TOPIC 14 - Redox II

Question 1

How is an electrochemical cell set up?

Answer 1

1. Clean the metals of the electrodes using sandpaper to remove impurities in the surface.
2. Wash surface of electrode metals with propanone to remove any grease from the surface of the metal.
3. Place each metal into a solution containing the ion of the same metal.
4. Make the salt bridge from filter paper soaked in saturated KNO3 (or KCl). Each end should be submerged in each solution.
5. Connect the electrodes with wires, crocodile clips and a voltmeter.

Question 2

What is a half cell?

Answer 2

Half of an electrochemical cell. Constructed of a metal cipped in its ions, or a platinum electrode with two aqueous ions (i.e. Fe2+, Fe3+)

Question 3

Why is platinum used for electrodes?

Answer 3

It is inhert and conducts electricity.

Question 4

How can you predict the reactions in electrochemical cells by looking at the electrode potentials?

Answer 4

The half-equation with the most -ve electrode potential will undergo the reverse reaction than the one shown in the electrode potential, while the most +ve will undergo the forward reaction.

Question 5

What is the standard hydrogen electrode?

Answer 5

The reference half cell to which all the standard electrode potentials are measured against.

Question 6

What are the standard conditions for the standard electrode potentials?

Answer 6

Temperature - 298K

Pressure - 100kPa

1M conc. of ions.

Measured against standard hydrogen electrode.

Question 7

How can you calculate the standard cell potential by using the standard electrode potential?

Answer 7

Add up the potentials of the two standard electrode potentials (remember to change the sign if the reaction is going the other way).

Question 8

Why do we use standard conditions to measure electrode potentials?

Answer 8

Most of the reactions are reversible and so different conditions will change the equilibrium of the reactions and therefore will change the potential. Therefore we use standard conditions to be able to compare different electrode potentials.

Question 9

How are cells described using cell notation?

Answer 9

Question 10

How can the cell potential describe if a reaction is feasible?

Answer 10

If Ecell is positive, reaction is feasible.

Question 11

Why might a reaction might not occur even though the cell potential is feasible (+ve)?

Answer 11

There might not be standard conditions, so the equilibrium may change, and so the electrode potentials might change.

There might be non-favourable kinetics: Reaction might be very slow or need a very high actication energy.

Question 12

How does charging electrochemical cells work?

Answer 12

A current is supplied, which forces electrons to flow in the opposite way, so the reverse reaction occurs, regenerating the reactants for the discharge reaction.

Question 13

How does an alkaline hydrogen-oxygen fuel cell work?

Answer 13

Question 14

How do methanol (or ethanol) fuel cells work?

Answer 14

Question 15

How do acidic hydrogen-oxygen fuel cells work?

Answer 15

Hydrogen arrives at anode and is oxidised: H2 > 2H+ + 2e-

Electrons flow through circuit, H+ travel through the acid.

Electrons + H+ travel to the cathode and react together with oxygen, givng water.

1/2O2 + 2H+ + 2e- > H2O

Question 16

Answer 16

Question 17

How would you find the amount of iodate ions by using an iodine-sodium thiosulfate titration?

Answer 17

1. Add excess acidified KI solution to the IO3- solution:
2. IO3 (-) + 5I- + 6H+ > 3H2O + 3I2
3. Add solution to flask and titrate sodium thiosulfate (Na2S2O3) and once you see a pale yellow colour, add 2cm^3 of starch (which turns deep blue in presence of Iodine) and tittrate until the solution turns colourless:
4. Using the oles of thiosulfate, calculate the moles of iodine:
5. I2 + 2S2O3 (2-) > S4O6 (2-) + 2I-
6. Use previous equation to work out the moles of IO3 (-) using the moles of I2.
7. Divide by moles if you want the concentration.

Question 18

How would you find the concentration of Fe2+ ions by titrating agains MnO4 (-) ions?

Answer 18

1. Place an MnO4 (-) solution of known concentration in the burette.
2. Have a known volume of Fe 2+ solution in the conical flask.
3. Add excess dilute sulfuric acid to ensure there’s enough H+ ions in order to reduce MnO4 (-)
4. Add MnO4 (-) until purple colour is seen. This is the colour of MnO4 (-) ions and so indicates that all the Fe 2+ ions have reacted.
5. Calculate the number of moles in MnO4 (-) in the titre.
6. Use it to calculate the number of moles of Fe2+ :
7. MnO4 (-) + 8H+ + 5Fe2+ > Mn2+ + 5Fe3+ + 4H2O
8. Divide by volume for conc.

Question 19

How would you use iodide ions to calculate the percentage of copper in an alloy?

Answer 19

1. Dissolve a mass of alloy in concentrated nitric acid and make up to the 250cm^3 with deionised water.
2. Pipette 25cm^3 into a flask. Add sodium carbonate solution o neutralise the acid. Stop adding when a precipitate is seen.
3. Add a few drops of ethanoic acid to remove the precipitate.
4. Add excess KI for the I- ions to react:
5. 2Cu2+ + 4I- > 2CuI (s) + I2 (aq)
6. Titrate the solution with sodium thiosulfate to find the number of moles of I2. (add starch when pale yellow colour is seen, end point is when starch turns colourless).
7. Using the eqt above find the numer of moles of copper ions by using the number of moles of iodine.
8. Find mass and percentage by mass of the copper ions.

Question 20

What should be taken into account during the iodine+sodium thiosulfate titration?

Answer 20

Starch solution should only be added when most of the iodine has reacted.

Starch solution should only be made when it is ready to use.

Keep the solution cool as I2 might eveaporate.

Question 21

Wit the iodine - thiosulfate titration, why might it be difficult to see a pale yellow colour?

Answer 21

CuI precipitate might make it difficult to see.

Question 22

What is the equation linking Gibb’s free energy and Ecell?

Answer 22

Question 23

What is the equation linking Ecell and K (equilibrium constant)?

Answer 23