In chemistry, the equivalent concentration or normality of a solution is defined as the molar concentration divided by an equivalence factor or -factor :
Normality is defined as the number of gram or mole equivalents of solute present in one liter of solution. The SI unit of normality is equivalents per liter (Eq/L).
where is normality, is the mass of solute in grams, is the equivalent weight of solute, and is the volume of the entire solution in liters.
There are three common types of chemical reaction where normality is used as a measure of reactive species in solution:
Normal concentration of an ionic solution is also related to conductivity (electrolytic) through the use of equivalent conductivity.
Although losing favor in the medical industry, reporting of serum concentrations in units of "eq/L" (= 1 N) or "meq/L" (= 0.001 N) still occurs.
Normality can be used for acid-base titrations. For example, sulfuric acid (H2SO4) is a diprotic acid. Since only 0.5 mol of H2SO4 are needed to neutralize 1 mol of OH−, the equivalence factor is:
feq(H2SO4) = 0.5
If the concentration of a sulfuric acid solution is c(H2SO4) = 1 mol/L, then its normality is 2 N. It can also be called a "2 normal" solution.
Similarly, for a solution with c(H3PO4) = 1 mol/L, the normality is 3 N because phosphoric acid contains 3 acidic H atoms.
The normality of a solution depends on the equivalence factor eq for a particular reaction, which presents two possible sources of ambiguity - namely, eq depends on the choice of reaction as well as which chemical species of the reaction is being discussed (e.g., acid/base species, redox species, precipitating salts, isotopes exchanged, etc.). That is to say, the same solution can possess different normalities for different reactions or potentially even the same reaction in a different context.
To avoid ambiguity, IUPAC[1] and NIST[2] discourage the use of the terms "normality" and "normal solution".