REDOX TITRATION ASSIGNMENT

Question 1: A piece of rusted iron was analysed to find out how much of the iron had been oxidised to rust [hydrated iron(III) oxide]. A small sample of the iron was dissolved in excess dilute sulfuric acid to give 250 cm³ of solution. The solution contains Fe²⁺ ions from the unrusted iron dissolving in the acid, and, Fe³⁺ ions from the rusted iron.

(a) 25.0 cm³ of this solution required 16.9 cm³ of 0.0200 mol dm^–3 KMnO₄ for complete oxidation of the Fe²⁺ ions.

Calculate the moles of Fe²⁺ ions in the sample titrated.

(b) To a second 25.00 cm³ of the rusted iron  solution an oxidising agent was added to convert all the Fe²⁺ ions present to Fe³⁺ ions. The Fe³⁺ ions were titrated with a solution of EDTA^4–_(aq) ions and 17.6 cm³ of 0.100 mol dm^–3 EDTA were required.

Assuming 1 mole of EDTA reacts with 1 mole of Fe³⁺ ions, calculate the moles of Fe³⁺ ions in the sample.

(c) From your calculations in (a) and (b) calculate the ratio of rusted iron to unrusted iron and hence the percentage of iron that had rusted.

QUESTION 2 Given the following two half–reactions

(a) Given (i) S₄O₆^2–_(aq) + 2e^– ==> 2S₂O₃^2–_(aq)

and (ii) I_2(aq) + 2e^– ==> 2I^–_(aq)

construct the full ionic redox equation for the reaction of the thiosulfate ion S₂O₃^2– and iodine I₂.

(b) what mass of iodine reacts with 23.5 cm³ of 0.0120 mol dm^–3 sodium thiosulfate solution.

(c) 25.0 cm³ of a solution of iodine in potassium iodide solution required 26.5 cm³ of 0.0950 mol dm^–3 sodium thiosulfate solution to titrate the iodine.

What is the molarity of the iodine solution and the mass of iodine per dm³?

Question 3: 2.83 g of a sample of haematite iron ore [iron (III) oxide, Fe₂O₃] were dissolved in concentrated hydrochloric acid and the solution diluted to 250 cm³.

25.0 cm³ of this solution was reduced with tin(II) chloride (which is oxidised to Sn⁴⁺ in the process) to form a solution of iron(II) ions.

This solution of iron(II) ions required 26.4 cm³ of a 0.0200 mol dm^–3 potassium dichromate(VI) solution for complete oxidation back to iron(III) ions.

(a) given the half–cell reactions 

(i) Sn⁴⁺_(aq) + 2e^– ==> Sn²⁺_(aq)

and (ii) Cr₂O₇^2–_(aq) + 14H⁺_(aq) + 6e^– ==> 2Cr³⁺_(aq) + 7H₂O_(l)

deduce the fully balanced redox equations for the reactions

(i) the reduction of iron(III) ions by tin(II) ions

(ii) the oxidation of iron(II) ions by the dichromate(VI) ion

(b) Calculate the percentage of iron(III) oxide in the ore.

(c) Suggest why potassium manganate(VII) isn’t used for this titration? (though it was ok in Q1)

If you don’t know, the following half-cell potential data will help!

E^θ = +1.33 for Cr₂O₇^2–_(aq) + 14H⁺_(aq) + 6e^– 2Cr³⁺_(aq) + 7H₂O_(l)

E^θ = +1.36 for Cl_2(aq) + 2e^– 2Cl^–_(aq)

E^θ = +1.51 for MnO₄^–_(aq) + 8H⁺_(aq) + 5e^– Mn²⁺_(aq) + 4H₂O_(l)

Question 4: An approximately 0.02 mol dm^–3 potassium manganate(VII) solution was standardized against precisely 0.100mol dm^–3 iron(II) ammonium sulfate solution. 25.0 cm³ of the solution of the iron(II) salt were oxidized by 24.15 cm³ of the manganate(VII) solution.

What is the molarity of the potassium manganate(VII) solution ?

Question 5: 10.0 g of iron(II) ammonium sulfate crystals were made up to 250 cm³ of acidified aqueous solution. 25.0 cm³of this solution required 21.25 cm³ of 0.0200 mol dm^–3 potassium dichromate(VI) for oxidation.

Calculate x in the formula FeSO₄.(NH₄)₂SO₄.xH₂O

Question 6: Given the half–reaction C₂O₄^2–_(aq) – 2e^– ==> 2CO_2(g)

or H₂C₂O_4(aq) – 2e^– ==> 2CO_2(g) + 2H⁺_(aq)

(a) write out the balanced redox equation for manganate(VII) ions oxidising the ethanedioate ion (or ethanedioic acid).

(b) 1.520 g of ethanedioic acid crystals, H₂C₂O₄.2H₂O, was made up to 250.0 cm³ of aqueous solution and 25.00 cm³ of this solution needed 24.55 cm³ of a potassium manganate(VII) solution for oxidation.

Calculate the molarity of the manganate(VII) solution and its concentration in g dm^–3.

Question 7: A standardization of potassium manganate(VII) solution yielded the following data:

0.150 g of potassium tetraoxalate dihydrate, KHC₂O₄.H₂C₂O₄.2H₂O needed 23.20 cm³ of the manganate(VII) solution.

What is the molarity of the manganate(VII) solution? Use the equation and reasoning from Q6 to help you.

