A mole is a counting unit, like a dozen, but much larger: 1 mol = 6.022 × 10²³ particles (Avogadro's number). It lets chemists connect the mass they can weigh to the number of atoms or molecules that react.
Molar mass
Add the atomic masses from the periodic table. Water, H₂O: 2(1.008) + 16.00 = 18.02 g/mol. Oxygen gas, O₂: 2(16.00) = 32.00 g/mol.
The mole map
- grams ÷ molar mass = moles
- moles × 6.022 × 10²³ = particles
- moles of A × (coefficient of B / coefficient of A) = moles of B
Example 1: grams to molecules
How many molecules are in 36.04 g of water? 36.04 g ÷ 18.02 g/mol = 2.000 mol, then 2.000 × 6.022 × 10²³ = 1.204 × 10²⁴ molecules.
Example 2: predicting product mass
For 2H₂ + O₂ → 2H₂O, how much water forms from 4.0 g of H₂ with plenty of oxygen?
- Moles of H₂:
4.0 ÷ 2.016 = 1.98 mol - Mole ratio H₂O : H₂ is 2 : 2, so 1.98 mol of water.
- Mass:
1.98 × 18.02 ≈ 36 gof water (two significant figures).
Example 3: the limiting reactant
Now react 4.0 g of H₂ with only 16.0 g of O₂. Oxygen: 16.0 ÷ 32.00 = 0.500 mol. The 1.98 mol of H₂ would need 1.98 ÷ 2 = 0.99 mol of O₂, but only 0.500 mol is available, so O₂ is the limiting reactant. Water formed: 0.500 × 2 = 1.000 mol = 18.0 g. The leftover hydrogen is the excess reactant.
Common mistakes
- Using grams directly in a mole ratio. Ratios only work in moles.
- Starting from an unbalanced equation.
- Assuming the reactant with the smaller mass is limiting.