The method
- Treat percentages as grams (imagine a 100 g sample).
- Convert each element's grams to moles (÷ atomic mass).
- Divide every mole value by the smallest one.
- If ratios aren't whole, multiply all of them by 2, 3, 4… until they are.
Worked example
A compound is 69.9% iron and 30.1% oxygen. Find its empirical formula.
- Fe: 69.9 ÷ 55.85 = 1.25 mol • O: 30.1 ÷ 16.00 = 1.88 mol
- Divide by 1.25 → Fe: 1.00, O: 1.50
- 1.5 isn't whole — multiply both by 2 → Fe: 2, O: 3
Answer: Fe₂O₃ — iron(III) oxide, the main component of rust.
Empirical vs molecular formulas
The empirical formula is the simplest ratio; the molecular formula counts real atoms per molecule. Glucose's molecular formula C₆H₁₂O₆ reduces to the empirical CH₂O. To go from empirical to molecular you need the compound's molar mass: divide it by the empirical formula's mass to find the multiplier (180 ÷ 30 = 6 for glucose).
⚠ Common mistake: Rounding too early is the classic trap: a ratio of 1.33 is ⅓ short of whole — multiply by 3 (→ 4), don't round down to 1. This calculator warns when ratios don't settle cleanly, which usually signals measurement or rounding error in the input data.