Question 8: Given the half–cell equation  O_2(g) +2H⁺_(aq) + 2e^– ==> H₂O_2(aq)

(a) construct the fully balanced redox ionic equation for the oxidation of hydrogen peroxide by potassium manganate(VII)

(b) 50.0 cm³ of solution of hydrogen peroxide were diluted to 1.00 dm³ with water.

25.0 cm³ of this solution, when acidified with dilute sulfuric acid, reacted with 20.25 cm³ of 0.0200 mol dm^–3KMnO₄.

What is the concentration of the original hydrogen peroxide solution in mol dm^–3?

Question 9: 13.2 g of iron(III) alum were dissolved in water and reduced to an iron(II) ion solution by zinc and dilute sulfuric acid. The mixture was filtered and the filtrate and washings made up to 500 cm³ in a standard volumetric flask.

If 20.0 cm³ of this solution required 26.5  cm³ of 0.0100 mol dm^–3 KMnO₄ for oxidation.

(a) give the ionic equation for the reduction of iron(III) ions by zinc metal.

(b) Calculate the percentage by mass of iron in iron alum.

Question 10: Calculate the concentration in mol dm^–3 and g dm^–3, of a sodium ethanedioate (Na₂C₂O₄) solution, 5.00 cm³of which were oxidized in acid solution by 24.50 cm³ of a potassium manganate(VII) solution containing 0.05 mol dm^–3.

Question 11: Calculate x in the formula FeSO₄.xH₂O from the following data:

12.18 g of iron(II) sulfate crystals were made up to 500 cm³ acidified with sulfuric acid.

25.0 cm³ of this solution required 43.85 cm³ of 0.0100 mol dm^–3 KMnO₄ for complete oxidation.

Question 12:  Given the half–reaction NO₃^–_(aq) + 2H⁺_(aq) + 2e^– ==> NO₂^–_(aq) + H₂O_(l)

(a) give the ionic equation for potassium manganate(VII) oxidising nitrate(III) to nitrate(V)

(b) 24.2 cm³ of sodium nitrate(III) [sodium nitrite] solution, added from a burette, were needed to discharge the colour of25.0 cm³ of an acidified 0.0250 mol dm^–3 KMnO₄ solution.

What was the concentration of the nitrate(III)  solution in grammes of anhydrous salt per dm³?

Question 13: 2.68 g of iron(II) ethanedioate, FeC₂O₄, were made up to 500 cm³ of acidified aqueous solution. 25.0 cm³ of this solution reacted completely with 28.0 cm³ of 0.0200 mol dm^–3 potassium manganate(VII) solution.

Calculate the mole ratio of KMnO₄ to FeC₂O₄ taking part in this reaction. Give the full redox ionic equation for the reaction.

Question 14: Given the half–cell reaction IO₃^–_(aq) + 6H⁺_(aq) + 5e^– ==> ¹/₂I_2(aq) + 3H₂O_(l) (see also Q2)

(a) Deduce the redox equation for iodate(V) ions oxidising iodide ions.

(b) What volume of 0.0120 mol dm^–3 iodate(V) solution reacts with 20.0 cm³ of 0.100 mol dm^–3 iodide solution?

(c) 25.0 cm³ of the potassium iodate(V) solution were added to about 15 cm³ of a 15% solution of potassium iodide (ensures excess iodide ion). On acidification, the liberated iodine needed 24.1 cm³ of 0.0500 mol dm^–3 sodium thiosulfate solution to titrate it.

(i) Calculate the concentration of potassium iodate(V) in g dm^–3

(ii) What indicator is used for this titration and what is the colour change at the end–point?

Question 15: 25.0 cm³ of an iodine solution was titrated with 0.100 mol dm^–3 sodium thiosulfate solution and the iodine reacted with 17.6 cm³ of the thiosulfate solution.

(a) give the reaction equation.

(b) what indicator is used? and describe the end–point in the titration.

(c) calculate the concentration of the iodine solution in mol dm^–3 and g dm^–3.

Question 16: This question involves titrating ethanedioic acid (oxalic acid), H₂C₂O₄ or (COOH)₂ (i) with standard sodium hydroxide solution and then with potassium manganate(VII) solution (potassium permanganate, KMnO₄).

The titration data is as follows:

10 cm³ of a H₂C₂O₄ solution required 8.50 cm³ of a 0.20 mol dm^-3 (0.20M) solution of sodium hydroxide for complete neutralisation using phenolphthalein indicator (first permanent pink end-point).

10 cm³ of the same H₂C₂O₄ solution required 8.20 cm³ of a KMnO₄ solution for complete oxidation to carbon dioxide and water in the presence of dilute sulfuric acid to further acidify the ethanedioic acid solution (first permanent pink end-point).

(a) Write an equation for the neutralisation reaction of ethanedioic acid with sodium hydroxide.

(b) Calculate the moles of H₂C₂O₄ in the solution and the molarity of the ethanedioic acid solution.

(c) Given the following half-reactions:

(i)   MnO₄^–_(aq)  +  8H⁺_(aq)  +  5e^–  ===>  Mn²⁺_(aq)  +  4H₂O_(l)

(ii)   H₂C₂O_4(aq)  – 2e^–   ===> 2CO_2(g)   +   2H⁺_(aq)

Deduce the full redox titration equation for the oxidation of ethanedioic acid by potassium manganate(VII).

(d) From the equation in (c) and the titration data, deduce the molarity of the potassium manganate(VII) solution.