The Mole & Chemical Formulas Cheat Sheet

What you're counting

CORE RULE

Molecular compounds: count molecules (H₂O, CO₂). Ionic compounds: count formula units (NaCl). Elements: atoms (Na, Fe) or molecules (H₂, O₂, P₄, S₈).

Moles of each element

CORE RULE

Each subscript = moles of that element per mole of compound. 1 mol Al₂(SO₄)₃ = 2 mol Al, 3 mol S, 12 mol O. Hydrate waters add to the count.

Molar mass

CORE RULE

Sum the atomic masses of every atom in one formula unit (expand parentheses, include hydrate water), in g/mol. Al₂(SO₄)₃ = 2(26.98) + 3(32.06) + 12(16.00) = 342.14 g/mol.

Mass ↔ moles

CORE RULE

Mass → moles: × (1 mol / molar mass). Moles → mass: × (molar mass / 1 mol). g/mol cancels; carry units to self-check.

Avogadro's number

CORE RULE

Nₐ = 6.022 × 10²³ entities/mol. Moles → particles: × Nₐ. Particles → moles: ÷ Nₐ. Atoms of an element = mol compound × subscript × Nₐ.

Hydrates

CORE RULE

Use the full hydrate molar mass (including water). 1 mol CuSO₄·5H₂O = 1 mol anhydrous CuSO₄ + 5 mol H₂O.

Atom count includes the waters: CuSO₄·5H₂O has 9 O (4 + 5) and 10 H (2 × 5).

Dissociation uses the anhydrous ions only: CuSO₄·5H₂O → Cu²⁺ + SO₄²⁻ + 5H₂O (2 ions, not 7).

Percent composition

CORE RULE

Mass % of X = (atomic mass of X × subscript) / molar mass × 100%. All element percents sum to 100.00% (± 0.05% rounding).

Empirical vs molecular formula

CORE RULE

Empirical formula = smallest whole-number atom ratio. Molecular formula = actual atoms = (empirical)ₙ, where n is a positive integer (n = 1 if they match).

Empirical formula from % or mass

METHOD

Assume 100 g (so % = g); convert each element to moles; divide all by the smallest; if a ratio isn't near an integer, multiply all by 2 (for .5), 3 (.33/.67), or 4 (.25/.75).

Combustion analysis

METHOD

All C → CO₂ (mass C = mol CO₂ × 12.01), all H → H₂O (mass H = 2 × mol H₂O × 1.008). For C-H-O, mass O = sample − C − H (by difference).