Doc Brown's Chemistry  Quantitative redox reaction analysis

REDOX VOLUMETRIC TITRATION CALCULATIONS

Titrations and calculations based on oxidation-reduction techniques-reactions - solved problems

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses)

Relative atomic masses that may be needed, in alphabetical order of symbol ...

C=12.0,  Cr = 52.0, Fe=55.9,  H=1.0,  I=126.9,  K=39.1,  Mn=54.9,  N=14.0,  Na=23.0,  O=16.0,  S=32.1,

Equations needed * FULL ANSWERS and WORKING * REDOX REACTION THEORY * Qualitative Analysis

and also acid-base and other non-redox titrations Questions * EMAIL query?comment

NOTE: Half reactions are usually quoted as the half-cell reduction equation. Reuse half-cell or full equations in later questions from earlier questions. It is assumed you will work through them in numerical question order. If you cannot work out the redox equations, you can just download the equations so that you can at least practice the 'pure' volumetric calculation aspects of the questions. 

Questions 1/4/6/7/8/9/10/11/12/13/15/16 based potassium manganate(VII)-iron(II)/ethanedioate-ethanedioc acid (oxalate, oxalic acid)/hydrogen peroxide/sodium nitrite titrations, Q2/14/17 on sodium thiosulphate-iodine titration, Q3/5 on potassium dichromate(VI)-iron(II) titration,  further Q's will be added - suggestions?

Question 1: Given the following two half-reactions: (Q1 can be done as an experimental 'word-fill' version)

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

and (ii) Fe³⁺_(aq) +  e^- ==> Fe²⁺_(aq)

(a) Construct the fully balanced redox ionic equation for the manganate(VII) ion oxidising the iron(II) ion

(b) 24.3 cm³ of 0.02 mol dm^-3 KMnO₄ reacted with 20.0 cm³ of an iron(II) solution.

(i) Calculate the molarity of the iron(II) ion. (ii) How do recognise the end-point in the titration?

(c) Calculate the percentage of iron in a sample of steel wire if 1.51 g of the wire was dissolved in excess of dilute sulphuric acid and the solution made up to 250 cm³ in a standard graduated flask.  25.0 cm³ of this solution was pipetted into a conical flask and needed 25.45 cm³ of  O.02 mol dm^-3 KMnO₄ for complete oxidation.

(d) Suggest reasons why the presence of dil. sulfuric acid is essential for an accurate titration and why dil. hydrochloric and nitric acids are unsuitable.

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 thiosulphate ion S₂O₃^2-,and iodine.

(b) what mass of iodine reacts with 23.5 cm³ of 0.012 mol dm^-3 sodium thiosulphate solution.

(c) 25cm³ of a solution of iodine in potassium iodide solution required 26.5 cm³ of 0.095 mol dm^-3 sodium thiosulphate 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, Fe₂0₃] were dissolved in concentrated hydrochloric acid and the solution diluted to 250 cm³.

25.00 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 required 26.4 cm³ of 0.02 mol dm^-3 potassium dichromate(VI) for oxidation.

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

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

(c) Why isn't potassium manganate(VII) used for this titration? (as in Q1)

Question 4: An approximately 0.02 mol dm^-3 potassium manganate(VII) solution was standardized against precisely 0.1 mol dm^-3 iron(II) ammonium sulphate solution. 25.00 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 sulphate crystals were made up to 250 cm³ of acidified aqueous solution. 25 cm³ of this solution required 21.25 cm³ of 0.02 mol dm^-3 potassium dichromate(VI) for oxidation.

Calculate x in the formula FeSO₄.(NH₄)₂S0₄.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 ethane-dioic acid).

(b) 1.520 g of ethanedioic acid crystals, H₂C₂O₄.2H₂O, was made up to 250 cm³ of aqueous solution and 25.0 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:

0.15 g of potassium tetroxalate (tetraoxalate?), KHC₂O₄.H₂C₂O₄.2H₂O needed 23.2 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 cm³ of solution of hydrogen peroxide were diluted to 1 dm³ with water. 25.0 cm³ of this solution, when acidified with dilute sulphuric acid, reacted with 20.25 cm³ of 0.02 mol dm^-3 KMnO₄.

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 sulphuric acid. The mixture was filtered and the filtrate and washings made up to 500 cm³ in a standard volumetric flask. 20.0 cm³ of this solution required 26.5  cm³ of 0.01 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 cm³ of which were oxidized in acid solution by 24.5 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) sulphate crystals were made up to 500 cm³ acidified with sulphuric acid.

25.0 cm³ of this solution required 43.85 cm³ of 0.01 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 of 25 cm³ of an acidified 0.025 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 0.02 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 0.012 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 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 0.05 mol dm^-3 sodium thiosulphate 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: Calculate the molarities of iron(II) and iron(III) ions in a mixed solution from the following data.

(i) 25.0 cm³ of the original mixture was acidified with dilute sulphuric acid and required 22.5 cm³ of 0.02 mol dm^-3 KMnO₄ for complete oxidation.

(ii) a further 25.0 cm3 of the original iron(II)/iron(III) mixture was reduced with zinc and acid and it then required 37.6 cm³ of the KMnO₄ for complete oxidation.

Question 16: 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 sulphuric 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 0.020 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 25.0 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 0.10 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 17: 25.0 cm³ of an iodine solution was titrated with 0.1 mol dm^-3 sodium thiosulphate solution and the iodine reacted with 17.6 cm³ of the thiosulphate 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 18: 1.01g of an impure sample of potassium dichromate(VI), K₂Cr₂O₇, was dissolved in dil. sulphuric acid and made up to 250 cm³ in a calibrated volumetric flask. A 25 cm³ aliquot of this solution pipetted into a conical flask and excess potassium iodide solution and starch indicator were added. The liberated iodine was titrated with 0.1 mol dm^-3 sodium thiosulphate and the starch turned colourless after 20.0 cm³ was added.

(a) Using the half-equations from Q3(a)(ii) and Q2(a)(ii), construct the full balanced equation for the reaction between the dichromate(VI) ion and the iodide ion.

(b) Using the half-equations from Q2(a) construct the balanced redox equation for the reaction between the thiosulphate ion and iodine.

(c) Calculate the moles of sodium thiosulphate used in the titration and hence the number of moles of iodine titrated.

(d) Calculate the moles of dichromate(VI) ion that reacted to give the iodine titrated in the titration.

(e) Calculate the formula mass of potassium dichromate(VI) and the mass of it in the 25 cm³ aliquot titrated.

(f) Calculate the total mass of potassium dichromate(VI) in the original sample and hence its % purity.

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