• Core Practical 13a: Redox titration between iron(II) and manganate(VII) — self-indicating, no indicator needed.
• Core Practical 13b: Redox titration between iodine and thiosulfate — uses starch as an indicator near the end point.
• Core Practical 12: Preparing a transition metal complex — tetraamminecopper(II) sulfate crystals — and calculating percentage yield.
• Core Practical 15: Qualitative analysis — identifying unknown positive ions (Group 2 metals, ammonium), negative ions (halides, hydroxide, carbonate, sulfate), and organic functional groups (alkenes, alcohols, carbonyls, carboxylic acids).
• Core Practical 16: Preparing aspirin from salicylic acid, then purifying it by recrystallisation and checking its purity with a melting point test.

Think of a titration as a chemical "weighing scale." You know the exact concentration of one solution (the one in the burette), and you slowly add it to a known volume of the other solution until the reaction is exactly complete. The volume you needed tells you how much of the unknown substance was there.

In a redox titration, instead of an acid neutralising a base, you have an oxidising agent reacting with a reducing agent — electrons are being transferred from one species to the other. The clever part is that many transition metal ions change colour when they change oxidation state, so you often don't even need to add a separate indicator — the reaction indicates itself.

Why Sulfuric Acid, and Nothing Else?

In this titration, manganate(VII) ions (MnO₄⁻) are the oxidising agent — they get reduced to Mn²⁺. Iron(II) is the reducing agent — it gets oxidised to Fe³⁺. This reaction only works in acidic conditions, so excess acid is added to the iron(II) solution before you even start titrating.

But you can't just grab any acid off the shelf. The acid has one job: provide H⁺ ions, and stay completely out of the redox chemistry. Dilute sulfuric acid is the only one that does this cleanly:

Spotting the End Point

Potassium manganate(VII) is intensely purple. As you drip it into the (acidified) iron(II) solution, it reacts instantly and is decolourised — the Mn²⁺ produced is such a pale pink that the solution just looks colourless. This continues until every Fe²⁺ ion has reacted. The very next drop of manganate(VII) has nothing left to react with, so it survives, and the whole flask turns a persistent pale pink. That's your end point.

Building and Balancing the Equation

You're given two half-equations and asked to combine them. The rule is always the same: multiply each half-equation so the number of electrons matches, then add them together and cancel the electrons.

Notice the mole ratio buried in that equation: 1 mole of MnO₄⁻ reacts with 5 moles of Fe²⁺. Every calculation in this practical hinges on that 1:5 ratio.

Worked Example — Analysing an Iron Tablet

This is a two-stage idea, and it's easy to get lost if you only look at the final titration equation. Here's the full logic chain:

1. You have an oxidising agent of unknown concentration (e.g. chlorate(I) ions in bleach).
2. You react it with an excess of iodide ions — the oxidising agent converts iodide into iodine (I₂). The amount of iodine produced is directly linked to how much oxidising agent was present.
3. You then titrate that iodine against a sodium thiosulfate solution of known concentration. The volume of thiosulfate needed tells you exactly how much iodine was made — and working backwards, how much oxidising agent you started with.

Reading the Colour Changes

Iodine solution starts off light brown/yellow. As thiosulfate is added from the burette, the iodine is steadily converted to colourless iodide ions, and the solution gets paler and paler. Once it fades to a pale straw colour, that's your cue to add starch indicator — not at the start, because starch forms such an intense blue-black complex with iodine that it would mask the gradual colour change and make it impossible to judge when you're getting close to the end point.

Worked Example — Chlorate(I) in Household Bleach

This practical makes crystals of tetraamminecopper(II) sulfate-1-water, Cu(NH₃)₄SO₄·H₂O, starting from ordinary copper(II) sulfate. The whole thing is really a story about solubility — you dissolve something, add a reagent to build a new, different compound in solution, and then force it back out as a solid by making the solvent "unfriendly" to it.

The Method, Step by Step

1. Weigh 1.4–1.6 g of copper(II) sulfate accurately (weigh test tube empty, then with the solid — the difference is the mass).
2. Dissolve it in 4 cm³ of water, warming gently in a water bath.
3. In the fume cupboard, wearing gloves, add 2 cm³ of concentrated ammonia solution while stirring — the solution turns deep blue as the [Cu(NH₃)₄(H₂O)₂]²⁺ complex ion forms.
4. Pour this into 6 cm³ of ethanol and cool in an ice bath — dark blue Cu(NH₃)₄SO₄·H₂O crystals precipitate out.
5. Filter using a Büchner funnel (vacuum filtration — much faster than gravity filtration), wash with cold ethanol.
6. Dry the crystals between filter paper, then weigh them.

Calculating Percentage Yield

This is a classic "actual vs. theoretical" calculation. You work out the maximum mass you could have made (assuming a perfect 1:1 reaction with no losses), then compare it to what you actually collected.

Worked Example

Safety in This Practical

This practical is a toolkit of "detective tests." You're handed an unknown substance, and each test gives you a clue — a colour change, a precipitate, a smell, a gas — that lets you rule ions and functional groups in or out. The trick to mastering this section isn't memorising every single line; it's understanding why each reagent produces the result it does, so you can reconstruct the table under exam pressure even if your memory blanks.

Testing for Positive Ions: Group 2 Metals

Add sodium hydroxide (or ammonia solution, or sulfuric acid) dropwise, in excess, to a solution of the unknown metal ion, and watch for a precipitate.

Testing for Ammonium Ions (NH₄⁺)

Add sodium hydroxide to the unknown solution, then gently warm it in a water bath. Ammonium ions react with hydroxide to release ammonia gas: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l). Hold a piece of damp red litmus paper near (not in) the mouth of the tube using tongs.

