Mark Riley 3107631608 Investigating Redox Reactions 1

Investigating Redox Reactions

Mark Riley

Introduction Redox reactions are oxidation-reduction reactions which are complementary chemical reactions characterised by the loss or gain, respectively, of one or more electrons by a substance.

Task Practical Report

Aim To determine the redox reaction that has taken place after mixing some common oxidants and reactants.

Procedure This experiment has been split into 7 parts. A separate procedure for each test is given.

Also- See page 6 of the Practical Activities handbook

Equipment Safety glasses, test tube rack, 7 test tubes, reagents in dropping bottle

Mark Riley 3107631608 Investigating Redox Reactions 2

The 𝐸0 given for each half and overall equations have been highlighted as they are only hypothetical values

(eg. if one mole of each was used, at 25oc, oxidation and reduction separated etc in a voltaic battery)

1. Reaction of acidified hydrogen peroxide with iron sulphate

5 drops of 𝐻2𝑂2, 5 drops of 𝐻2𝑆𝑂4 and 10 drops of 𝐹𝑒𝑆𝑂4 were added to a test tube. 1 drop of

KSCN was added and as a result the solution turned blood red indicating the presence of 𝐹𝑒3+.

𝐻2𝑂2 has a higher oxidising strength (oxidant) and 𝐹𝑒2+ has a higher reducing strength

(reductant) ∴ the 𝐹𝑒2+ was oxidized and 𝐻2𝑂2 was reduced according to the following equation.

𝐹𝑒2+ → 𝐹𝑒3+ + 𝑒 × 2 − 0.77V

𝐻2𝑂2 + 2𝐻+ + 2𝑒 → 2𝐻2𝑂 + 1.78V

𝐻2𝑂2 + 2𝐻+ + 2𝐹𝑒2+ → 2𝐻2𝑂 + 2𝐹𝑒3+ ∆V = 1.01V

2. Reaction of acidified hydrogen peroxide with potassium iodide

5 drops of 𝐻2𝑂2, 5 drops of 𝐻2𝑆𝑂4 and 10 drops of 𝐾𝐼(white crystalline solid, colourless in

solution) were added to a test tube. 3 drops of starch were added turning the solution blue-black

indicating the presence of iodine 𝐼2.

𝐻2𝑂2 has a higher oxidising strength (oxidant) and 𝐼− has a higher reducing strength (reductant) ∴

the 𝐼− was oxidized and 𝐻2𝑂2 was reduced according to the following equation.

2𝐼− → 𝐼2 + 2𝑒 − 0.54V

𝐻2𝑂2 + 2𝐻+ + 2𝑒 → 2𝐻2𝑂 + 1.78V

𝐻2𝑂2 + 2𝐻+ + 2𝐼− → 2𝐻2𝑂 + 𝐼2 ∆V = 1.24V

Mark Riley 3107631608 Investigating Redox Reactions 3

3. Reaction of acidified potassium permanganate with iron(II)sulphate

3 drops of 𝐾𝑀𝑛𝑂4, 6 drops of 𝐻2𝑆𝑂4 and 20 drops of 𝐹𝑒𝑆𝑂4 were added to a test tube. The

solution was initially purple but then turned colourless indicating that the 𝑀𝑛𝑂4− had been

reduced to 𝑀𝑛2+. 1 drop of KSCN was added and as a result the solution turned blood red

indicating the presence of 𝐹𝑒3+.

𝑀𝑛𝑂4− has a higher oxidising strength (oxidant) and 𝐹𝑒2+ has a higher reducing strength

(reductant) ∴ the 𝐹𝑒2+ is oxidized and 𝑀𝑛𝑂4− is reduced according to the following equation.

𝐹𝑒2+ → 𝐹𝑒3+ + 𝑒 × 5 − 0.77V

𝑀𝑛𝑂4− + 8𝐻+ + 5𝑒 → 𝑀𝑛2+ + 4𝐻2𝑂 + 1.51V

𝑀𝑛𝑂4− + 8𝐻+ + 5𝐹𝑒2+ → 𝑀𝑛2+ + 4𝐻2𝑂 + 5𝐹𝑒3+ ∆V = 0.74

4. Reaction of acidified potassium permanganate with potassium iodide

3 drops of 𝐾𝑀𝑛𝑂4, 6 drops of 𝐻2𝑆𝑂4 and 20 drops of 𝐾𝐼(white crystalline solid, colourless in

solution) were added to a test tube. The solution was initially purple but then turned colourless

indicating that the 𝑀𝑛𝑂4− had been reduced to 𝑀𝑛2+. 3 drops of starch were added turning the

solution blue-black indicating the presence of iodine 𝐼2. A precipitate was also present.

𝑀𝑛𝑂4− has a higher oxidising strength (oxidant) and 2𝐼− has a higher reducing strength (reductant)

∴ the 2𝐼− is oxidized and 𝑀𝑛𝑂4− is reduced according to the following equation.

