reaction

Collision theory explains how chemical reactions occur and why rates of reaction differ. For a reaction to occur, particles must collide. If the collision causes a chemical change it is referred to as a fruitful collision.

The rate of reaction decreases with time while the concentration of product increases.

In chemistry, the coming together, or interaction, of two or more atoms, ions, or molecules with the result that a chemical change takes place and a new substance is formed, with a different chemical composition. The nature of the reaction is described by a chemical equation. For example, in the chemical reaction that occurs when magnesium burns in oxygen, a new substance, magnesium oxide is made:

2Mg_(s) + O_2(g) → 2MgO_(s)

Chemical equations show the reactants and products of a chemical reaction by using chemical symbols and formulae. State symbols and the energy symbol (ΔH) can be used to show whether reactants and products are solids (s), liquids (l), gases (g), or are in solution (aq); and whether energy has been released or absorbed during the reaction.

In addition reactions two or more compounds react together to form one compound. For example hydrogen chloride reacts with ethene to give chloroethane:

CH2=CH2 + HCl → CH3CH2Cl

Chain reactions produce very fast, exothermic reactions, as in the formation of flames and explosions.

In displacement reactions a less reactive element is replaced in a compound by a more reactive one. For example, the addition of powdered zinc to a solution of copper(II) sulphate displaces copper metal, which can be detected by its characteristic colour.

In endothermic reactions there is a physical or chemical change where energy is absorbed by the reactants from the surroundings. The energy absorbed is represented by the symbol +ΔH. Photosynthesis is an example.

An exothermic reaction is a chemical reaction during which heat is given out (see energy of reaction).

In heterogeneous reactions there is an interface between the different components or reactants. Examples of heterogeneous reactions are those between a gas and a solid or between two immiscible liquids.

In homogeneous reactions there is no interface between the components. The term applies to all reactions where only gases are involved or where all the components are in solution.

In photochemical reactions light is produced or light initiates the reaction. Light can initiate reactions by exciting atoms or molecules and making them more reactive: the light energy becomes converted to chemical energy.

A redox reaction is where one reactant is reduced and the other reactant oxidized. The reaction can only occur if both reactants are present and each changes simultaneously. For example, hydrogen reduces copper(II) oxide to copper while it is itself oxidized to water:

CuO + H2 → Cu + H2O

A reversible reaction proceeds in both directions at the same time, as the product decomposes back into reactants as it is being produced. Such reactions do not run to completion, provided that no substance leaves the system. The reaction reaches chemical equilibrium. The manufacture of ammonia from hydrogen and nitrogen is an example:

N2 + 3H2 ⇌ 2NH3

The term is also applied to those reactions that can be made to go in the opposite direction by changing the conditions, but these run to completion because some of the substances escape from the reaction. An example is the decomposition of calcium hydrogencarbonate on heating:

Ca(HCO3)2 → CaCO3 + CO2 + H2O

In substitution reactions one atom or functional group in an organic molecule is replaced by another.

The reactions of metals are determined by the position of the metal in the reactivity series of metals.