Molar Mass Converter
Converters · Added 15 August 2026
Enter a chemical formula and get its molar mass in grams per mole, the percentage each element contributes, and a conversion between grams and moles. Brackets and hydrate notation are handled, so Ca(OH)₂ and CuSO₄·5H₂O parse correctly, and an unrecognised symbol produces an error rather than a plausible wrong answer.
Result
Molar mass of H2O
18.0150 g/mol
3 atoms per formula unit
- Mass
- 100 g
- Amount of substance
- 5.55092978 mol
- Formula units (particles)
- 3.3428e+24
Percentage composition by mass
| Element | Atoms | Atomic weight | Mass (g/mol) | % of total |
|---|---|---|---|---|
| OOxygen | 1 | 15.9990 | 15.9990 | 88.81% |
| HHydrogen | 2 | 1.0080 | 2.0160 | 11.19% |
Standard atomic weights are averages over the isotopes found in normal terrestrial material, so this is the molar mass of ordinary laboratory material rather than the mass of one specific isotope. Enriched or depleted samples differ.
How to use the molar mass converter
- 1Type a chemical formula — case matters, since Co is cobalt and CO is carbon monoxide.
- 2Use brackets for groups, as in Al2(SO4)3, and a dot or full stop for hydrates, as in CuSO4·5H2O.
- 3Read the molar mass and the percentage composition by element.
- 4Enter a mass in grams to see the equivalent in moles, or switch the direction.
- 5Tap one of the example formulas to see the format.
Examples
Water
- Input
- H2O
- Result
- 18.015 g/mol
Oxygen accounts for 88.81% of the mass despite being one atom in three.
A hydrate
- Input
- CuSO4·5H2O
- Result
- 249.681 g/mol
The five water molecules add over 90 g/mol — which is why hydrated and anhydrous forms are not interchangeable by weight.
Grams to moles
- Input
- 100 g of NaCl
- Result
- 1.7112 mol
100 divided by the 58.44 g/mol molar mass.
About the molar mass converter
Why chemistry counts in moles
Chemical reactions happen between fixed numbers of particles, not fixed masses. Two hydrogen molecules react with one oxygen molecule to give two waters, regardless of what those molecules weigh. But you cannot count molecules on a balance — you can only weigh them.
The mole bridges that gap. Because one mole of any substance contains the same number of particles, and because molar mass in grams per mole matches the molecular mass in atomic mass units, weighing out a mass in grams becomes a way of counting particles. Working out how much of each reactant to weigh is the everyday use, and it is why molar mass is the first calculation in most practical chemistry.
The number itself was chosen to make that correspondence exact rather than for any independent reason. Avogadro's constant is now fixed by definition at 6.02214076 × 10²³, which means the mole is defined by counting rather than by reference to a kilogram of anything.
Reading a formula correctly
A chemical formula is a compact notation with a few rules that all matter. A subscript applies to the symbol immediately before it, so H2O is two hydrogens and one oxygen. A subscript after a closing bracket applies to everything inside, so Ca(OH)2 has two oxygens and two hydrogens, not one of each.
The dot in a hydrate is not a multiplication or a decimal point. CuSO4·5H2O means one formula unit of copper sulfate accompanied by five water molecules in the crystal lattice — water that is genuinely part of the compound and is driven off by heating. The anhydrous and pentahydrate forms have very different molar masses, so a recipe calling for one and weighed as the other is out by roughly 36%.
Where the same element appears more than once, the counts add: CH3COOH is acetic acid with two carbons, four hydrogens and two oxygens, even though the hydrogens are written in two separate places. This tool accumulates repeated symbols rather than overwriting them, which is a common failing in simpler parsers and produces silently wrong answers on exactly this kind of structural formula.