Normality from Molarity
Worked example: 0.500 M H2SO4 at 2 eq/mol → 1.00 N — press Try an example to run it live, then adjust anything.
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Normality from Molarity explained
Normality counts reactive capacity rather than molecules. One mole of sulfuric acid delivers two protons, so 0.500 M H₂SO₄ is 1.00 N — and 1.00 N of any acid neutralises 1.00 N of any base volume for volume, which is exactly why the unit survived so long in analytical labs. For redox work the equivalents are electrons: potassium permanganate in acid picks up five, making 0.0200 M KMnO₄ a 0.100 N oxidant.
The catch, and the reason IUPAC deprecated normality decades ago, is that the equivalent count depends on the reaction, not the substance. Phosphoric acid is 3 N when fully neutralised to phosphate but only 1 N in a titration stopped at the first endpoint, so a bottle labelled "1 N H₃PO₄" is ambiguous without knowing the intended reaction. You will still meet it constantly in water-treatment and clinical chemistry, where hardness, alkalinity, and electrolyte balances are quoted in equivalents per litre or milliequivalents per litre — blood sodium at 140 mEq/L being the most familiar example.
Normality from Molarity formula
- = Normality in eq/L (M)
- = Molarity (M)
- = Equivalents per mole
Missing one of these? Work it out first, then come back
- Normality in eq/L — Molarity (C = n/V), pH from Hydrogen Ion Concentration
- Molarity — Molarity (C = n/V), pH from Hydrogen Ion Concentration
- Equivalents per mole — Equivalent Weight from Molar Mass and Valence