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Quantitative redox reaction analysis

GCE Advanced A Level REDOX Volumetric Analysis Titration Revision QUESTIONS

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Quantitative volumetric analysis – exam practice redox titration questions based potassium manganate(VII)–iron(II)/ethanedioate–ethanedioic acid (oxalate, oxalic acid)/hydrogen peroxide/sodium nitrite titrations, sodium thiosulfate/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–ethanedioic acid (oxalate, oxalic acid)/hydrogen peroxide/sodium nitrite titrations, Q2/14/17 on sodium thiosulfate–iodine titration, Q3/5 on potassium dichromate(VI)–iron(II) titration,  further Q's will be added – suggestions?

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 + electrons 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.

The relative atomic mass of iron = 55.9

Q1(e) was a late addition and the worked out answers are better presented than maybe others on this page?

(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.

(i) Calculate the molarity of the iron(II) ion.

hint: first work out mol of manganate(VII), then from equation mol of iron(II), then molarity

(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 sulfuric acid and the solution made up to 250 cm³ in a standard graduated flask.

A 25.0 cm³ aliquot of this solution was pipetted into a conical flask and needed 25.45 cm³ of  O.0200 mol dm^–3 KMnO₄ for complete oxidation.

hints as above, then mass of iron, but need to take into account the dilution.

(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.

8.25g of an iron(II) salt was dissolved in 250 cm³ of pure water. 25.0 cm³ aliquots were pipetted from this stock solution and titrated with 0.0200 mol dm^–3 potassium manganate(VII) solution.

The titration values obtained were 23.95 cm³, 23.80 cm³ and 23.85 cm³.

(i) What titration value should be used in the calculation and why?

(ii) Calculate the moles of manganate(VII) used in the titration.

(iii) calculate the moles of iron(II) ion titrated

(iv) Calculate the mass of iron(II) titrated

(v) Calculate the total mass of iron in the original sample of the iron(II) salt.

(vi) calculate the % iron in the salt.

ANSWERS and WORKING

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.

hints: work out moles of thiosulfate, then from equation moles of iodine, then mass iodine

(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?

ANSWERS and WORKING

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.

hints: calculate moles dichromate(VI), from equation moles of iron(II), mass of iron and take into account the dilution

(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)

ANSWERS and WORKING

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 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 ?

hints: equation from Q1, work out mol of Fe(II), from equation mol of manganate(VI), molarity via volume.

ANSWERS and WORKING

Question 5: 10.0 g of iron(II) ammonium sulfate 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)₂SO₄.xH₂O

hints: equation from Q3,  work out mol dichromate(VI), then from equation mol Fe(II), allow for dilution, calculate molar mass of salt, molar mass - water, hence total mass of water in terms of x x 18.

ANSWERS and WORKING

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 hydrated 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.

hints: calculate molar mass then mol of the salt, allow for dilution, then mol manganate(VII) via equation, then molarity via volume.

ANSWERS and WORKING

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

0.150 g of the salt 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.

hints: calculate molar mass then mol of the salt, allow for dilution, then mol manganate(VII) via equation from Q6, then molarity via volume.

ANSWERS and WORKING

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

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

(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^–3 KMnO₄.

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

hints: calculate mol manganate(VII) in titration, from equation calculate mole peroxide, scale up from dilution for total moles in the original solution and its molarity.

ANSWERS and WORKING

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.

hints: See Q1, particularly Q1(e)

ANSWERS and WORKING

Question 10: Calculate the concentration in mol dm^–3 and g dm^–3, the salt 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.

hints: calculate mol manganate(VII), calculate moles of salt from equation in Q6, hence molarity via volume, then mass/volume in appropriate units.

ANSWERS and WORKING

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.

hints: See Q5, but manganate(VII) instead of dichromate(VI) and different ratio.

ANSWERS and WORKING

Question 12:  Given the half–reaction equation:

(i)  NO₃^–_(aq) + 2H⁺_(aq) + 2e^– ==> NO₂^–_(aq) + H₂O_(l)

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

hints: calculate mol of manganate(VII), from equation calculate mol of nitrate(V), hence molarity via volume

ANSWERS and WORKING

Question 13: 2.68 g of the salt 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 - tricky to balance - triple checks of - oxidation numbers up/down, ionic charge (= both sides of equation), and atom count.

hints: calculate molar mass of salt and number of moles, calculate mol of manganate(VII) in titration.

