B. Precipitation method

• The analyte is precipitated in the form of a pure, sparingly soluble compound of constant, known chemical composition.
• The precipitate is separated by filtration through ashless filter paper (a paper which leaves no ash when burnt), the paper with the precipitate is transferred to an ignition crucible and ignited until the paper is completely burnt off (its constituents volatilize), and the mass of the precipitate that remains is weighed and used to find the mass of the analyte.
• Example: barium precipitated as BaSO₄.

Worked example. Excess sodium sulphate is added to a solution containing BaCl₂; the BaSO₄ precipitate weighs 1.165 g (Ba = 137, S = 32, O = 16, Cl = 35.5). BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl 208 g BaCl₂ → 233 g BaSO₄ x → 1.165 g ⇒ x = 208 × 1.165/233 = 1.04 g BaCl₂.

Worked example (percentage of chloride). 1.5 g of impure NaCl gives 2.87 g AgCl (Ag = 108, Cl = 35.5): moles AgCl = 2.87/143.5 = 0.02 mol, so 0.02 mol Cl⁻ = 0.71 g Cl; percentage of chlorine = 0.71/1.5 × 100 = 47.3 %. (Mass of NaCl = 0.02 × 58.5 = 1.17 g, i.e. 78 % purity.)

Worked example (sulphate in a sample). A 0.60 g sample containing sodium sulphate is dissolved and excess BaCl₂ added; the dried precipitate of BaSO₄ weighs 0.699 g. Moles BaSO₄ = 0.699/233 = 0.00300 mol. Sulphate present = 0.00300 × 96 = 0.288 g. Mass of Na₂SO₄ = 0.00300 × 142 = 0.426 g; purity = 0.426/0.60 × 100 = 71.0 %.

Worked example (iron as the oxide). The iron in a sample is precipitated as Fe(OH)₃ and ignited to Fe₂O₃, which weighs 0.800 g (Fe = 56, O = 16). Fe₂O₃ (160 g) contains 112 g Fe, so the mass of iron = 0.800 × 112/160 = 0.560 g.