The distinction between voluntary and involuntary reactions is equivalent to the distinction between exergonic and endergonic reactions.

An example of a chemical reaction that is endothermic but nevertheless voluntary is the decomposition of dinitrogen trioxide into nitrogen monoxide and nitrogen dioxide:

{\displaystyle \mathrm {N_{2}O_{3}\longrightarrow \ NO\ +\ NO_{2}\ ;\quad } \Delta _{\mathrm {R} }H=+39{,}7\;\mathrm {kJ/mol} ,\quad \Delta _{\mathrm {R} }G=-1{,}6\;\mathrm {kJ/mol} }

The enthalpy of reaction Δ \Delta _{{\mathrm {R}}}Hof this decay is positive, so the reaction is endothermic. The Gibbs reaction energy Δ \Delta _{\mathrm {R} }G, however, is negative, so the reaction is exergonic.

The change in Gibbs energy G=H-TS is under isothermal conditions

{\displaystyle \Delta G=\Delta H-T\ \Delta S}.

At small temperatures, Δ {\displaystyle \Delta G\approx \Delta H}and minimizing the Gibbs energy G is approximately equivalent to minimizing the enthalpy Hof the system. In this case, the exergonic reactions are usually also exothermic reactions, and the principle of Thomsen and Berthelot predicts the equilibrium states approximately correctly by considering the enthalpy change. Even at higher temperatures (such as room temperature), the principle remains approximately correct, since the temperature dependences of Δ \Delta Gand Δ \Delta Hare similar at temperatures that are not too high (as can be shown by considering the Third Law), and the similarity of Δ \Delta Gand Δ \Delta Htherefore preserved with temperature increase over a larger temperature range.

However, if a reaction is accompanied by a sufficiently large entropy increase Δ \Delta S (as in the aforementioned cases of evaporating liquid or the decay of dinitrogen trioxide), then the term may {\displaystyle -T\ \Delta S}predominate and the reaction may proceed voluntarily (exergonic, Δ {\displaystyle \Delta G<0}), although its enthalpy increases in the process (endothermic, Δ \Delta H>0), i.e. the reaction "runs uphill" in terms of enthalpy.

The thermite reaction (the reduction of iron(III) oxide by aluminium) is strongly exothermic and proceeds very violently.
The thermite reaction (the reduction of iron(III) oxide by aluminium) is strongly exothermic and proceeds very violently.

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