Skip to content
Dabour HubLEARN · PRACTICE · EXPLORESign inStudent

6.2Unit 6. Practical Skills in Chemistry II

Advanced Inorganic & Organic — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

5 min read 9 parts Lesson 3 of 14 1 quiz

📚 IAL Edexcel Linear (Chemistry) Hub / Topical Revision
🩺
100% Curriculum Specification Match Lesson Doctor Verified
The only platform with 100% curriculum specification match and verified interactive question banks.
0%read
Lesson summary

6.2 Redox Titrations, Complex Ions and Aspirin (Core Practicals 12, 13 and 16) — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Carry out and interpret manganate(VII)/iron(II) and thiosulfate/iodine titrations, relate complex ion colours to ligands and describe the preparation of aspirin.…

  • Indicators
  • Two common titrations
  • Oxidising agent
  • Reducing agent
  • Half equations
  • Overall equation

Part 1 of 9Redox titration principles

Redox titration
  • Redox titration an oxidising agent is titrated against a reducing agent and electrons transfer between them.
Indicators
  • Indicators often unnecessary because transition metal ions change colour when their oxidation state changes.
Two common titrations
  • Two common titrations manganate(VII) with iron(II), and thiosulfate with iodine.

Part 2 of 9Core Practical 13a: manganate(VII) with iron(II)

Oxidising agent
  • MnO4− is reduced to Mn2+.
Reducing agent
  • Fe2+ is oxidised to Fe3+.
Half equations
  • MnO4− + 8H+ + 5e− → Mn2+ + 4H2O and Fe2+ → Fe3+ + e−.
Overall equation
  • MnO4− + 8H+ + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+, a 1 : 5 mole ratio.
Uses
  • Uses finding the percentage of iron in iron tablets, and the formula of hydrated ethanedioic acid (MnO4− : C2O42− is 2 : 5, warmed to about 60 °C).

Part 3 of 9Choice of acid and end point in 13a

Acid used
  • Acid used excess dilute sulfuric acid is added to the Fe2+ before titrating because it does not react with manganate(VII).
Hydrochloric acid
  • Cl− is oxidised to chlorine, giving too large a titre.
  • Nitric acid — is itself an oxidising agent.
  • Ethanoic acid — is too weak to supply enough H+.
Insufficient acid
  • Insufficient acid brown MnO2 forms instead of Mn2+.
Self-indicating
  • Self-indicating purple MnO4− becomes almost colourless Mn2+, and the end point is the first permanent pale pink.
Reading the burette
  • Reading the burette read the top of the meniscus because the solution is dark, and use a burette with white numbering.

Part 4 of 9Procedure for 13a

1
Weigh the iron(II) sample accurately and dissolve it in dilute sulfuric acid.
2
Make up to 250.0 cm3 in a volumetric flask and invert to mix.
3
Pipette 25.0 cm3 into a conical flask and add excess dilute H2SO4.
4
Titrate with standard KMnO4 until the first permanent pale pink colour.
5
Repeat to concordant titres within 0.10 cm3.
6
Calculate moles of Fe2+ from the 1 : 5 ratio and scale up to the whole sample.

Part 5 of 9Core Practical 13b: thiosulfate with iodine

Reaction
  • I2 + 2S2O32− → 2I− + S4O62−, a 1 : 2 mole ratio.
Purpose
  • Purpose determines an oxidising agent that liberates iodine from excess iodide, such as Cu2+, IO3−, ClO− or H2O2.
Liberation equations
  • 2Cu2+ + 4I− → 2CuI + I2, IO3− + 5I− + 6H+ → 3I2 + 3H2O and ClO− + 2I− + 2H+ → Cl− + I2 + H2O.
1
Add the oxidising agent to excess KI so brown iodine is released.
2
Titrate with thiosulfate until the solution is pale straw or yellow.
3
Add starch, giving a blue-black colour.
4
Continue until the blue-black colour just disappears.
5
Use the 1 : 2 ratio to calculate the amount of oxidant.
  • Starch timing
    • added only near the end point, because with a lot of iodine it forms an insoluble complex.

