Grams to Moles Calculator

How to use this calculator

Two inputs. That’s the entire interface.

Chemical formula — Type it in standard notation. The calculator reads each element symbol, subscript, and parenthetical group and parses the full molecular composition. Accepted formats include simple compounds (H2O, NaCl, CO2, NH3), compounds with parentheses (Ca(OH)2, Al2(SO4)3), organic molecules (C6H12O6, C2H5OH), hydrates (CuSO4·5H2O, Na2CO3·10H2O), and complex inorganics (Fe2(SO4)3, KMnO4).

Mass in grams — Enter the mass of your sample. If you weighed it on a lab balance, enter the number directly. If your problem gives you milligrams, convert first by dividing by 1,000.

After clicking Calculate, you get the total moles (primary result), molar mass in g/mol, and an element-by-element breakdown showing how molar mass was computed.

Why the gram-to-mole conversion matters

Chemistry equations don’t describe mass. They describe particle ratios.

When you see 2H₂ + O₂ → 2H₂O, the equation says two molecules of hydrogen gas react with one molecule of oxygen to produce two molecules of water. Scale that up by Avogadro’s number (6.022 × 10²³) and you get the same ratio in moles: two moles of H₂ react with one mole of O₂ to produce two moles of H₂O.

The problem is that you can’t measure moles on a scale. You measure grams. To actually run this reaction in a lab, you need to convert the mole ratio into grams. That conversion step sits at the centre of nearly every stoichiometry calculation. Get it wrong once and the entire downstream result is off by the same ratio.

Grams tell you how much the sample weighs. Moles tell you how many particles it contains. Chemistry equations need the second number.

The formula

One division. That’s it.

The molar mass is the number you look up or calculate. It’s the mass of one mole of the substance in grams per mole. For any compound, it equals the sum of the atomic masses of every atom in the formula.

Hydrogen has an atomic mass of 1.008. Oxygen is 15.999. For H₂O: (2 × 1.008) + 15.999 = 18.015 g/mol.

How molar mass is calculated element by element

Take Ca(OH)₂ as a worked example, because it’s where most students make their first parentheses error.

Element	Atom count	Atomic mass (g/mol)	Contribution
Ca	1	40.078	40.078
O	2 (from OH × 2)	15.999	31.998
H	2 (from OH × 2)	1.008	2.016
Total			74.092 g/mol

The subscript 2 after the closing parenthesis multiplies every atom inside the brackets by 2. So (OH)₂ means 2 oxygen atoms and 2 hydrogen atoms.

A common mistake is reading it as Ca + O + 2H = 40.078 + 15.999 + 2.016 = 58.093 g/mol. That’s wrong by 16 g/mol — a 27% error — because one oxygen atom was lost.

Molar masses of common compounds

Compound	Formula	Molar mass (g/mol)	Note
Water	H₂O	18.015	Universal reference
Carbon dioxide	CO₂	44.009	Common in gas problems
Sodium chloride	NaCl	58.443	Table salt
Ammonia	NH₃	17.031	Industrial base
Glucose	C₆H₁₂O₆	180.156	Common in biology
Sulfuric acid	H₂SO₄	98.072	Strong acid
Sodium hydroxide	NaOH	39.997	Strong base (lye)
Calcium carbonate	CaCO₃	100.086	Limestone, chalk
Hydrochloric acid	HCl	36.461	Strong acid
Ethanol	C₂H₅OH	46.068	Alcohol
Copper sulfate pentahydrate	CuSO₄·5H₂O	249.679	Hydrate example
Iron oxide	Fe₂O₃	159.687	Rust
Potassium permanganate	KMnO₄	158.034	Oxidising agent

Real worked examples

36 grams of water

117 grams of sodium chloride

180 grams of glucose

Copper sulfate pentahydrate

Stoichiometry: why mole ratios matter

Balanced equations give you mole ratios, not gram ratios. This is the entire foundation of stoichiometry.

Consider the synthesis of water. The balanced equation is 2H₂ + O₂ → 2H₂O, which gives a mole ratio of 2 mol H₂ : 1 mol O₂ : 2 mol H₂O.

If you start with 4 grams of hydrogen:

Every step in that chain goes through the grams-to-moles conversion. The conversion itself is simple. The complexity is in tracking which substance you’re converting and which ratio from the balanced equation applies.

Common mistakes

Subscript errors. CO has a molar mass of 28.010 g/mol (C + O). CO₂ has a molar mass of 44.009 g/mol (C + 2O). These are different compounds. The subscript 2 isn’t a minor detail — calculating with CO instead of CO₂ gives an answer that’s off by one oxygen atom (15.999 g/mol), and there’s no way to know from the number alone that anything went wrong.

