Required and core practicals

8.1 · Free notes

Practical: preparing a standard solution and carrying out an acid-base titration

Key points

• A standard solution is a solution of accurately known concentration, measured in mol dm⁻³. It is made from a primary standard: a pure, stable, soluble solid with a known formula and a fairly high molar mass that is not hygroscopic.
• To make one, first weigh by difference: weigh the container with the solid, tip the solid out, then reweigh the container. Record the mass to the same number of decimal places each time.
• Dissolve the solid in a beaker with less than the final volume of distilled water, stirring with a glass rod. Transfer the solution to a volumetric flask through a funnel, rinsing the beaker, rod and funnel into the flask so that no solute is lost.
• Add distilled water until the bottom of the meniscus sits on the graduation mark, adding the last few cm³ drop by drop. Stopper the flask and invert several times to mix. Never warm the flask to speed up dissolving, because the volume changes with temperature.
• Amount of substance: n = c × V with V in dm³ (V in cm³ ÷ 1000). Mass needed = n × Mr. Example: 250.0 cm³ of 0.0400 mol dm⁻³ Na₂CO₃ needs 0.0100 mol × 106.0 g mol⁻¹ = 1.06 g.
• Titration apparatus: the burette and pipette are rinsed with the solution they will contain. The conical flask is rinsed with distilled water only, because water does not change the moles of solute in it. Remove the funnel from the burette before reading, and fill the jet below the tap so there is no air bubble.
• Use a pipette (e.g. 25.0 cm³) to measure the solution in the flask. Add a few drops of indicator. Run in the titrant from the burette while swirling, over a white tile, adding it drop by drop near the end-point. Read the burette at eye level at the bottom of the meniscus.
• Indicator choice: methyl orange (red in acid, yellow in alkali, orange at the end-point) suits strong acid–weak base and strong–strong titrations. Phenolphthalein (colourless in acid, pink in alkali) suits weak acid–strong base and strong–strong titrations. Use a single indicator, and never one that changes colour gradually over a wide pH range.
• Do a rough titration first, then repeat until you have two concordant titres, meaning within 0.10 cm³ of each other. The mean titre uses concordant results only. Never include the rough titre or an outlier.
• Titration calculation: (1) n(known) = c × V; (2) use the mole ratio from the balanced equation; (3) c(unknown) = n ÷ V. For H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O the ratio is 1 : 2. Quote answers to 3 significant figures, with units.
• Uncertainty: the burette has ±0.05 cm³ per reading, so a titre (two readings) has ±0.10 cm³. Percentage uncertainty = (uncertainty ÷ measured value) × 100, so 0.10 ÷ 24.00 × 100 = 0.42%. Using a larger sample volume or a more concentrated titrant reduces the percentage uncertainty.
• Common errors: overshooting the end-point (titre too large); an air bubble in the jet that disappears (titre too large); leaving the funnel in the burette; rinsing the flask with the solution (extra moles); and reading the top of the meniscus.

In more detail

• Worked standard solution: 250 cm³ of 0.100 mol dm⁻³ Na₂CO₃ needs 0.0250 mol × 106 = 2.65 g of Na₂CO₃.
• Worked titration with Na₂CO₃: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂. 25.0 cm³ of 0.100 Na₂CO₃ (2.50 × 10⁻³ mol) needs 5.00 × 10⁻³ mol HCl; titre 24.00 cm³ → [HCl] = 0.208 mol dm⁻³.
• Effect of errors: water left in the burette dilutes the titrant (titre too large); an unrinsed pipette dilutes the sample (titre too small).
• Swirling and white tile: swirling mixes the reactants; the white tile makes the colour change at the end point easier to see.