Lesson 14.3.1

14.3.1 Redox titration calculations Quiz: Pearson Edexcel Chemistry, Unit 14

20 questions

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Lesson 14.3.1, Redox titration calculations: 20 multiple choice questions for the Pearson Edexcel Chemistry (9CH0), Unit 14: Redox II, written with Revision Ninja.

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The 20 questions

  1. In the titration of Fe2+ with MnO4- in acid, what is the mole ratio of MnO4- to Fe2+?

    • 1:8
    • 1:1
    • 1:5
    • 5:1

  2. What is a redox titration?

    • A titration that measures pH using a single indicator only, in which the end point is found from the pH curve
    • A precipitation titration using silver nitrate only, where the end point is shown by a white solid forming
    • A titration in which a known oxidising or reducing agent reacts with the analyte, with a colour change marking the end point
    • A titration in which the volume of gas evolved is measured at room temperature and pressure for each addition

  3. Why is potassium manganate(VII) described as self-indicating?

    • A faint pink colour persists at the end point because MnO4- is in slight excess
    • It becomes colourless at the end point because Mn2+ is colourless and no MnO4- remains
    • It needs starch to show a blue-black colour at the end point
    • It turns green at the end point because chromium(III) forms

  4. Why is starch used as an indicator in iodine-thiosulfate titrations?

    • Starch turns red with iodine, showing the end point
    • Starch is only used in permanganate titrations
    • Starch forms a blue-black complex with iodine
    • Starch forms a precipitate with thiosulfate

  5. What is the mole ratio of I2 to S2O3^2- in the iodine-thiosulfate titration?

    • 1:2
    • 1:4
    • 1:1
    • 2:1

  6. Why is dilute sulfuric acid used in permanganate titrations rather than hydrochloric acid?

    • It supplies H+ ions for the MnO4- reduction and is not oxidised, whereas hydrochloric acid would be oxidised
    • It precipitates Mn2+ ions as an insoluble sulfate so that they are removed before the end point
    • It oxidises Fe2+ to Fe3+ before the titration begins, so the titre measured is always too small
    • It raises the pH to 14 so that the MnO4- is stabilised and reacts more slowly with the analyte

  7. What is the first step in converting a titre volume to the moles of analyte?

    • Multiply the moles of titrant by the mole ratio from the balanced equation
    • Add the titre volume to the molar mass of the analyte to find the mass of analyte in the flask
    • Take the square root of the moles of titrant and then multiply by the volume of the analyte
    • Divide the moles of titrant by the molar mass of water to find the concentration of the analyte

  8. What colour change marks the end point when Fe2+ is titrated with MnO4- in acid?

    • Purple to colourless, with the colour disappearing at the end point of the titration
    • Blue-black to colourless, as the starch complex is destroyed by the excess titrant
    • Yellow to orange, as chromium(VI) is reduced by the iron(II) during the titration
    • Colourless to a permanent faint pink

  9. 25.0 cm3 of Fe2+ solution needs 22.50 cm3 of 0.0200 mol dm-3 MnO4- for complete reaction. What is the concentration of Fe2+?

    • 0.00900 mol dm-3
    • 0.0180 mol dm-3
    • 0.0900 mol dm-3
    • 0.450 mol dm-3

  10. 20.0 cm3 of 0.100 mol dm-3 Na2S2O3 is needed to titrate iodine. How many moles of iodine react?

    • 5.00 x 10^-4 mol
    • 2.00 x 10^-3 mol
    • 1.00 x 10^-3 mol
    • 4.00 x 10^-3 mol

  11. What is the mass of Fe2+ in 25.0 cm3 of a solution of concentration 0.0900 mol dm-3? (Ar Fe = 55.8)

    • 1.26 g
    • 0.126 g
    • 0.0126 g
    • 0.0403 g

  12. A 25.0 cm3 aliquot of Fe2+ solution needs 18.60 cm3 of 0.0200 mol dm-3 MnO4-. How many moles of Fe2+ are in the aliquot?

    • 1.86 x 10^-3 mol
    • 9.30 x 10^-3 mol
    • 7.44 x 10^-4 mol
    • 3.72 x 10^-4 mol

  13. How many electrons does one MnO4- ion gain when it is reduced to Mn2+ in acid?

    • 5 electrons, as Mn falls from +7 to +2
    • 8 electrons, because eight H+ ions are used
    • 3 electrons, as Mn falls from +7 to +4
    • 1 electron, as Mn falls from +7 to +6

  14. How many electrons does one S2O3^2- ion lose when it is oxidised to S4O6^2-?

    • 4 electrons per thiosulfate ion
    • 1 electron per thiosulfate ion
    • 2 electrons per thiosulfate ion
    • 0.5 electrons per thiosulfate ion

  15. A 0.300 g sample containing Fe2+ needs 22.00 cm3 of 0.0200 mol dm-3 MnO4- for titration. What is the percentage of iron by mass? (Ar Fe = 55.8)

    • 40.9%
    • 81.8%
    • 4.09%
    • 12.3%

  16. Why would a sample of Fe2+ that was partly oxidised by air before titration give a result that is too low?

    • Air adds extra Fe2+ to the solution, so the result is too high
    • Some Fe3+ was formed, so more MnO4- is needed and the result is too high
    • Oxidation does not change the amount of Fe2+, so the result is unchanged
    • Some Fe2+ was lost to oxidation

  17. Why is starch added only near the end point in an iodine-thiosulfate titration?

    • Starch reacts with thiosulfate ions, making the titre smaller because the thiosulfate is used up early
    • Adding starch early has no effect on the end point, so the timing of addition is not important at all
    • Starch added early traps iodine in a complex that releases it slowly, blurring the end point
    • Starch increases the iodine concentration in the flask and so reduces the titre measured at the end point

  18. What are the moles of H+ needed to react with 15.00 cm3 of 0.0200 mol dm-3 MnO4- in acid?

    • 2.40 x 10^-3 mol
    • 1.50 x 10^-3 mol
    • 5.00 x 10^-3 mol
    • 3.00 x 10^-4 mol

  19. 25.00 cm3 of Fe2+ is titrated with 30.00 cm3 of 0.0100 mol dm-3 MnO4-. What is the concentration of Fe2+?

    • 0.0600 mol dm-3
    • 0.0120 mol dm-3
    • 0.0300 mol dm-3
    • 0.300 mol dm-3

  20. A student uses hydrochloric acid instead of sulfuric acid in a permanganate titration and gets a titre that is too large. Why?

    • HCl reduces MnO4-, so the titre is too small
    • HCl has no effect because it is a strong acid
    • HCl forms a precipitate, so the titre is too small
    • HCl is oxidised by MnO4-, which consumes extra titrant

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