chemical reaction

chemical reaction

reaction, chemical: see chemical reaction.
chemical reaction, process by which one or more substances may be transformed into one or more new substances. Energy is released or is absorbed, but no loss in total molecular weight occurs. When, for example, water is decomposed, its molecules, each of which consists of one atom of oxygen and two of hydrogen, are broken down; the hydrogen atoms then combine in pairs to form hydrogen molecules and the oxygen atoms to form oxygen molecules. In a chemical reaction, substances lose their characteristic properties. Water, for example, a liquid which neither burns nor supports combustion, is decomposed to yield flammable hydrogen and combustion-supporting oxygen. In some reactions heat is given off (exothermic reactions), and in others heat is absorbed (endothermic reactions). Furthermore, the new substances formed differ from the original substances in the energy they contain. Chemical reactions are classified according to the kind of change that takes place. When a compound, which consists of two or more elements or groups of elements, is broken down into its constituents, the reaction is called simple decomposition. When two compounds react with one another to form two new compounds, the reaction is called double decomposition. In so-called replacement reactions the place of one of the elements in a compound is taken by another element reacting with the compound. When elements combine to form a compound, the reaction is termed chemical combination. Oxidation and reduction reactions are extremely important. Reversible reactions are those in which the chemical change taking place may be paralleled by another change back to the original substances. The rates at which chemical reactions proceed depend upon various factors, e.g., upon temperature, pressure, and the concentration of the substances involved and, sometimes, upon the use of a chemical called a catalyst. In some chemical reactions, such as that of photographic film, light is an important factor. The changes taking place in a chemical reaction are represented by a chemical equation. An element's activity, i.e., its tendency to enter into compounds, varies from one element to another.

Any chemical process in which substances are changed into different ones, with different properties, as distinct from changing position or form (phase). Chemical reactions involve the rupture or rearrangement of the bonds holding atoms together (see bonding), never atomic nuclei. The total mass and number of atoms of all reactants equals those of all products, and energy is almost always consumed or liberated (see heat of reaction). The speed of reactions varies (see reaction rate). Understanding their mechanisms lets chemists alter reaction conditions to optimize the rate or the amount of a given product; the reversibility of the reaction and the presence of competing reactions and intermediate products complicate these studies. Reactions can be syntheses, decompositions, or rearrangements, or they can be additions, eliminations, or substitutions. Examples include oxidation-reduction, polymerization, ionization (see ion), combustion (burning), hydrolysis, and acid-base reactions.

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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:


  • 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 definition: 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. An example of a redox reaction 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 usually used for 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

Organic reactions encompass a wide assortment of reactions involving compounds which have carbon as the main element in their molecular structure. The reactions in which an organic compound may take part are largely defined by its functional groups. Defined in opposition to inorganic reactions. Reactions can also be classified according to their mechanism, some typical examples being:

*Reactions of ions, e.g. disproportionation of hypochlorite
*Reactions with reactive ionic intermediates, e.g. reactions of enolates
*Radical reactions, e.g. combustion at high temperature
*Reactions of carbenes

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:

  • 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


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