Chemical reaction

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Editor-In-Chief: C. Michael Gibson, M.S., M.D. [1]

A chemical reaction is a process that results in the interconversion of chemical substances.[1] The substance or substances initially involved in a chemical reaction are called reactants. Chemical reactions are usually characterized by a chemical change, and they yield one or more products which are, in general, different from the reactants. Classically, chemical reactions encompass changes that strictly involve the motion of electrons in the forming and breaking of chemical bonds, although the general concept of a chemical reaction, in particular the notion of a chemical equation, is applicable to transformations of elementary particles, as well as nuclear reactions.

Different chemical reactions are used in combinations in chemical synthesis in order to get a desired product. In biochemistry, series of chemical reactions catalyzed by enzymes form metabolic pathways, by which syntheses and decompositions ordinarily impossible in conditions within a cell are performed.

Reaction types

The large diversity of chemical reactions and approaches to their study results in the existence of several concurring, often overlapping, ways of classifying them. Below are examples of widely used terms for describing common kinds of reactions:

N2 + 3 H2 → 2 NH3
2 H2O → 2 H2 + O2
2 Na(s) + 2 HCl(aq) → 2 NaCl(aq) + H2(g)
  • Metathesis or Double displacement reaction, in which two compounds exchange ions or bonds to form different compounds:
NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
  • Precipitation (chemistry) Reactions where species in solution combine to form a solid product (precipitate). A typical example would be the reaction of methatesis described above.
  • Acid-base reactions, broadly characterized as reactions between an acid and a base, can have different definitions depending on the acid-base concept employed. Some of the most common are:
  • Arrhenius definition: Acids dissociate in water releasing H3O+ ions; bases dissociate in water releasing OH- ions.
  • Brønsted-Lowry definition: Acids are proton (H+) donors; bases are proton acceptors. Includes the Arrhenius definition.
  • Lewis definiton: Acids are electron-pair acceptors; bases are electron-pair donors. Includes the Brønsted-Lowry definition.
  • Redox reactions, in which changes in oxidation numbers of atoms in involved species occur. Those reactions can often be interpreted as transferences of electrons between different molecular sites or species. A typical example of redox rection is:
2 S2O32−(aq) + I2(aq) → S4O62−(aq) + 2 I(aq)

In which I2 is reduced to I- and S2O32- (thiosulfate anion) is oxidized to S4O62-.

  • Combustion, a kind of redox reaction in which any combustible substance combines with an oxidizing element, usually oxygen, to generate heat and form oxidized products. The term combustion is used usually only large-scale oxidation of whole molecules, i.e. a controlled oxidation of a single functional group is not combustion.
C10H8+ 12 O2 → 10 CO2 + 4 H2O
CH2S + 6 F2CF4 + 2 HF + SF6

Reactions can also be classified according to their mechanism, some typical examples being:

Some branches of chemistry include any detectable changes in chemical conformation in the reaction types, while others consider these changes merely as physical properties of a compound.

The collision of more than two particles into the ordered structure necessary to perform chemical transformations is extremely unlikely; which is why ternary reactions in practice are not observed. A chemical reaction may require three or more reactants, but the process can generally be best described as a stepwise series of elementary reactions.

Chemical kinetics

Main article: Chemical kinetics

The rate of a chemical reaction is a measure of how the concentration or pressure of the involved substances changes with time. Analysis of reaction rates is important for several applications, such as in chemical engineering or in chemical equilibrium study. Rates of reaction depends basically on:

  • Reactant concentrations, which usually make the reaction happen at a faster rate if raised through increased collisions per unit time,
  • Surface area available for contact between the reactants, in particular solid ones in heterogeneous systems. Larger surface area leads to higher reaction rates.
  • Pressure, by increasing the pressure, you decrease the volume between molecules. This will increase the frequency of collisions of molecules.
  • Activation energy, which is defined as the amount of energy required to make the reaction start and carry on spontaneously. Higher activation energy implies that the reactants need more energy to start than a reaction with a lower activation energy.
  • Temperature, which hastens reactions if raised, since higher temperature increases the energy of the molecules, creating more collisions per unit time,
  • The presence or absence of a catalyst. Catalysts are substances which change the pathway (mechanism) of a reaction which in turn increases the speed of a reaction by lowering the activation energy needed for the reaction to take place. A catalyst is not destroyed or changed during a reaction, so it can be used again.
  • For some reactions, the presence of electromagnetic radiation, most notably ultra violet, is needed to promote the breaking of bonds to start the reaction. This is particularly true for reactions involving radicals.

Reaction rates are related to the concentrations of substances involved in reactions, as quantified by the rate law of each reaction. Note that some reactions have rates that are independent of reactant concentrations. These are called zero order reactions.

See also


  1. International Union of Pure and Applied Chemistry. "chemical reaction". Compendium of Chemical Terminology Internet edition.

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