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Endothermic reaction

A chemical process that absorbs heat from its surroundings. Covers definitions, thermodynamic sign, examples, detection, applications and distinctions from related concepts.

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Overview

An endothermic reaction is a chemical change that takes in energy from its surroundings, usually in the form of heat. In thermochemical notation it is characterized by a positive enthalpy change (ΔH > 0), indicating that the products contain more chemical energy than the reactants. Endothermic processes contrast with exothermic reactions, which release heat to the surroundings. Whether an endothermic reaction occurs spontaneously depends on both enthalpy and entropy and is determined by the Gibbs free energy, not heat flow alone.

Characteristics and thermodynamic context

Key features of endothermic reactions include absorption of thermal energy, cooling of the immediate environment, and the need for a continuous energy supply for the reaction to proceed in many cases. The term describes energy transfer as heat; it does not by itself specify the reaction speed, mechanism, or whether energy is supplied as light or electricity. In many chemical equations the absorbed heat can be shown explicitly as a reactant, or the reaction enthalpy is reported as a positive value.

Common examples and applications

  • Photosynthesis in plants, where light energy is stored in chemical bonds, is commonly described as overall endothermic.
  • Dissolving certain salts such as ammonium nitrate in water absorbs heat and is the basis for disposable cold packs used in first aid.
  • Phase changes like melting and evaporation require heat input and are endothermic physical processes closely related to chemical energetics.
  • Industrial thermal decomposition or calcination reactions (for example the conversion of some carbonates to oxides and CO2) require sustained heating and are treated as endothermic steps in manufacturing.

Detection and measurement

Endothermic changes are detected by a drop in temperature of the surroundings, which can be measured with a thermometer or, more precisely, with calorimetry. In a calorimeter the heat absorbed by the reaction is measured and the enthalpy change calculated per mole of reactant. Laboratory demonstrations often show an immediate temperature decrease when a reagent dissolves or when a solid decomposes upon heating. For reversible reactions, the reverse direction may be exothermic; this reciprocity is useful in teaching about energy balance and reaction coupling.

Distinctions and notable points

Endothermic should not be confused with endergonic: endothermic describes heat absorption, while endergonic refers to a positive change in Gibbs free energy (non-spontaneity under standard conditions). A process can be endothermic but spontaneous if entropy increase compensates the enthalpy term. Understanding both heat flow and free energy is essential when assessing reaction feasibility. For accessible introductions to related terms, see general resources on endothermic reactions and on exothermic processes.

Practical importance

Recognizing endothermic steps is important in chemical engineering, environmental science, and everyday technology: from designing reactors that supply heat efficiently, to leveraging evaporative cooling in cooling systems, to employing endothermic dissolution in medical cold packs. The concept also underpins natural phenomena such as weather-driven cooling and the energy input required for biological synthesis of organic molecules.

Sequence of an endothermic reaction

As with exothermic reactions, the process of endothermic reactions also takes place in two steps. First, a certain activation energy must be applied, and then part of this energy is released again. The difference to the exothermic reaction is that this released energy is less than the activation energy and is therefore not sufficient to drive the reaction further. The reaction energy is positive. Therefore, in order not to interrupt the reaction, energy must be continuously supplied from outside during the reaction.

For an endothermic reaction to take place at all (to be exergonic), the reaction must be favored by an increase in entropy and thus have a negative free enthalpy. Endothermic reactions therefore often take place at high temperatures, since at these, according to the Gibbs-Helmholtz equation, the entropy portion of the free enthalpy is greater. This can be seen, for example, in the Boudouard equilibrium, where the endothermic reaction to carbon monoxide takes place at high temperatures.

Enthalpieprofil einer endothermen Reaktion

Legend:
 {\displaystyle \!\ H:={\text{Enthalpie}}}
 \!\ \Delta ^{{\ddagger }}H:={\text{Aktivierungsenthalpie}}
 \!\ \Delta _{{\mathrm {R}}}H:={\text{Reaktionsenthalpie}}

left: Initial state of reactants: stablemiddle
: Transition state of activated complex: unstableright
: Final state of products: metastable

For example: If you add water vapor to a coke layer, an endothermic reaction takes place: {\displaystyle \mathrm {C+H_{2}O\ {\xrightarrow {\triangle }}\ CO+H_{2}} }.

Industrial chemistry

Endothermic reactions in industrial chemistry are referred to as allothermic. Important examples are allothermal pyrolysis, in which the biomass is split by externally supplied heat, or steam reforming in the production of synthesis gas. In contrast, autothermal reactions are used in industrial terminology.

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AlegsaOnline.com Endothermic reaction

URL: https://en.alegsaonline.com/art/31401

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