Overview

An empirical formula gives the simplest whole-number ratio of atoms or components in a compound. It does not always show the actual molecule, but it shows the basic composition. For example, glucose has the molecular formula C6H12O6, while its empirical formula is CH2O because the atom ratio reduces to 1:2:1. Empirical formulas are common in general chemistry, laboratory reports, combustion analysis, and composition problems.

This calculator converts mass percentages and atomic masses into a mole ratio. Because the shared calculator shell uses numeric inputs only, the results use generic labels A, B, and C rather than element symbols. You can treat A as carbon, B as hydrogen, and C as oxygen, or replace them with any three elements by entering the correct atomic masses. If a compound has only two elements, set the third mass percentage to 0.

How to use the calculator

Enter the mass percent for each component and its atomic mass in grams per mole. Mass percentages should usually add to about 100%. Small deviations are common because experimental data are rounded. If the total is far from 100%, check whether a component is missing or whether the values are mass fractions rather than percentages.

The default values model a compound with 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen. With atomic masses of approximately 12.011, 1.008, and 15.999, those values reduce to a 1:2:1 ratio. The molecular multiplier lets you scale the empirical formula if you know the molecular formula is a whole-number multiple of the empirical formula. Without an independent molar mass, that multiplier is only a scenario input.

Formula

The standard method assumes a 100 g sample. A mass percent then becomes grams directly:

moles of element = mass percent / atomic mass

Each mole amount is divided by the smallest positive mole amount. The resulting decimal ratios are converted into small whole numbers. Sometimes the ratios are close to halves, thirds, or quarters, so the calculator tests small multipliers to find a clean integer ratio. The final empirical formula is the smallest whole-number version of that ratio.

Example

For 40.0% A with atomic mass 12.011, 6.7% B with atomic mass 1.008, and 53.3% C with atomic mass 15.999, the mole amounts are roughly 3.33, 6.65, and 3.33 moles on a 100 g basis. Dividing by the smallest value gives about 1, 2, and 1. The empirical ratio is therefore 1:2:1, and the generic formula is AB2C.

If the actual molar mass were twice the empirical formula mass, the molecular multiplier would be 2 and the molecular formula estimate would become A2B4C2. That distinction is why empirical and molecular formulas should not be treated as the same thing.

Interpretation and limitations

The empirical ratio is most reliable when the input percentages come from a complete and accurate elemental analysis. Rounding can make ratios look slightly off, especially when one element has a very small percentage. Hydrates, salts, mixtures, isotopic masses, impurities, and incomplete combustion data can also affect results. The calculator does not identify element symbols or validate chemical plausibility; it only performs the ratio math. Use it as a transparent worksheet for chemistry calculations and confirm final formulas with lab context or a chemistry reference.