Molar Mass Reference Calculator

Look up molar mass for common compounds and calculate formula mass from atomic weights.

Compound

Mass of sample (g) (g)

Desired concentration (mol/L) (mol/L)

Volume of solution (L) (L)

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The Mole as a Counting Unit

Chemical reactions happen between numbers of particles, but laboratories measure mass. The mole bridges the two: molar mass converts between grams on a balance and particles in a reaction.

n = m/M m = n × M c = n/V

Since the 2019 SI redefinition, a mole is exactly 6.02214076 × 10²³ entities — a fixed number rather than one derived from the mass of carbon-12. In practice the change is invisible at ordinary precision.

Step	Operation	Watch for
Mass to moles	Divide by molar mass	Use the formula mass, not atomic mass
Moles to particles	Multiply by 6.022 × 10²³	—
Moles to volume	Divide by concentration	Litres, not millilitres
Hydrates	Include the water	CuSO₄·5H₂O is 249.7, not 159.6

Hydrates and Purity

Two practical traps dominate real solution preparation. Hydrated salts carry water in their formula mass — copper sulfate pentahydrate is 249.7 g/mol against 159.6 for the anhydrous form, a 56% difference that would badly misprepare a solution.

Purity matters equally. A reagent labelled 98% requires dividing the calculated mass by 0.98 to deliver the intended amount. For analytical work this correction is routine; primary standards are chosen precisely because their purity is certified and stable.

Worked Examples

Example 1: 10g NaCl dissolved: MW=58.44

moles=10/58.44=0.171 mol

Result: To make 0.1mol/L in 1L: need 5.844g

Standard NaCl solution preparation

Example 2: To make 1L of 0.5M glucose

mass=0.5×1×180.16

Result: mass=90.08g glucose needed

Biological media preparation

Example 3: Hydrated salt

Prepare 500 mL of 0.1 M CuSO₄ from the pentahydrate

Result: 12.48 g, not 7.98 g

n = 0.05 mol, and M = 249.7 g/mol including the five waters. Using the anhydrous mass would give a solution 36% too dilute.

Example 4: Purity correction

Need 5.844 g of NaCl, reagent is 99.5% pure

Result: Weigh 5.873 g

Divide by 0.995. For most work the difference is small, but it matters for standards and analytical calibration.

Common Mistakes

⚠️

Ignoring water of crystallisation

CuSO₄·5H₂O has a molar mass of 249.7 g/mol, not 159.6. Using the anhydrous value gives a solution over a third too dilute.

⚠️

Forgetting the purity correction

A 98% pure reagent needs the calculated mass divided by 0.98. Skipping this systematically under-delivers the solute.

⚠️

Confusing molar mass with molecular mass

Molar mass carries units of g/mol; molecular mass is in atomic mass units. Numerically equal, but the distinction matters in dimensional analysis.

⚠️

Mixing millilitres with litres

Concentration is moles per litre. Using millilitres without converting gives an answer wrong by a factor of 1,000.

Frequently Asked Questions

Mole concept?

1 mole = 6.022×10²³ particles (Avogadro's number). Molar mass = mass of 1 mole = numerically equal to molecular weight in amu. 18g water = 1 mol = 6.022×10²³ water molecules.

Percent composition?

Mass% of element = (n × atomic mass / molar mass) × 100. For NaCl: Na% = 22.99/58.44 × 100 = 39.3%. Used to verify compound identity and for elemental analysis calculations.

How do I convert mass to moles?

Divide the mass in grams by the molar mass in g/mol. The molar mass must be the full formula mass, including any water of crystallisation.

Why does a hydrate have a different molar mass?

Because water molecules are part of the crystal structure and contribute to the mass weighed out. CuSO₄·5H₂O is 249.7 g/mol against 159.6 anhydrous.

Do I need to correct for reagent purity?

For accurate work, yes. Divide the calculated mass by the purity fraction. A 98% reagent needs about 2% more mass.

What is the current definition of the mole?

Since 2019 it is exactly 6.02214076 × 10²³ entities, a fixed number rather than one tied to the mass of carbon-12.

Formula Explorer connections

Interpretation: This relationship converts chemical amount, mass, composition or balanced-equation ratios into a reaction quantity. Assumption: Use a balanced reaction, consistent units and the correct molar mass. Purity, side reactions and limiting reagents can change experimental results.

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