Parentheses errors. Al₂(SO₄)₃ contains 2 aluminium atoms, 3 sulfur atoms, and 12 oxygen atoms (4 inside the brackets × 3 outside). Missing the parenthesis multiplier gives 8 oxygens instead of 12. The molar mass drops from 342.15 g/mol to 294.15 g/mol — a 16% error.

Unit confusion: grams vs milligrams. The formula uses mass in grams. If your balance reads in milligrams (common in analytical chemistry), divide by 1,000 first. 500 mg = 0.5 g, not 500 g. Skip the conversion and your mole count is off by a factor of 1,000.

Wrong direction: dividing vs multiplying. Grams to moles means dividing by molar mass. Moles to grams means multiplying by molar mass. These directions get swapped on a significant number of student worksheets — write them down separately if you need to keep them straight.

Using anhydrous molar mass for a hydrated compound. If the solid in front of you is the blue hydrated form of copper sulfate, its molar mass is 249.68 g/mol — not 159.60 g/mol. Using the anhydrous value gives a mole count that’s 56% too high.

Hydrates: the extra water that counts

Hydrated salts have water molecules chemically bound into the crystal structure. When you weigh out copper sulfate pentahydrate (the blue solid in most school chemistry labs), you’re weighing CuSO₄·5H₂O. The “·5H₂O” part is not an impurity — it’s part of the compound. Those 5 water molecules per formula unit add 5 × 18.015 = 90.075 g/mol to the molar mass.

Hydrate	Anhydrous molar mass	Hydrated molar mass	Difference
CuSO₄·5H₂O	159.60 g/mol	249.68 g/mol	+56.4%
Na₂SO₄·10H₂O	142.04 g/mol	322.19 g/mol	+126.8%
MgSO₄·7H₂O (Epsom salt)	120.36 g/mol	246.47 g/mol	+104.8%
Na₂CO₃·10H₂O (washing soda)	105.99 g/mol	286.14 g/mol	+170.0%

Using the anhydrous molar mass for a hydrated compound gives a calculation that’s off by 57% to 170%. This is the single largest source of error in students preparing solutions from hydrated salts.

Significant figures and when they matter

A molar mass of 18.015 g/mol has five significant figures. If your mass measurement is 36 g (two significant figures), your answer should be reported to two significant figures: 2.0 mol, not 1.999 mol.

The rule: your answer can’t be more precise than your least precise input.

In practice, homework problems often use exact round numbers (36 g of H₂O) and expect a round answer (2 mol). Lab work uses the actual precision of your balance (typically 4 decimal places for analytical balances). Exam mark schemes usually specify significant figures explicitly.

Limitation: standard atomic masses vs isotope-specific masses

The calculator uses standard atomic weights from the IUPAC periodic table. These are weighted averages across all naturally occurring isotopes of each element.

Hydrogen’s standard atomic weight is 1.008, which reflects the fact that natural hydrogen is 99.985% ¹H (mass 1.00783) and 0.015% deuterium (mass 2.01410).

For most chemistry, standard atomic weights are exactly what you want. But for specialized applications — mass spectrometry, isotope labelling, carbon-13 NMR, radioactive decay calculations — the standard masses don’t apply. Those fields use monoisotopic masses: the mass of the most abundant single isotope, not the weighted average. If you’re working in one of those areas, this calculator will give you the wrong molar mass.

The standard atomic masses in every general chemistry course are weighted averages. They're right for bench chemistry. They're wrong for mass spectrometry. Know which one you need before you calculate.

What to do with the moles result

Number of molecules — multiply moles by 6.022 × 10²³. Two moles of water contains 1.204 × 10²⁴ molecules.

Stoichiometry — apply the mole ratio from your balanced equation to find how many moles of reactant you need or how much product you’ll produce.

Molarity — if you’re dissolving the substance in solution, moles ÷ volume in litres = molarity (mol/L). Two moles of NaCl dissolved in 1 litre of solution = 2M.

Percent yield — actual moles produced ÷ theoretical moles expected × 100%. Both numbers need to be in moles before you can compare them fairly.

Limiting reagent — convert grams of each reactant to moles, divide each by its stoichiometric coefficient, and the smallest result identifies the limiting reagent.

Every one of these starts with the grams-to-moles step. Get that right and the rest of the calculation follows the rules.

The bottom line

Grams to moles is one division. The difficulty is in assembling the right molar mass before you divide — getting the formula right, parsing the parentheses correctly, including the hydrate water if it’s there, and using the right atomic masses.

The calculator handles the molar mass assembly automatically. Your job is entering the formula correctly. Double-check subscripts, include parentheses where they belong, and match the formula to the actual compound you’re working with — hydrated or anhydrous.

Get those inputs right and the output is reliable for every downstream calculation that depends on it.