ANSWERS and WORKING

Question 14: Given the half–cell reactions:

(i) I_2(aq)  +  2e^-  ===>  2I^-_(aq)

(ii)  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 to iodine.

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

hints: calculate mol of iodide reacted, from equation calculate mol of iodate(V), then calculate volume

(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

Need 2S₂O₃^2–_(aq)  +  I_2(aq)  ===>  S₄O₆^2–_(aq) + 2I^–_(aq)

hints: calculate mol of thiosulfate, then from equations mol of iodine, then moles of iodate(V), hence volume of iodate(V)

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

ANSWERS and WORKING

Question 15: Calculate the molarities of iron(II) and iron(III) ions in a mixed solution from the following data.

MnO₄^–_(aq) + 8H⁺_(aq)  + 5Fe²⁺_(aq) ===> Mn²⁺_(aq) + 5Fe³⁺_(aq) + 4H₂O_(l)

(i) 25.0 cm³ of the original mixture was acidified with dilute sulfuric acid and required 22.5 cm³ of 0.0200 mol dm^–3 KMnO₄ for complete oxidation.

hints: calculate moles of Fe(II) then molarity from given volume, Fe(III) NOT titrated

(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.

hints: calculate moles of Fe(II) + Fe(III), then total molarity from given volume, all Fe titrated, a little subtraction in the end.

ANSWERS and WORKING

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 rusted 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.

Given: MnO₄^–_(aq) + 8H⁺_(aq)  + 5Fe²⁺_(aq) ===> Mn²⁺_(aq) + 5Fe³⁺_(aq) + 4H₂O_(l)

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 a 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.

ANSWERS and WORKING

Question 17: 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.

ANSWERS and WORKING

Question 18: 1.01g of an impure sample of potassium dichromate(VI), K₂Cr₂O₇, was dissolved in dil. sulfuric 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 thiosulfate 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 thiosulfate ion and iodine.

(c) Calculate the moles of sodium thiosulfate 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.

ANSWERS and WORKING

Question 19: 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 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.

ANSWERS and WORKING

Question 20 Lawn sand containing the salt iron(II) sulfate is used to treat moss.

2.50 g of the lawn sand was mixed with dilute sulfuric acid to extract the iron(II) salt.

In a volumetric titration, it required 24.50 cm³ of a 0.0200 mol dm^-3 solution of potassium manganate(VII) to fully oxidise the iron salt.

(a) Write out the equation for the redox reaction involved in the titration.

(b) Calculate the % by mass of Fe²⁺ ions in the lawn sand. Atomic mass of iron = 55.8

Hint: First calculate mol of KMnO₄ used in titration, then from redox equation calculate the mol of Fe²⁺ titrated ...

Question 21 Analysing a medicinal iron tablet containing an iron(II) salt.

(For people showing signs of iron deficiency e.g. low red blood cell count, suffering from anaemia).

A 1.02 g iron tablet was dissolved in dilute sulfuric acid and titrated with a standard solution of potassium dichromate(VI).

With a suitable redox indicator, it took 23.80 cm³ of a 0.02 mol dm^-3 solution of the dichromate reagent to completely oxidise all the Fe²⁺ ions in the tablet.

(a) Give the full redox equation of the titration.

(b) Calculate the % by mass of Fe²⁺ ions in the tablet. Atomic mass of iron = 55.8

Hint: First calculate mol of K₂Cr₂O₇ used in titration, then from redox equation calculate the mol of Fe²⁺ titrated ...

Question 22 ?

