Normality is a measure of concentration of a chemical species in solution. It is defined as the number of equivalents of solute per volume of solution in liters.

This definition is dependent upon the type of chemical reaction under consideration. For acid-base reactions, one equivalent of acid is the quantity that supplies one mole of H⁺ ions; whereas one equivalent of base is the amount that reacts with one mole of H⁺ ions. In oxidation-reduction reactions, an equivalent is the quantity of substance that either gains (oxidant) or loses (reductant) one mole of electrons. Though not as commonly used, an equivalent in precipitation and/or complexation reactions is the amount of the substance that provides or reacts with one mole of reacting cation if it is univalent, one-half of a mole if it is divalent, one-third of a mole if it is trivalent, etc. At the equivalence point of a titrimetric analysis, irrespective of the reaction type, one equivalent of analyte will have reacted with one equivalent of titrant.

Products formed as the result of acid-base and/or oxidation-reduction reactions often depend upon the experimental conditions. For example, when H₃PO₄ reacts as an acid to form the HPO₄^2- ion it loses two H⁺ ions. Thus, one mole of H₃PO₄ (98.0 grams) will be two equivalents. If one mole of H₃PO₄ is dissolved in sufficient water to form 1 liter of solution, its concentration is given as 1 Molar (abbreviated as 1 M) or 2 Normal (abbreviated as 2 N). If only a single H⁺ is removed, the concentration is only 1 N. The normality of the H₃PO₄ depends upon the number of H⁺ ions that are removed. Similarly, in oxidation- reduction reactions, the calculated normality depends upon the number of electrons gained or lost be one formula unit of the substance (i.e., Normality = number of electrons transferred x Molarity). Concentrations expressed as normalities are meaningless unless accompanied by knowledge of how the solution is to be used and the chemical reaction(s) involved.