2𝐼− → 𝐼2 + 2𝑒 × 5 − 0.54V

𝑀𝑛𝑂4− + 8𝐻+ + 5𝑒 → 𝑀𝑛2+ + 4𝐻2𝑂 × 2 + 1.51V

2𝑀𝑛𝑂4− + 2 × 8𝐻+ + 5 × 2𝐼− → 2𝑀𝑛2+ + 2 × 4𝐻2𝑂 + 5𝐼2

2𝑀𝑛𝑂4− + 16𝐻+ + 10𝐼− → 2𝑀𝑛2+ + 8𝐻2𝑂 + 5𝐼2 ∆V = 0.97

Mark Riley 3107631608 Investigating Redox Reactions 4

5. Reaction of acidified potassium dichromate with potassium iodide

2 drops of 𝐾2𝐶𝑟2𝑂7, 6 drops of 𝐻2𝑆𝑂4 and 15 drops of 𝐾𝐼 (white crystalline solid, colourless in

solution) were added to a test tube. The orange solution turned blue/green indicating that the

𝐶𝑟2𝑂72− had been reduced to 𝐶𝑟3+. 3 drops of starch were added turning the solution blue-black

indicating the presence of iodine 𝐼2.

𝐶𝑟2𝑂72− has a higher oxidising strength (oxidant) and 2𝐼− has a higher reducing strength

(reductant) ∴ the 2𝐼− is oxidized and 𝐶𝑟2𝑂72− is reduced according to the following equation.

2𝐼− → 𝐼2 + 2𝑒 × 3 − 0.54V

𝐶𝑟2𝑂72− + 14𝐻+ + 6𝑒 → 2𝐶𝑟3+ + 7𝐻2𝑂 + 1.23V

𝐶𝑟2𝑂72− + 14𝐻+ + 6𝐼− → 2𝐶𝑟3+ + 7𝐻2𝑂 + 3𝐼2 ∆V = 0.69V

6. Reaction of iron(III)chloride with acidified hydrogen peroxide

5 drops of 𝐹𝑒𝐶𝑙3 , 5 drops of 𝐻2𝑆𝑂4 and 5 drops of 𝐻2𝑂2 were added to a test tube. 1 drop of

KSCN was added, the colour of the solution was unchanged indicating that no 𝐹𝑒3+ was present.

The solution fizzled indicating a the release of oxygen gas 𝑂2.

𝐹𝑒3+ has a higher oxidising strength (oxidant) and 𝐻2𝑂2 has a higher reducing strength

(reductant) ∴ the 𝐻2𝑂2 is oxidized and 𝐹𝑒3+ is reduced according to the following equation.

𝐹𝑒3+ + 𝑒 → 𝐹𝑒2+ × 2 + 0.77V

𝐻2𝑂2 → 𝑂2 + 2𝐻+ + 2𝑒 − 0.70V

𝐻2𝑂2 + 2𝐹𝑒3+ → 𝑂2 + 2𝐻+ + 2𝐹𝑒2+ ∆V = 0.07V

Mark Riley 3107631608 Investigating Redox Reactions 5

7. Reaction of acidified potassium permanganate with hydrogen peroxide

3 drops of 𝐾𝑀𝑛𝑂4, 6 drops of H2𝑆𝑂4 and 10 drops of 𝐻2𝑂2 were added to a test tube. The

solution fizzled indicating the release of oxygen gas 𝑂2.

𝑀𝑛𝑂4− has a higher oxidising strength (oxidant) and 𝐻2𝑂2 has a higher reducing strength

(reductant) ∴ the 𝐻2𝑂2 is oxidized and 𝑀𝑛𝑂4−is reduced according to the following equation.

𝐻2𝑂2 → 𝑂2 + 2𝐻+ + 2𝑒 × 5 − 0.70V

𝑀𝑛𝑂4− + 8𝐻+ + 5𝑒 → 𝑀𝑛2+ + 4𝐻2𝑂 × 2 + 1.51V

2𝑀𝑛𝑂4− + 2 × 8𝐻+ + 5𝐻2𝑂2 → 2𝑀𝑛2+ + 2 × 4𝐻2𝑂 + 5𝑂2 + 5 × 2𝐻+

2𝑀𝑛𝑂4− + 16𝐻+ + 5𝐻2𝑂2 → 2𝑀𝑛2+ + 8𝐻2𝑂 + 5𝑂2 + 10𝐻+

2𝑀𝑛𝑂4− + 6𝐻+ + 5𝐻2𝑂2 → 2𝑀𝑛2+ + 8𝐻2𝑂 + 5𝑂2 ∆V = 0.81

Conclusion

Redox reactions were balanced in the form of chemical equations by arranging the quantities of

the substances involved so that the number of electrons lost by one substance is equaled by the

number gained by another substance. In redox reactions, the substance losing electrons

(undergoing oxidation) is a good electron donor, or reductant because lost electrons are given to

and reduce the other substance. The other substance that gained electrons (undergoing

reduction) is an electron acceptor, or oxidant. Hydrogen peroxide was capable of acting as a

reductant as well as an oxidant.