Part 6 of 9Sequence for redox titration calculations

1
Write the half equations for the oxidant and the reductant.
2
Deduce the overall equation.
3
Calculate moles of the titrant used.
4
Find the mole ratio from the overall equation.
5
Calculate moles of the other reagent in the sample.
6
Scale to the original solution.
7
Give the concentration or the percentage by mass.
  • Accuracy
    • use pipettes and burettes, record to ±0.05 cm3 and calculate percentage uncertainties.

Part 7 of 9Complex ions and colour

Complex ion
  • Complex ion a central metal ion surrounded by ligands bonded by dative covalent bonds.
Ligands
  • H2O, NH3, Cl− and OH− donate lone pairs to the metal ion.
Colour change
  • Colour change changing the ligand, coordination number or oxidation state changes the colour.
Copper(II) example
  • [Cu(H2O)6]2+ + 4Cl− ⇌ [CuCl4]2− + 6H2O, blue to yellow-green.
Ion tests with NaOH(aq)
  • Cu2+ blue precipitate, Fe2+ green, Fe3+ orange-brown, Cr3+ grey-green and Mn2+ pale brown darkening in air.
    • (dissolving in excess to dark green)
Excess ammonia
  • Excess ammonia dissolves Cu(OH)2 to the deep blue [Cu(NH3)4(H2O)2]2+.

Part 8 of 9Core Practical 16: preparing aspirin

  • Reaction
    • 2-hydroxybenzoic acid (salicylic acid) reacts with ethanoic anhydride, with ethanoic acid as by-product.
1
Mix 6.0 g salicylic acid, 10 cm3 ethanoic anhydride and 5 drops of concentrated sulfuric acid in a conical flask.
2
Warm in a water bath at about 60 °C for 20 minutes.
3
Cool, then pour into 75 cm3 cold water so the crude aspirin crystallises.
4
Collect by Buchner filtration under reduced pressure.
5
Recrystallise from the minimum hot ethanol, pour over 40 cm3 cold water, filter and dry.
6
Test purity by melting point.
Choice of reagent
  • Choice of reagent ethanoic anhydride is cheaper, less corrosive and does not release HCl, unlike ethanoyl chloride.
Purity
  • Purity pure aspirin melts sharply at about 135 °C, and TLC comparison with pure aspirin also checks purity.

Part 9 of 9More from the specification

1Precipitation of insoluble solid
  • Precipitation of insoluble solid mixing aqueous reagents forms an insoluble precipitate that separates readily as a crystalline or amorphous solid.
2Positive result
  • The observation that shows a functional group is present.
3Carbonate
  • A compound containing the CO3 group, such as calcium carbonate.
4Carbon dioxide
  • A gas that turns limewater cloudy.
5Secondary alcohol
  • An alcohol whose –OH carbon is bonded to two other carbons.
6Slow cooling
  • Slow cooling gives larger, well-formed crystals that are easier to filter and dry.
7Adding ethanol
  • Adding ethanol lowers the solubility of the complex salt in the solvent, so crystals form.
8Percentage yield
  • Percentage yield actual mass obtained ÷ theoretical mass × 100.

Key terms

TermMeaning
Redox titrationA titration in which titrant and analyte exchange electrons.
Self-indicatingA titrant whose own colour change shows the end point, such as MnO4− becoming Mn2+.
Standard solutionA solution whose concentration is accurately known.
Iodometric titrationA method in which an oxidising agent liberates iodine from excess I− that is titrated with thiosulfate.
Starch indicatorForms a blue-black complex with iodine and is added near the end point.
Complex ionA central metal ion surrounded by ligands bonded by dative covalent bonds.

Avoid these mistakes

1
Using hydrochloric acid to acidify manganate(VII) — Cl− is oxidised to Cl2 so the titre is too high.
2
Adding starch at the start of the thiosulfate titration — add it near the end point when the solution is pale yellow.
3
Forgetting excess acid in 13a — without it brown MnO2 forms.
4
Using the wrong mole ratio — MnO4− : Fe2+ is 1 : 5 and I2 : S2O32− is 1 : 2.

Last updated:

\n