Empirical Formula Calculator
Enter up to three elements and their mass percentages to calculate the compound's empirical formula.
Last updated: September 26, 2026
Frequently asked questions
What is an empirical formula?
The simplest whole-number ratio of atoms in a compound - for example, glucose's molecular formula is C₆H₁₂O₆, but its empirical formula is CH₂O, the reduced ratio.
How do I calculate an empirical formula from mass percentages?
Divide each element's mass percentage by its atomic weight to get relative moles, then divide all the mole values by the smallest one to get the simplest ratio - round to whole numbers if they're close (like 1.98 → 2).
What's the empirical formula for a compound that's 40% C, 6.7% H, 53.3% O?
CH₂O - moles: C = 40/12.011 ≈ 3.33, H = 6.7/1.008 ≈ 6.65, O = 53.3/15.999 ≈ 3.33. Dividing all by the smallest (3.33) gives roughly C₁H₂O₁, so CH₂O.
Why might my mole ratios not come out to exact whole numbers?
Rounding in the mass percentages themselves, or a true ratio that isn't 1:1:1 (like 2:3), can leave you with values like 1.5 or 2.33 instead of whole numbers - if that happens, try multiplying all the ratios by a small whole number (2 or 3) to see if they land on whole numbers.
How is the empirical formula different from the molecular formula?
The empirical formula shows only the simplest whole-number ratio of atoms, while the molecular formula shows the actual number of atoms in one molecule - the molecular formula is always a whole-number multiple of the empirical formula (e.g., C₆H₁₂O₆ is 6× the empirical formula CH₂O).
What if I only know the mass percentages of two elements and they don't add up to 100%?
The missing percentage is typically oxygen or another unmeasured element - subtract your known percentages from 100% to find it, then include that value as the third element for a more accurate formula.
Do I need to know the actual masses of my sample, or do percentages work fine?
Percentages work fine, since ratios are all that matter - mass percentages behave exactly like grams in a 100 g sample, so the relative mole ratios come out the same whether you use percentages or actual measured masses.