Doc Brown's Advanced Level Chemistry
Quantitative redox reaction analysis
GCE Advanced Level REDOX Volumetric Analysis Titration Revision Questions
Quantitative volumetric analysis – exam practice redox titration questions based potassium manganate(VII)–iron(II)/ethanedioate–ethanedioc acid (oxalate, oxalic acid)/hydrogen peroxide/sodium nitrite titrations, sodium thiosulphate/thiosulfate–iodine titrations and potassium dichromate(VI)–iron(II) titration. Any suggestions for additional types of A level redox titration questions?
Volumetric analysis worksheet of structured questions of REDOX VOLUMETRIC TITRATION CALCULATIONS
Titrations and calculations based on oxidation–reduction techniques–reactions – solved problems
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,
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?
REDOX–ionic EQUATION CHECKS
I've tried to quote the data to the appropriate significant figures and associated 'trailing zeros'.
Note that it is standard convention to show half-cell reactions as reductions, i.e. atom/ion/molecule + electron(s) to give the reduction product. This means you have to judge whether the half-reaction needs to be reversed to derive the full ionic redox equation, and any multiples of it are needed, - so take care, and don't get confused by conventions! they are there help!
Question 1: Given the following two half–reactions: (Q1 can be done as an experimental 'word–fill' version)
Question 1 has many parts covering the titration of iron(II) ions with a standard solution of potassium manganate(VII) and the problems are solved.
(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 0.0200 mol dm–3 KMnO4 reacted with 20.0 cm3 of an iron(II) solution acidified with dilute sulfuric acid.
(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.
(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 to be used in this context.
(e) The analysis of a soluble iron(II) salt to obtain the percentage of iron in it.
Question 2: Given the following two half–reactions
(a) Given (i) S4O62–(aq) + 2e– ==> 2S2O32–(aq)
and (ii) I2(aq) + 2e– ==> 2I–(aq)
construct the full ionic redox equation for the reaction of the thiosulphate ion S2O32– and iodine I2.
(b) what mass of iodine reacts with 23.5 cm3 of 0.0120 mol dm–3 sodium thiosulphate 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 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, 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) Cr2O72–(aq) + 14H+(aq) + 6e– ==> 2Cr3+(aq) + 7H2O(l)
deduce the fully balanced redox equations for the reactions
(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)
Question 4: An approximately 0.02 mol dm–3 potassium manganate(VII) solution was standardized against precisely 0.100 mol dm–3 iron(II) ammonium sulphate 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 sulphate crystals were made up to 250 cm3 of acidified aqueous solution. 25.0 cm3 of this solution required 21.25 cm3 of 0.0200 mol dm–3 potassium dichromate(VI) for oxidation.
Calculate x in the formula FeSO4.(NH4)2S04.xH2O
Question 6: Given the half–reaction C2O42–(aq) – 2e– ==> 2CO2(g)
(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 250.0 cm3 of aqueous solution and 25.00 cm3 of this solution needed 24.55 cm3 of a potassium manganate(VII) solution for oxidation.
Question 7: A standardization of potassium manganate(VII) solution yielded the following data:
0.150 g of potassium tetroxalate (tetraoxalate?), 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 O2(g) +2H+(aq) + 2e– ==> H2O2(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.
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.
If 20.0 cm3 of this solution required 26.5 cm3 of 0.0100 mol dm–3 KMnO4 for oxidation.
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 24.50 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) sulphate crystals were made up to 500 cm3 acidified with sulphuric 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 of 25.0 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– ==> 1/2I2(aq) + 3H2O(l) (see also Q2)
(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 0.0500 mol dm–3 sodium thiosulphate solution to titrate it.
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 0.0200 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 0.0200 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 EDTA4–(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 17: 25.0 cm3 of an iodine solution was titrated with 0.100 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 25.0 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 0.100 mol dm–3 sodium thiosulphate and the starch turned colourless after 20.0 cm3 was added.
(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 25.0 cm3 aliquot titrated.
(f) Calculate the total mass of potassium dichromate(VI) in the original sample and hence its % purity.
I DO MY BEST TO CHECK MY CALCULATIONS, as you yourself should do, BUT I AM HUMAN! AND IF YOU THINK THERE IS A 'TYPO' or CALCULATION ERROR PLEASE EMAIL ME ASAP TO SORT IT OUT!
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