Chemical reaction

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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.

• Isomerisation, in which a chemical compound undergoes a structural rearrangement without any change in its net atomic composition; see stereoisomerism
• Direct combination or synthesis, in which 2 or more chemical elements or compounds unite to form a more complex product:

N₂ + 3 H₂ → 2 NH₃

• Chemical decomposition or analysis, in which a compound is decomposed into smaller compounds or elements:

2 H₂O → 2 H₂ + O₂

• Single displacement or substitution, characterized by an element being displaced out of a compound by a more reactive element:

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)

• 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. 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 S₂O₃^2- (thiosulfate anion) is oxidized to S₄O₆^2-.

• 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 F₂ → CF₄ + 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

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

• List of reactions
• Organic reaction
• Inorganic chemical reaction
• List of important publications in chemistry
• Stoichiometry
• Stoichiometric coefficient
• Transition state theory
• Gas stoichiometry
• Thermodynamic databases for pure substances
• Autocatalytic reaction
• Coefficients