Part 4: Solutions, dilution and Core Practical 2

Distinguish the amount of solute from its concentration, prepare a solution of known concentration and use it to standardise sodium hydroxide.

Concentration uses the final solution volume

Amount concentration is moles of solute per dm³ of solution: c = n/V. A concentration of 0.200 mol dm⁻³ does not mean a flask necessarily contains 0.200 mol; a 25.0 cm³ portion contains 0.200 × 0.0250 = 0.00500 mol. The final solution volume is used, not the initial volume of solvent added to dissolve the solid.

Dissolve 2.650 g anhydrous Na₂CO₃ and make the final solution volume 250.0 cm³. M = 106.0 g mol⁻¹, so n = 2.650/106.0 = 0.02500 mol and c = 0.02500/0.2500 = 0.1000 mol dm⁻³. Mass concentration = 2.650/0.2500 = 10.60 g dm⁻³. Multiplying 0.1000 mol dm⁻³ by 106.0 g mol⁻¹ gives the same answer.

Count ions, then conserve solute during dilution

Assuming complete dissolution, 0.100 mol dm⁻³ CaCl₂ supplies [Ca²⁺] = 0.100 mol dm⁻³ and [Cl⁻] = 0.200 mol dm⁻³. The two chloride ions per formula unit double the chloride amount without doubling the solution volume. Do not confuse the concentration of one ion with the total concentration of all ions.

During simple dilution no solute is created or lost, so c₁V₁ = c₂V₂. Pipetting 20.0 cm³ of 0.500 mol dm⁻³ solution into a 250.0 cm³ volumetric flask and making up to the mark gives c₂ = (0.500 × 20.0)/250.0 = 0.0400 mol dm⁻³. Volumes can share the same units in this ratio; if separately calculating moles, convert to dm³.

The final volume is 250.0 cm³, not 270.0 cm³. 'Make up to 250.0 cm³' differs from 'add 250.0 cm³ water'. In a volumetric preparation the calibration line defines the final solution volume, avoiding assumptions that solvent and solute volumes are exactly additive.

Core Practical 2: prepare a known acid solution

A standard solution has an accurately known concentration. A suitable solid standard has known purity and composition and is stable during weighing and preparation. Sodium hydroxide is unsuitable for simply weighing an exact amount of pure NaOH from a bench bottle: it absorbs water and reacts with carbon dioxide. Standardising its solution against a suitable acid determines its actual concentration.

Pearson's CP2 example uses sulfamic acid, NH₂SO₃H, M = 97.1 g mol⁻¹, a monoprotic solid acid. Measure the mass transferred, dissolve completely in a beaker, then transfer quantitatively to a volumetric flask. Rinse the beaker, rod and funnel into the flask so that measured solute is not left behind. If warmed to dissolve, cool to the flask's calibration temperature before making up the volume.

Add deionised water close to the mark, then dropwise until the bottom of the meniscus lies on the line at eye level. Stopper and invert repeatedly to mix throughout. The narrow neck makes a small volume change produce a visible meniscus movement; the bulb stores most of the solution. A beaker's graduations do not provide the same volume precision.

Use the acid to standardise NaOH

In the Pearson example, the sulfamic acid goes in the burette and a measured 25.0 cm³ NaOH aliquot goes in a conical flask. Add a few drops of methyl orange and deliver acid while swirling, dropwise near the end point. The flask changes from yellow through orange towards red as acid is added; use the specified end point consistently. Repeat to obtain concordant titres and calculate a mean from suitable concordant runs, excluding the rough result.

Because sulfamic acid is monoprotic, its reaction with NaOH is 1:1. In an original example, 2.4275 g acid made to 250.0 cm³ gives n = 2.4275/97.1 = 0.02500 mol and c = 0.1000 mol dm⁻³. A mean acid titre of 24.60 cm³ contains 0.1000 × 0.02460 = 0.002460 mol acid. The 25.0 cm³ alkali therefore contains 0.002460 mol NaOH: c(NaOH) = 0.002460/0.0250 = 0.0984 mol dm⁻³.

Use the acid titre volume to find the acid amount; use the pipetted alkali volume to find the alkali concentration. The two volumes perform different roles. A 1:1 reaction does not make their concentrations equal unless the reacting volumes are also equal.

Explain each preparation error through concentration

Failure to transfer all the weighed acid makes the true acid concentration lower than the calculated concentration. More acid volume is then required to neutralise the alkali; using the falsely high acid concentration with that titre overestimates NaOH concentration. Overfilling the volumetric flask has the same concentration effect. Do not attempt to fix an overfilled flask by removing mixed solution: that removes solute as well as water; prepare again.

Incomplete mixing gives aliquots with different concentrations and unpredictable errors. Rinsing the pipette with NaOH after water removes dilution of the measured alkali aliquot. Rinsing the conical flask with water is acceptable because it changes volume but not the moles already pipetted. Rinsing that flask with NaOH would add an unknown extra amount and increase the acid titre.

Wear goggles, use a safety pipette filler and fill a clamped burette below eye level. Avoid ingestion or skin contact with the solid acid. Match precautions to the supplied solutions and their hazard information. Practical records should retain actual mass, initial and final burette readings, units and observations, rather than just the final concentration.

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Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources