Redox and Electrode Potentials

Question 1

How do you write a redox half equation?

Answer 1

Reduction in the forward reaction and oxidation in the reverse reaction.

Question 2

What would you call a species that wants to be oxidised and why?

Answer 2

A good reducing agent because it reduces other things and is itself oxidised.
Eg: Li+ + e- <—> Li

Question 3

What would you call a species that wants to be reduced and why?

Answer 3

A good oxidising agent because it oxidises other things and is itself reduced.
Eg: F2 + 2e- <—> 2F-

Question 4

How would you combine a redox half equation?
Al3+ + 3e- <—> Al
Fe3+ + e- <—> Fe2+

Answer 4

Find which wants to be reduced or oxidised.
Write both the oxidation and reduction of each in the forward reaction.
Balance the electrons on each half equation.
Cancel out the electrons and combine the equations.
Check that the charges balance.
Al —> Al3+ + 3e-
3Fe3+ + 3e- —> 3Fe2+
Al + 3Fe3+ —> Al3+ + 2Fe2+

Question 5

What is needed to know the feasibility of half redox half equations at equilibrium?

Answer 5

Which are more readily oxidised or reduced. (READILY OXIDISED=REDUCING AGENT=FORWARD REACTION)

Question 6

Give an example of a redox tritration

Answer 6

Iodine and thiosulfate. They react to make peroxidised sulphate and iodide. The reaction goes colourless once all the thiosulfate has reacted with the iodine (it makes the starch a blue black)

Question 7

What are the steps to balancing a half equation?

Answer 7

Look in gallery.

Question 8

Why are transition metals used to reduce and oxidise other substances in redox tritrations?

Answer 8

Because they have variable oxidation states which means they can accept or donate electrons easily. They are coloured depending on their oxidation state, so this makes them useful in identifying end points without using an indicator.

Question 9

Give two examples of transition elements used in redox tritrations.

Answer 9

Eg acidified KMnO4 is an oxidising agent that dissociates to make Mno4- ions. It is reduced from Mn=7+ oxidation state to Mn2+ ions in the titration. (Fe2+ is oxidised) It goes from deep purple to colourless.
Acidified potassium dichromate, K2Fr2O7 is an oxidising agent that dissociates to make Cr2O72- ions. It is reduced from Cr=+6 oxidation state to Cr3+ ions in titrations.(Zn is oxidised) It goes from orange to green.

Question 10

What are redox tritrations used for?

Answer 10

To find the concentration of a reducing or oxidising agent.

Question 11

Explain the steps of a redox titration of unknown Fe2+ concentration titrated against an oxidising agent like Mn04-

Answer 11

Look at picture on phone.

Question 12

What can iodine/thiosulfate reactions be used to determine?

Answer 12

The amount of chlorate(I), ClO-, in bleach.
The amount of copper(II) ions, Cu2+, in copper(II) compounds.
The copper content of alloys.

Manganate(VII) titrations can be used to determine:
The percentage purity of iron supplements

Question 13

What is the standard electrode potential?

Answer 13

2H+ + 2e- <—> H2 (0.00 V)

Question 14

How do we quantitatively measure electron transfer in redox equations?

Answer 14

By measuring electron transfer, aka voltage (potential difference) between two halves of a reaction. Connect then using a Voltmeter.

Question 15

What makes up an electrode and what makes up a cell?

Answer 15

An electrode is made of one redox half equation, two redox half equations form a half cell. Connecting each electrode with a voltmeter determines how strong of an oxidising and reducing agent each is. This requires standard conditions.

Question 16

What are the standard conditions for electrode potential cells?

Answer 16

Room temperature and pressure, and 1 mol dm-3 standard solution.

Question 17

How do you measure the standard electrode potential of a non-hydrogen electrode.

Answer 17

By connecting it with a hydrogen electrode, which is the standard, with a voltmeter. One half cell is H, one half cell is another reaction. The negative electrode (right side) loses electrons (is oxidised) so is the reducing agent. (This means the electrode on the right will be in the backwards reaction of the redox equation.)

Question 18

Why is a platinum electrode used in non-metal half cells?

Answer 18

To make sure the electrode doesn’t interfere with the reaction as it is inert.

Question 19

What is a salt bridge and what does it do?

Answer 19

Used in cells to complete the circuit by allowing ion transfer. It is usually a piece of filter paper soaked in a salt (potassium nitrate).

Question 20

What electrode potential is the forward and backward reaction?

Answer 20

More positive electrode potential = being reduced = forward reaction
More negative electrode = being oxidised = backwards reaction

Question 21

How does changing equilibrium position affect electrode potential?

Answer 21

It can cause an increase in the side of the equilibrium with oxidation or reduction occurring. If oxidation increases, electrode potential gets more negative as more electrons are lost.
If reduction increases electrode potential gets more positive as more of the original ion/atom is made (less e- lost).

Question 22

How do you calculate the standard electrode potential?

Answer 22

More positive Eꝋ - less positive Eꝋ
Ecellꝋ = E right ꝋ - Eleft ꝋ
Ecellꝋ = E reduction ꝋ - Eoxidation ꝋ

Question 23

How do fuel cells work?

Answer 23

A combustion reaction where the reactants are separated from the fuel but both are in the same electrolyte. No collision occurs but the electrons are transferred via a wire and used to power something connected to the wire. Eg: H fuel cells.