Empirical Formula Calculator
Simplest whole number formula from two element masses with mole ratio shown.
What it does
Empirical formulas reduce a compound to its simplest whole atom ratio: rust analyses to Fe2O3, water to H2O. Masses convert to moles first, the smallest count divides the rest, and a small multiplier clears any halves.
The tool above takes two element masses with their symbols. It returns the empirical formula with the mole counts and the ratio plus multiplier beside them.
The default of 34.97 g iron with 15.03 g oxygen gives 0.6262 mol Fe and 0.9394 mol O, ratio 1 to 1.5, doubled to Fe2O3. Water style numbers of 11.11 g H with 88.89 g O give 11.02 and 5.56 mol, ratio 1.98 to 1, H2O.
The steps follow open textbook stoichiometry chapters with worked examples matching each move. Both divide by the smallest moles and multiply out halves like 1.5 before rounding.
How to use
- Pick the two elements in the sample.
- Enter each mass in grams (percents read as grams of 100 g).
- Read moles first for the raw counts.
- Read the ratio with its clearing multiplier.
- Read the formula for the simplest whole answer.
Formula
- moles
- Grams divided by atomic weight per element.
- divide
- Each count over the smallest count.
- x2-x6
- Multiplier clearing halves and thirds.
Oxygen prints last by convention here. Three element samples need the same path with one more division; this page pairs the two biggest shares or runs twice.
Worked example
Example: 34.97 g Fe with 15.03 g O
Defaults: Fe 34.97 g, O 15.03 g.
- Fe moles: 34.97 / 55.845 = 0.6262.
- O moles: 15.03 / 15.999 = 0.9394.
- Divide by 0.6262: Fe 1, O 1.5.
- Clear the half with x2: Fe2O3.
- Tool shows Fe2O3, the mole counts, 1 : 1.5 (x2).
Reference table
| Mole ratio | Multiplier | Formula shape |
|---|---|---|
| 1 : 1 | x1 | AB |
| 1 : 1.5 | x2 | A2B3 |
| 1 : 1.33 | x3 | A3B4 |
| 1 : 1.98 | x1 | AB2 (rounds) |
Limits
- Two elements only; third needs a second run.
- Small ratios only: exotic fractions stay unresolved.
- Analysis error bends ratios near rounding lines.
- Empirical is simplest, not molecular: CH2 pairs with C2H4.
Common mistakes
- Dividing grams directly without moles.
- Rounding 1.5 to 2 instead of doubling.
- Forgetting percents mean grams of 100 g.
- Writing oxygen first against convention.
- Reading empirical as the molecular formula.
Try it yourself
Quick checks. Answers sit tucked under each question.
11.11 g H with 88.89 g O?
11.02 and 5.56 mol, ratio 1.98:1, H2O.
Same element twice?
Rejected: pick two different elements.
1.33 ratio?
Triples to 3:4, e.g. A3B4 shape.
Questions
What is an empirical formula?
The simplest whole number atom ratio: Fe2O3 for rust, H2O for water.
How do I find it?
Masses to moles, divide by the smallest, multiply to clear fractions.
Can I use percents?
Yes. Read them as grams of a 100 g sample and run the same path.
Why multiply by 2?
Halves like 1.5 never round: doubling gives the true 2:3.
Empirical versus molecular?
Empirical is simplest; molecular multiplies it out with the molar mass.
What if no small ratio fits?
The tool says so: check analysis or split a third element out.
Sources
- LibreTexts: Percent composition, empirical and molecular formulas
- LibreTexts: Empirical and molecular formulas (CCSF)