Testing for Negative Ions

Halide Ions (Cl⁻, Br⁻, I⁻)

Add dilute nitric acid (to remove interference from other ions like carbonate), then add silver nitrate solution.

Hydroxide Ions (OH⁻)

Test the pH directly with red litmus paper or universal indicator. A positive result: red litmus turns blue, or the universal indicator paper reads clearly alkaline.

Carbonate Ions (CO₃²⁻)

Add dilute hydrochloric acid, then immediately attach a bung and delivery tube leading into a second test tube containing limewater (calcium hydroxide solution).

Sulfate Ions (SO₄²⁻)

Acidify the sample with dilute hydrochloric acid first, then add a few drops of aqueous barium chloride.

Testing for Organic Functional Groups

Alkenes / Unsaturation — Bromine Water Test

Shake the unknown compound with orange/yellow bromine water. If a carbon–carbon double bond is present, an addition reaction occurs and the bromine is used up.

Alcohols — PCl₅ Test

Add solid phosphorus(V) chloride. If an −OH group is present, a vigorous reaction occurs at room temperature, with no heating needed.

Distinguishing 1° / 2° Alcohols from Tertiary Alcohols — Oxidation Test

Warm the unknown with acidified potassium dichromate(VI), or acidified potassium manganate(VII).

Carbonyls (General) — 2,4-DNPH Test

2,4-dinitrophenylhydrazine (2,4-DNPH) undergoes a condensation reaction with the carbonyl group (C=O) in both aldehydes and ketones.

Distinguishing Aldehydes from Ketones

Both Tollens' reagent and Fehling's solution rely on the same underlying idea: aldehydes can be oxidised further (to a carboxylic acid); ketones cannot.

Carboxylic Acids

Add solid sodium carbonate, or aqueous sodium hydrogen carbonate.

This practical brings together several separate skills you've built up across the chapter: accurate measurement, controlled heating, purification by recrystallisation, and using a melting point to check purity. It's really two experiments glued together — first you make the aspirin, then you purify and verify it.

Stage 1 — Synthesis

1. Add 6.0 g of salicylic acid to a conical flask with 10 cm³ of ethanoic anhydride and 5 drops of concentrated sulfuric acid (a catalyst).
2. Swirl and hold in a warm water bath at around 60°C for about 20 minutes.
3. Cool the flask, then pour the contents into 75 cm³ of cold water — the aspirin crystallises out.
4. Recover the crude aspirin by Büchner (vacuum) filtration, and leave to dry.

Stage 2 — Recrystallisation (Purification)

The crude aspirin you've collected is impure — it contains left-over reactants and by-products. Recrystallisation is the standard technique for purifying an organic solid, and the logic is elegant:

1. Dissolve the impure solid in the minimum volume of hot solvent (here, ethanol) needed to fully dissolve it.
2. If any insoluble impurities remain, filter the hot solution (a "hot filtration").
3. Let the solution cool slowly to room temperature — the desired product crystallises out, but the impurities (present in much smaller amounts) stay dissolved in the solvent.
4. Filter again (typically Büchner filtration, since it's faster — filtration under reduced pressure) to collect the pure crystals.
5. Wash with a small amount of fresh, cold solvent to remove any remaining surface impurities, then dry.

Stage 3 — Melting Point Analysis

The melting point of a solid is a fingerprint for both its identity and its purity. Pure aspirin melts sharply at 135°C.

Practical Skills for a Good Melting Point Test

• The sample must be completely dry and finely powdered — crush it with the back of a spatula on filter paper or a white tile to absorb any remaining moisture.
• Do a quick first run to find the approximate melting range, heating fairly fast.
• Repeat with a much slower heating rate for an accurate reading — heating too fast means the thermometer reading lags behind the true temperature of the sample, giving you an inaccurately high value.
• Take repeat readings (three is standard) and quote the result as a range, ideally referenced against a data book value.

• I can explain why dilute sulfuric acid — and no other common acid — is used in the manganate(VII)/iron(II) titration.
• I can describe how potassium manganate(VII) acts as its own indicator, and describe the colour change at the end point.
• I can balance and combine two half-equations into a full redox equation.
• I can carry out a full redox titration calculation, from titre volume to final concentration or percentage purity.
• I understand why starch indicator is added late (not at the start) in an iodine–thiosulfate titration.
• I can describe the method for preparing tetraamminecopper(II) sulfate crystals, including safety precautions.
• I can explain, using solubility, why adding ethanol causes the copper complex to crystallise.
• I can calculate theoretical yield and percentage yield, and explain why a >100% yield usually means undried crystals.
• I can identify Group 2 metal ions from their reactions with NaOH, ammonia, and sulfuric acid.
• I can test for ammonium ions and describe the positive result.
• I can identify halide, hydroxide, carbonate, and sulfate ions, including why HCl is added before BaCl₂ in the sulfate test.
• I can distinguish chloride, bromide, and iodide precipitates by colour and by their reaction with ammonia.
• I can describe tests for alkenes, alcohols, carbonyls, and carboxylic acids, and state the positive result for each.
• I can distinguish primary/secondary alcohols from tertiary alcohols using an oxidising agent.
• I can distinguish aldehydes from ketones using Tollens' reagent and Fehling's solution, and explain the underlying chemistry.
• I can describe the synthesis of aspirin from salicylic acid, including the role of concentrated sulfuric acid.
• I can explain the principle of recrystallisation and why the minimum volume of hot solvent must be used.
• I can explain how melting point range and value indicate the purity of a product.