Stoichiometry

Stoichiometry /stoi-kE-'a-m&-trE/ (from Greek stoicheion meaning element or principle, and Middle English metrie meaning to measure) refers to the relative number of atoms of various elements found in a chemical substance and is often useful in characterizing a chemical reaction.

It rests upon the law of definite proportions (i.e., the law of constant composition) and the law of multiple proportions.

Stoichiometry is often used to balance chemical equations. For example, the two diatomic gases hydrogen and oxygen can combine to form a liquid, water, in an exothermic reaction, as described by Equation 1.

  (1) H2 + O2 --> H2O

Equation 1 does not depict the proper stoichiometry of the reaction—that is, it does not reflect the relative proportions of the reactants and products.

 (2) 2H2 + O2 --> 2H2O 

Equation 2 does have proper stochiometry and is therefore said to be a "balanced" equation, depicting the same number of atoms of each type on each side of the equation.

The term stoichiometry is also often used for the molar proportions of elements in stoichiometric compounds. For example, the stoichiometry of hydrogen and oxygen in H₂O is 2:1. In stoichiometric compounds, the molar proportions are whole numbers (that is what the law of multiple proportions is about).

Compounds for which the molar proportions are not whole numbers are called nonstoichiometric compounds. Such compounds can be produced by sputtering in a plasma. They are not in chemical equilibrium.

Solids that actually are a mixture of very small crystallites of compounds of different stoichiometry also have been loosely called nonstoichiometric compounds. This is incorrect and probably due to the difficulty in observing the very small crystallites. If a solid was misinterpreted as homogeneous, it was consequently misinterpreted as nonstoichiometric.