Empirical & Molecular Formula Calculator
Empirical and molecular formula calculator. Converts percent composition into the empirical formula, then the molecular formula given the molar mass. Educational.
Empirical and Molecular Formula Calculator
Enter each element's symbol and its mass percent (leave a row blank to skip). Percentages should total about 100%.
How to use the empirical formula calculator
Enter each element and its mass %
Type the standard chemical symbol (C, H, O, N, S…) and the percentage of the compound's mass it contributes. The percentages should add up to roughly 100%.
Add the molar mass (optional)
If you know the compound's molar mass, enter it to get the molecular formula as well as the empirical formula.
Calculate
The tool converts percentages to moles, finds the simplest whole-number ratio for the empirical formula, then scales it to the molecular formula using your molar mass.
Check the result
If the ratios don't land cleanly on whole numbers, your input percentages may be rounded or slightly off. Verify against the known compound where possible.
Empirical and molecular formulas — from percentages to atoms
The simplest ratio, then the real molecule
A compound's formula can be written two ways. The empirical formula gives the simplest whole-number ratio of the atoms — for glucose it is CH₂O — while the molecular formula gives the actual numbers of each atom in one molecule — for glucose, C₆H₁₂O₆. The molecular formula is always a whole-number multiple of the empirical one. Deriving them from experimental data is a classic chemistry exercise, because elemental analysis tells you the percentage by mass of each element, not the atom counts directly. The route from one to the other is a tidy three-step recipe. First, treat the percentages as grams (imagine a 100 g sample) and convert each element's mass to moles by dividing by its atomic mass — this turns mass fractions into atom counts. Second, divide every mole value by the smallest of them, which scales the smallest element to 1 and expresses the others relative to it. Third, if the resulting ratios aren't already whole numbers, multiply them all by a small integer to clear the fractions (a ratio of 1.5 becomes 3 when doubled, for instance). The cleaned-up integers are the subscripts of the empirical formula.
Getting from the empirical formula to the molecular formula needs one more piece of information: the compound's molar mass, found separately by methods such as mass spectrometry. Divide the molar mass by the empirical formula's mass and round to the nearest whole number — that multiplier tells you how many empirical units make up one molecule. Multiply every subscript by it and you have the molecular formula. For glucose, the empirical CH₂O has a mass of about 30 g/mol; the molar mass of 180 g/mol divided by 30 is 6, so the molecular formula is six times CH₂O, giving C₆H₁₂O₆.
"Percent composition gives the empirical formula — the simplest atom ratio. Add the molar mass and you recover the real molecule, because the molecular formula is always a whole-number multiple of the empirical one."
Where the method needs care
The arithmetic is straightforward, but a few practical points decide whether you get the right answer. The element percentages must be reasonably accurate, because the method hinges on the ratios landing close to whole numbers; rounded or imprecise data can make a true ratio of, say, 2:3 look like 1.97:3.04, which is usually obvious but occasionally ambiguous. When a ratio comes out near a simple fraction — 1.5, 1.33, 1.25 — that's the signal to multiply through by 2, 3, or 4 rather than rounding to the nearest integer, which would lose information. The percentages should sum to about 100%; a large shortfall often means an element (frequently oxygen) was omitted and should be included as the balance. The empirical formula alone can't distinguish compounds that share it — formaldehyde (CH₂O), acetic acid (C₂H₄O₂), and glucose (C₆H₁₂O₆) all reduce to CH₂O — which is exactly why the molar mass is needed to pin down the molecular formula. And even the molecular formula doesn't reveal structure: isomers with identical molecular formulas can be very different substances. This calculator handles the percentage-to-formula computation and the molar-mass scaling; use it to learn the method and check your work, and confirm against known data for anything that matters.
10 Facts About Chemical Formulas
The empirical formula is the simplest atom ratio.
The molecular formula gives the actual atom counts.
Molecular formula = empirical × a whole number.
Step 1: moles = mass% ÷ atomic mass.
Step 2: divide all by the smallest.
Step 3: multiply to clear fractions.
Glucose: empirical CH₂O, molecular C₆H₁₂O₆.
CH₂O is shared by several compounds.
You need the molar mass for the molecular formula.
Formulas don't show structure — isomers differ.
Frequently asked questions
The empirical formula is the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of each atom in one molecule. They differ by a whole-number factor: glucose's empirical formula is CH₂O and its molecular formula is C₆H₁₂O₆ — six times the empirical unit. Percent composition gives you the empirical formula directly; you need the molar mass to find the molecular formula.
Three steps. Assume a 100 g sample so each element's percentage becomes grams, then divide each by its atomic mass to get moles. Next, divide every mole value by the smallest one, which sets the smallest element to 1. Finally, if the resulting ratios aren't whole numbers, multiply them all by a small integer to clear the fractions. The whole-number results are the subscripts. The calculator performs all three steps automatically.
Divide the compound's molar mass by the mass of the empirical formula and round to the nearest whole number — that's the multiplier. Multiply every subscript in the empirical formula by it to get the molecular formula. For glucose, the empirical CH₂O has a mass near 30 g/mol; 180 ÷ 30 = 6, so the molecular formula is C₆H₁₂O₆. You must measure the molar mass separately (for example by mass spectrometry); enter it in the calculator to get this step.
Because the true ratio isn't 1:1. A value near 1.5 means the real ratio is 2:3 (multiply by 2); 1.33 means 3:4 (multiply by 3); 1.25 means 4:5 (multiply by 4). You should multiply all the ratios by the small integer that turns them into whole numbers, rather than rounding 1.5 down to 1, which would give the wrong formula. The calculator detects these near-fractions and applies the right multiplier automatically.
A small shortfall (a percent or two) is just rounding. A large shortfall usually means an element was left out — very often oxygen, which isn't always measured directly and is reported "by difference." If your listed elements fall well short of 100%, add the missing element with the balance of the percentage. The method still works as long as every element is accounted for; an omitted element throws off all the ratios.
Yes, and this is exactly why the molar mass is needed. Formaldehyde (CH₂O), acetic acid (C₂H₄O₂), and glucose (C₆H₁₂O₆) all reduce to the empirical formula CH₂O but are completely different substances. The empirical formula fixes only the ratio of atoms, not their absolute number, so on its own it can't identify the molecule. The molar mass distinguishes them by setting the multiplier that scales the empirical formula to the real molecular formula.
It recognises the standard one- and two-letter symbols for the common elements found in most compounds — carbon, hydrogen, oxygen, nitrogen, sulfur, phosphorus, the halogens, the common metals, and more — using their standard atomic masses. Enter the symbol exactly as written on the periodic table (case matters: "Co" is cobalt, "CO" would be misread). If a symbol isn't recognised, the calculator flags it so you can correct a typo.
No. The molecular formula gives the atom counts but not how the atoms are connected. Different compounds can share the same molecular formula — these are isomers — and have very different properties; ethanol and dimethyl ether are both C₂H₆O, for example. Determining the actual arrangement of atoms (the structural formula) requires further techniques such as spectroscopy. The formula is a necessary first step, but identifying a compound fully goes beyond it.
Use it to check your answers and learn the procedure, but make sure you can perform the mole conversion and ratio steps yourself, since that's what's assessed. Enter symbols carefully and confirm the percentages account for every element. This tool runs the standard method and is educational; reproducing the calculation by hand builds the understanding you'll need in an exam or lab.
No. The values you enter are processed entirely in your browser. Nothing is sent to a server, stored, or shared, and no account is required. The calculation runs on your device only.
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