Moles to Grams Converter
Common Conversions
| mol | g |
|---|---|
| 0.001 | 0.018 |
| 0.01 | 0.18 |
| 0.05 | 0.901 |
| 0.1 | 1.802 |
| 0.25 | 4.504 |
| 0.5 | 9.008 |
| 1 | 18.015 |
| 2 | 36.03 |
| 5 | 90.075 |
| 10 | 180.15 |
Why this conversion matters in chemistry
Moles to grams is the conversion you run every single time you step up to a balance. A procedure calls for 0.25 mol of sodium bicarbonate, but the balance reads in grams, so you multiply by the molar mass (84.007 g/mol for NaHCO₃) and weigh out 21.00 g. It's arithmetic that never stops being useful — stoichiometry, titration prep, reaction scale-up, anywhere a paper reports something in moles and a bench gives you a weighing boat.
Formula
Where the factor comes from
Every other conversion in this family runs on a defined power of ten. This one runs on a property of the substance in front of you, and that is the whole difference. Molar mass is the sum of the standard atomic weights in the formula multiplied by the molar mass constant, and none of those inputs is stipulated. Standard atomic weights are measured, revised as better measurements arrive, and for elements whose isotopic composition varies with source — hydrogen, carbon, oxygen, sulfur and chlorine among them — IUPAC publishes an interval rather than a single number. The molar mass constant is measured too: since the 2019 revision of the SI it is no longer exactly 1 g/mol, though it departs from that value by only about one part in 10⁹. The factor here is never exact, and it changes every time the compound does.
Precision and significant figures
A molar mass carries more digits than the weighing that follows it will ever justify. Sodium bicarbonate at 84.007 g/mol gives 21.00175 g for 0.250 mol, but the 0.250 mol claimed three figures, so the answer is 21.0 g and the rest is calculator residue. Match the output to the weaker input, which is almost always the mole figure rather than the mass — a four-place balance resolving 0.1 mg is rarely what limits you. Carry the extra decimals through the arithmetic, since they cost nothing, and drop them at the end. For elements published as intervals, the fourth decimal place of a molar mass is not defensible in the first place.
Worked Examples
A mole of water, the reference point most chemists keep in their head.
Two moles of sodium chloride for a 2 M stock in a liter of water — close to the upper end of routine benchwork.
Half a mole of sulfuric acid. You'd measure this by volume rather than mass in practice, but the number anchors the density calculation.
A tenth of a mole of glucose for a standard curve. Handy coincidence that 0.1 mol glucose weighs almost exactly the same as 1 mol water.
Common mistakes
Hydrate weighed as the anhydrous salt
Copper(II) sulfate is 159.6 g/mol anhydrous and 249.7 g/mol as the pentahydrate, so 0.100 mol is either 15.96 g or 24.97 g depending on which bottle you reached for. The label on the shelf decides the molar mass, not the formula written in the procedure, and a partly dehydrated bottle fading from blue toward white sits somewhere between the two with no way to tell by eye.
Moles of O₂ read as moles of O
One mole of O₂ weighs 32.00 g; one mole of oxygen atoms weighs 16.00 g. Balanced equations written with diatomic elements — O₂, N₂, H₂, Cl₂ — give amounts of the molecule, while a mass pulled straight off a periodic table gives the atom. The resulting factor of two survives every later step of the stoichiometry and rarely announces itself.
Reagent purity ignored at the balance
Multiplying moles by molar mass tells you the mass of pure compound. A bottle assayed at 97% needs that target divided by 0.97 before anything goes on the weighing paper. Technical grades, deliquescent salts that have taken up water in a humid room, and anything quoted on a dried basis all break the assumption that mass in equals substance in.