Empirical vs Molecular Formula

The empirical formula is the simplest whole-number ratio of atoms in a compound. The molecular formula is the actual number of atoms in one molecule — always a whole-number multiple of the empirical formula.

The difference in one line

Glucose has the molecular formula C₆H₁₂O₆ but the empirical formula CH₂O — the same 1 : 2 : 1 ratio, scaled down. Many compounds share an empirical formula (formaldehyde is also CH₂O) but differ in molecular formula.

Finding the empirical formula

1. Take the mass (or percent) of each element. Percent works directly — assume 100 g.
2. Convert each to moles by dividing by its atomic mass.
3. Divide every mole value by the smallest one to get a ratio.
4. If the ratio is not whole numbers, multiply all of them by a small integer until it is.

Finding the molecular formula

You also need the molar mass of the compound. Divide it by the empirical formula mass to get the multiplier n, then multiply every subscript by n:

Worked example (glucose)

A compound is 40.0% C, 6.71% H, 53.3% O, molar mass 180 g/mol.

• Moles in 100 g: C 40.0/12.01 = 3.33; H 6.71/1.008 = 6.66; O 53.3/16.00 = 3.33.
• Divide by smallest (3.33): C 1, H 2, O 1 → empirical CH₂O (mass 30.03 g/mol).
• n = 180 ÷ 30.03 ≈ 6 → molecular formula C₆H₁₂O₆.

Check empirical-formula masses quickly with the Molar Mass Calculator, and see how moles drive everything in How to Approach Stoichiometry.