How do I calculate percentage composition?

For each element, multiply how many atoms of it the formula contains by its atomic mass, divide by the compound’s molar mass, and multiply by a hundred. The molar mass is just the sum of all those contributions, so the percentages necessarily add to 100 — if yours do not, the arithmetic has gone wrong somewhere rather than the chemistry being unusual.

Why do the percentages always add up to 100?

Because every one of them is a share of the same denominator. The molar mass is defined as the total of each element’s contribution, so dividing each contribution by that total and summing gives the total divided by itself. It is not a coincidence or a check that happened to pass — it is arithmetic that cannot come out otherwise, which makes a sum that misses a reliable sign of a mistake.

How do I find the empirical formula from percentages?

Treat the percentages as grams in a hundred-gram sample. Divide each by its atomic mass to get moles, divide all the results by the smallest, and you have a ratio. If that ratio is near whole numbers, round it. If it is not, multiply through by small integers until it is — a value near .5 needs doubling, near .33 or .67 needs tripling, near .25 or .75 needs quadrupling.

Why does the ratio need multiplying sometimes?

Because rounding a fractional ratio invents a compound. Phosphorus and oxygen at 43.6% and 56.4% give a ratio of 1 : 2.50 — round that and you get PO₃, which does not exist. Multiplying by two gives 2 : 5, or P₂O₅, which does. The step is easy to skip because 2.50 looks close enough to 2 or 3 that rounding feels reasonable, and it is the single most common error in this calculation.

What is the difference between an empirical and a molecular formula?

The empirical formula gives the simplest whole-number ratio of atoms; the molecular formula gives the actual count in one molecule. CH₂O is the empirical formula for formaldehyde, acetic acid and glucose alike — the same ratio at three different sizes. A composition analysis can only ever give the empirical formula, because ratios have no scale. Turning it into a molecular formula needs a separately measured molar mass, and the molecular formula is then a whole-number multiple of the empirical one.

Can two different compounds have the same percentage composition?

Routinely, and this is the measure’s real limit. Glucose and fructose are both C₆H₁₂O₆ and their compositions agree to every decimal place. Ethanol and dimethyl ether are both C₂H₆O. Butane and isobutane are both C₄H₁₀. Composition fixes how many of each atom are present and says nothing at all about how they are joined, so isomers are invisible to it. Identifying which compound you have needs spectroscopy, not analysis of the mass percentages.

How do I handle a formula with brackets or a hydrate dot?

Brackets multiply everything inside them by the subscript that follows, so Ca(OH)₂ has one calcium, two oxygens and two hydrogens. A hydrate dot separates the compound from the attached water, and the number before the water multiplies it: CuSO₄·5H₂O has five waters, contributing ten hydrogens and five oxygens on top of the four already in the sulfate. Both are handled here, and brackets can be nested.

Why does hydrogen contribute so little mass?

Because it is the lightest element, at 1.008 against carbon’s 12.011 and oxygen’s 15.999. Glucose contains twice as many hydrogen atoms as carbon atoms and they account for about a sixth as much mass. It is a useful corrective to the way formulas are read: the subscripts count atoms and the percentages weigh them, and for hydrogen those two pictures diverge sharply.

What if my percentages do not add to 100?

A ratio ignores scale, so a formula will still come out — which is exactly why it is worth flagging. A total a little under 100 is usually rounding. A total well under usually means an element was not measured: in combustion analysis the oxygen is often obtained by difference rather than directly, and forgetting it is a common slip. A total above 100 means at least one figure is too high, and the resulting ratio will be wrong in proportion.

Which atomic masses should I use?

The IUPAC standard atomic weights, which are what this page uses — the conventional single values rather than the interval forms published for elements with variable isotopic composition. School textbooks sometimes round further, using 12 for carbon or 16 for oxygen, which shifts a percentage by a few hundredths. If your answer differs from a textbook in the second decimal place, that is almost certainly why rather than an error in the working.

Percentage Composition Calculator

Percentage composition is the share of a compound's mass contributed by each element: the count of that element times its atomic mass, divided by the molar mass, times a hundred. The arithmetic is one line. Parsing the formula is the part that takes work, so this page does it — nested brackets, hyd...

% of an element = count × atomic mass ÷ molar mass × 100

Type a formula and it is parsed — brackets, hydrate dots and two-letter symbols included. Or work backwards from percentages to an empirical formula, which is the harder direction and the one exams tend to ask about.

THE COMPOUND

CHEMICAL FORMULA

SAMPLE MASS (G) — OPTIONAL

MOLAR MASS

180.156 g/mol

3 elements: C×6, H×12, O×6. The percentages below sum to exactly 100, which is a property of the definition rather than a check that happened to pass.

Element	Count	Atomic mass	Contributes	%
C	6	12.011	72.066	40.00%
H	12	1.008	12.096	6.71%
O	6	15.999	95.994	53.28%

These percentages cannot tell you what the compound is. Glucose and fructose are both C₆H₁₂O₆ and share every figure above to the last decimal place. So do ethanol and dimethyl ether, and butane and isobutane. Composition fixes how many of each atom are present and says nothing whatever about how they are joined — which is why an analysis yields an empirical formula and a structure needs spectroscopy.

THE BREAKDOWN

Read the bar as a hundred grams of the compound. Each block is the mass that element contributes, and because every block is a share of the same molar mass they necessarily add to the whole. Hydrogen is the usual surprise — twelve hydrogens in glucose weigh less than one oxygen.

IUPAC standard atomic weights · live calculation

Created with❤️byeaglecalculator.com

HOW TO USE

1
Type a chemical formula. Case matters — CO is carbon monoxide and Co is cobalt — and brackets may be nested. Hydrates take a dot, as in CuSO4·5H2O.
2
Read the breakdown. Each element contributes its count times its atomic mass, and each percentage is that contribution as a share of the molar mass.
3
Add a sample mass if you want the percentages in grams rather than as proportions. It changes nothing about the composition, only the units it is reported in.
4
Switch to the reverse mode to go from percentages to an empirical formula, and supply a molar mass there if you want the molecular formula as well.

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Last updated: August 2, 2026 · IUPAC standard atomic weights · Parser verified against 18 published molar masses · Composition cannot distinguish isomers.