Empirical and molecular formulae — Chemistry, 14–17 years
How measurements reveal the simplest ratio of atoms in a substance, and how that ratio can be expanded to its real molecular formula.
The simplest atom ratio
A chemical formula can show the smallest whole-number ratio of atoms in a compound. That is its empirical formula: glucose, for example, is CH₂O in its simplest ratio, although each glucose molecule is actually C₆H₁₂O₆. The molecular formula tells how many complete groups of that ratio one molecule contains.
Why reduce the ratio?
Chemists often measure a sample’s mass, not the atoms one by one. Dividing each element’s amount by its atomic mass gives a comparison of how many particles are present, but the first numbers may not be whole numbers. Reducing them to the smallest whole-number ratio creates a formula that can be compared with other substances and tested against molecular mass.
Worked example: an oxide
A compound contains 2.4 g of magnesium and 1.6 g of oxygen. Step 1: divide by relative atomic masses: 2.4 ÷ 24 = 0.10 mol Mg and 1.6 ÷ 16 = 0.10 mol O. Step 2: divide both by the smallest value, giving 1:1. The empirical formula is MgO; if its molecular mass were 80, it would be twice MgO, or Mg₂O₂.
The tempting decimal
A common mistake is to write a decimal subscript directly, such as C₁.₅H₃O. The decimal is reasonable because measured masses rarely divide neatly, but atoms occur in whole particles. Multiply every ratio by the same small number—2 in this case—to turn 1.5 into 3, giving C₃H₆O₂. Then check that the numbers cannot all be reduced again.
From analysis to identification
Laboratories use measured percentages to work out the formula of an unknown material. This helps identify minerals, medicines and substances found in environmental samples. The formula is not a full picture of a molecule’s shape, but it gives a firm starting point: which elements are present and in what simplest ratio.
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