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 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!

practice redox titration questions, how do you do redox titrations? how do you do redox titration calculations? what apparatus do you need to do redox titrations? describe redox titration methods and procedures, what is the definition of a redox titration? how to do redox titration experiments, describe examples of redox titrations, (potassium permanganate) potassium manganate(VII) - iron iron(II) sulfate titration calculation method, sodium thiosulfate Na2S2O3 - I2 iodine titration calculation method, potassium manganate(VII) - iron(II) ammonium sulfate (ferrous ammonium sulfate) standardising titration calculation method, potassium manganate(VII) - ethanedioic acid (oxalic acid) titration calculation method, KMnO4 - potassium tetraoxalate dihydrate titration calculation method, hydrogen peroxide H2O2 KMnO4 potassium manganate (VII) titration calculation method, potassium manganate(VII) KMnO4 - sodium ethanedioate Na2C2O4 titration calculation method, determining water of crystallisation in FeSO4 hydrated iron(II) sulfate, KMnO4 potassium manganate(VII) - nitrite nitrate(III) potassium nitrite titration calculation method, potassium manganate (VII) - iron(II) ethanedioate FeC2O4 titration calculation method, KIO3 potassium iodate iodate(V) - iodine - sodium thiosulfate titration calculation method, analysing iron content of rust or iron ore iron(III) oxide Fe2O3 with potassium manganate(VII) titration calculation method, analysing potassium dichromate dichromate(VI) by oxidising iodide to iodine and titrating iodine with Na2S2O3 sodium thiosulfate keywords–phrases etc.: MnO4–(aq) + 8H+(aq) + 5Fe2+(aq) ==> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) * 2S2O32–(aq) + I2(aq) ==>  S4O62–(aq) + 2I–(aq) * Cr2O72–(aq) + 14H+(aq) + 6Fe2+(aq) ==> 2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l) * 2MnO4–(aq) + 16H+(aq) + 5C2O42–(aq) ==> 2Mn2+(aq) + 8H2O(l) + 10CO2(g) or  2MnO4–(aq) + 6H+(aq) + 5H2C2O4(aq) ==> 2Mn2+(aq) + 8H2O(l) + 10CO2(g) * 2MnO4–(aq) + 6H+(aq) + 5H2O2(aq) ==> 2Mn2+(aq) + 8H2O(l) + 5O2(g) * 2MnO4–(aq) + 6H+(aq) + 5NO2–(aq) ==> Mn2+(aq) + 5NO3–(aq) + 3H2O(l) * IO3–(aq) + 5I–(aq) + 6H+(aq) ==> 3I2(aq) + 3H2O(l) * Cr2O72–(aq) + 14H+(aq) + 6I–(aq) ==> 2Cr3+(aq) + 3I2(aq) + 7H2O(l) * MnO4– + 8H+ + 5Fe2+ ==> Mn2+ + 5Fe3+ + 4H2O * 2S2O32– + I2 ==>  S4O62– + 2I– * Cr2O72– + 14H+ + 6Fe2+ ==> 2Cr3+ + 6Fe3+ + 7H2O * 2MnO4– + 16H+ + 5C2O42– ==> 2Mn2+ + 8H2O + 10CO2 or  2MnO4– + 6H+ + 5H2C2O4 ==> 2Mn2+ + 8H2O + 10CO2 * 2MnO4– + 6H+ + 5H2O2 ==> 2Mn2+ + 8H2O + 5O2 * 2MnO4– + 6H+ + 5NO2– ==> Mn2+ + 5NO3– + 3H2O * IO3– + 5I– + 6H+ ==> 3I2 + 3H2O * Cr2O72– + 14H+ + 6I– ==> 2Cr3+ + 3I2 + 7H2O advanced level chemistry redox titration questions for AQA AS chemistry, redox titration questions for Edexcel A level AS chemistry, redox titration questions for A level OCR AS chemistry A, redox titration questions for OCR Salters AS chemistry B, redox titration questions for AQA A level chemistry, redox titration questions for A level Edexcel A level chemistry, redox titration questions for OCR A level chemistry A, redox titration questions for A level OCR Salters A level chemistry B redox titration questions for US Honours grade 11 grade 12 redox titration questions for pre-university chemistry courses pre-university A level revision notes for redox titration questions  A level guide notes on redox titration questions for schools colleges academies science course tutors images pictures diagrams for redox titration questions A level chemistry revision notes on redox titration questions for revising module topics notes to help on understanding of redox titration questions university courses in science careers in science jobs in the industry laboratory assistant apprenticeships technical internships USA US grade 11 grade 11 AQA A level chemistry notes on redox titration questions Edexcel A level chemistry notes on redox titration questions for OCR A level chemistry notes WJEC A level chemistry notes on redox titration questions CCEA/CEA A level chemistry notes on redox titration questions for university entrance examinations with advanced level chemistry

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