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General Chemistry I · Stoichiometry

Molar Mass

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Want it in plain words first? Jump to Eli explains — the same idea, no jargon.
On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Study tools
  7. Sources & references

In 30 seconds

Molar mass is the mass in grams of one mole of a substance, expressed in g/mol. It is numerically equal to the formula mass in atomic mass units: you sum the atomic masses of every atom in the chemical formula. For molecular compounds this is the molecular mass (one molecule); for ionic compounds it is the formula mass (one formula unit). Because one mole of any element weighs its atomic mass in grams, adding up the atoms in a formula gives, in one step, the mass of 6.022 × 10²³ formula units.

Why this matters

Molar mass is the number you type into every "grams to moles" conversion in the laboratory. Preparing a 0.100 M solution, dosing a reactant, or computing a percent yield all begin by computing molar mass correctly. A missed parenthesis (writing CaNO₃ instead of Ca(NO₃)₂) changes the mass by a factor of two — a real error that ruins an experiment.

The college version

Key Ideas

Atomic mass vs. molar mass vs. formula mass

  • Atomic mass is the weighted-average mass of one atom of an element, in amu (e.g., C = 12.01 amu).
  • Molar mass of an element is that same number in grams per mole (C = 12.01 g/mol).
  • Molecular mass is the sum for one molecule; formula mass is the sum for one formula unit of an ionic compound.

Reading a chemical formula

  • A subscript after an atom multiplies only that atom (H₂O → 2 H + 1 O).
  • Parentheses distribute a subscript to everything inside: Ca(NO₃)₂ → 1 Ca, 2 N, 6 O.
  • Hydrates add whole water molecules: CuSO₄·5H₂O → 1 CuSO₄ plus 5 H₂O.

Equations and Variables

  • Molar mass (M) = Σ (number of atoms × atomic mass), units g/mol.
  • Moles from mass: n = m / M
  • Mass from moles: m = n × M
  • Molar mass of water: M(H₂O) = 2(1.008) + 16.00 = 18.016 ≈ 18.02 g/mol

How It Works (Problem-Solving Method)

  1. Write the formula and list each distinct element.
  2. Count each element's atoms, expanding parentheses and hydrates fully.
  3. Look up atomic masses on the periodic table (typically to 2–4 significant figures).
  4. Multiply count × atomic mass for each element and sum the products.
  5. Attach units g/mol and round to the appropriate significant figures.

Worked Example

Calculate the molar mass of calcium nitrate, Ca(NO₃)₂.

  • Expand the formula: 1 Ca, 2 N, and 2 × 3 = 6 O.
  • Atomic masses: Ca = 40.08, N = 14.01, O = 16.00 g/mol.
  • Ca: 1 × 40.08 = 40.08 g/mol
  • N: 2 × 14.01 = 28.02 g/mol
  • O: 6 × 16.00 = 96.00 g/mol
  • Total = 40.08 + 28.02 + 96.00 = 164.10 g/mol

Hydrate example — copper(II) sulfate pentahydrate, CuSO₄·5H₂O:

  • CuSO₄ = 63.55 + 32.07 + 4(16.00) = 159.62 g/mol
  • 5 H₂O = 5(18.02) = 90.10 g/mol
  • Total = 159.62 + 90.10 = 249.72 g/mol (the water of hydration is counted).

Common Confusions

  • "Molar mass and atomic mass are the same thing." — Same number, different units and meaning: atomic mass is one atom in amu; molar mass is one mole in grams.
  • "Ca(NO₃)₂ has 2 N and 2 O." — The subscript 2 outside the parentheses multiplies the whole nitrate group, giving 2 N and 6 O.
  • "The water in a hydrate is not part of the molar mass." — It is. CuSO₄·5H₂O includes all five water molecules in its 249.72 g/mol.
  • "Rounding atomic masses to whole numbers is fine." — It changes results meaningfully; use at least the tenths place (16.00, 12.01) unless told otherwise.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of the periodic table as a price list for atoms. Carbon "costs" about 12, oxygen about 16, hydrogen about 1. To find how much one "box" of a molecule weighs, you add up the prices of all its pieces. H₂O is two hydrogens (1 + 1) plus one oxygen (16) = 18. The one catch: these aren't dollars, they're atomic mass units — but the number happens to be the same as the grams in one whole mole, which is why chemists can use the "price list" to weigh real chemicals.

Key takeaways

  • Molar mass has units of g/mol; atomic mass has units of amu.

  • Numerically, molar mass (g/mol) = formula mass (amu).

  • A subscript multiplies only the atom it follows; parentheses distribute to everything inside.

  • In a hydrate, include the mass of the water molecules (each H₂O = 18.02 g/mol).

  • Molar mass = sum of (atoms × atomic mass) over the whole formula.

  • Molar mass = mass of one mole, in g/mol.

  • Sum (atoms × atomic mass) across the formula.

  • Expand parentheses and count hydrate water.

  • Use molar mass as the grams↔moles conversion factor.

Keep learning

Ready to build on this? Continue to the next lesson.

Practice General Chemistry I

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Define molar mass and explain its units (g/mol).
  • Calculate the molar mass of a compound from its chemical formula and periodic-table atomic masses.
  • Handle parentheses and hydrate water molecules when summing atomic masses.
  • Distinguish molar mass (g/mol) from atomic mass (amu) and molecular/formula mass.

Sources & references

  1. OpenStax, "3.1 Formula Mass and the Mole Concept." *Chemistry 2e*.
  2. Chemistry LibreTexts, "10: The Mole."
  3. NIST, "CODATA Value: Avogadro constant."

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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