Empirical and molecular formulas
The empirical formula is the simplest whole-number ratio of atoms in a compound. The molecular formula is the actual number of each atom in one molecule, which is a whole-number multiple of it. Glucose, C6H12O6, has the empirical formula CH2O.
The method
- Assume 100 g of the compound, so each percentage becomes grams.
- Divide each mass by the element's atomic weight to get moles.
- Divide every mole value by the smallest one.
- If the ratios are not whole numbers, multiply them all by the smallest number that makes them whole.
A worked example
A compound is 40.0% carbon, 6.7% hydrogen and 53.3% oxygen.
- C: 40.0 ÷ 12.011 = 3.33 mol
- H: 6.7 ÷ 1.008 = 6.65 mol
- O: 53.3 ÷ 15.999 = 3.33 mol
Dividing by 3.33 gives 1 : 2 : 1, so the empirical formula is CH2O.
Ratios that are not whole numbers
Do not round 1.5 to 2. Recognize the common fractions and multiply:
- about 1.5 → multiply everything by 2
- about 1.33 or 1.67 → multiply by 3
- about 1.25 or 1.75 → multiply by 4
Iron oxide with Fe : O = 1 : 1.5 becomes Fe2O3. Small differences, such as 1.98 for 2, are measurement error and can be rounded.
From empirical to molecular formula
Divide the molar mass of the compound by the mass of the empirical formula and round to a whole number. CH2O has a mass of 30.03 g/mol; if the compound's molar mass is 180.16 g/mol, the multiple is 6 and the molecular formula is C6H12O6.
Calculators
The Empirical Formula Calculator takes percentages or masses, finds the multiplier and gives the molecular formula when you add the molar mass. The Molar Mass Calculator works out the mass of any formula.