The mole is the chemist’s counting unit. One mole of any substance contains 6.02 × 10²³ particles (Avogadro’s number). This lets us count atoms and molecules by weighing them.
Key Definitions
- Relative atomic mass (Aᵣ): the mass of one atom compared to 1/12 the mass of a carbon-12 atom.
- Molar mass: the mass of one mole of a substance, in g/mol. Numerically equal to the relative molecular mass (Mᵣ).
The Mole Formula
Also: number of particles = moles × 6.02 × 10²³.
Concentration of Solutions
- Concentration (mol/dm³) = moles ÷ volume (dm³)
- 1 dm³ = 1000 cm³.
Worked Examples
Example 1: Calculate the molar mass of sodium carbonate, Na₂CO₃ (Na = 23, C = 12, O = 16).
Mᵣ = (2 × 23) + 12 + (3 × 16) = 46 + 12 + 48 = 106 g/mol.
Example 2: How many moles are in 8 g of NaOH (Na = 23, O = 16, H = 1)?
Mᵣ of NaOH = 23 + 16 + 1 = 40 g/mol.
Moles = 8 ÷ 40 = 0.2 mol.
Example 3: How many moles are in 2.0 dm³ of a 0.5 mol/dm³ solution?
Moles = concentration × volume = 0.5 × 2.0 = 1.0 mol.
Using Moles in Equations
Balanced equations give the mole ratio of reactants and products. For example, in 2H₂ + O₂ → 2H₂O, the ratio H₂ : O₂ : H₂O is 2 : 1 : 2. Use this ratio to calculate reacting masses.
Summary
One mole = 6.02 × 10²³ particles. Moles = mass ÷ molar mass, and concentration = moles ÷ volume. Balanced equations provide mole ratios for calculating the masses of reactants and products.
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