Laboratory Techniques and Practical Chemistry
What this domain tests
Practical items test whether you can choose the right apparatus and separation method, carry out a titration calculation, purify a solid, and interpret the direction of an experimental error.
Core rules
- Volumetric glassware for accurate volumes: burette (±0.05 cm³), pipette (25.0 cm³), volumetric flask; measuring cylinders are for approximate volumes only.
- Titration: rinse the burette with the titrant, the pipette with the analyte; read the bottom of the meniscus at eye level; a rough titration locates the end point before three concordant accurate runs within 0.10 cm³.
- Indicator choice: phenolphthalein (8.3–10.0) for strong base–weak acid; methyl orange (3.1–4.4) for strong acid–weak base; either for strong acid–strong base.
- Separation: filtration for insoluble solids, recrystallisation for purifying solids (choose a solvent in which the compound is soluble hot and insoluble cold), simple distillation for a volatile liquid from non-volatile impurities, fractional distillation for liquids with close boiling points, chromatography for very small amounts or similar compounds.
- Errors: systematic errors shift every result the same way (miscalibrated balance, impure standard); random errors scatter and are reduced by repetition.
Worked example
25.0 cm³ of 0.100 mol dm⁻³ NaOH is titrated with HCl and requires 23.45 cm³ to reach the phenolphthalein end point. Moles NaOH = 0.100 × 0.0250 = 2.50 × 10⁻³ mol. The 1:1 equation gives the same moles of HCl, so [HCl] = 2.50 × 10⁻³/0.02345 = 0.1066 mol dm⁻³, reported as 0.107 mol dm⁻³ to three significant figures.
Traps that cost marks
- Using a measuring cylinder for the analyte - it introduces avoidable random error.
- Rinsing the burette with water rather than titrant, which dilutes the solution and raises the titre.
- Recrystallising with too much hot solvent, which lowers recovery rather than purity.
- Reporting a titre to more decimal places than the burette can deliver.