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 stoichiometry 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.