1. Redox Processes AHL

2. Electrochemical Cells
All half cell potentials are based on an arbitrary comparison to the potential of a standard hydrogen electrode.
The SHE consists of an inert platinum electrode in contact with 1 mol dm-3 hydrogen ion solution and hydrogen gas at 100kpa and 298K.

3. The Platinum Electrode
Conducts electricity.
Unreactive.
Can have a high surface area to increase rate of reaction.

4. Electrochemical Cells
The potential of a voltaic (galvanic) cell is determined by the difference in potential between the cathode and the anode.
Eϴcell = Eϴrhe - Eϴlhe
rhe = right hand electrode
lhe = left hand electrode

5. Electrochemical Cells
Use the activity series to determine which of the following half cells would contain the anode and which would be the cathode.
Write half equations for each reaction, including the half cell potential (Eϴ).
Determine the overall cell potential.
Ag/Ag+
Zn/Zn2+

6. Electrochemical Cells
If the cell potential is positive, the reaction is spontaneous.
If Gibbs free energy is negative, the reaction is spontaneous.
Gibbs free energy relates to cell potential according to the following reaction:
ΔGϴ = -nFEϴcell
F = Faraday’s constant, C mol-1

7. Important Note
It is necessary to balance electrons when combining half equations.
The coefficient of electrons DOES NOT affect the cell potential.
It DOES, however, affect the Gibbs free energy.

8. Electrochemical Cells
Calculate the value of Eϴcell for the following reaction and hence determine whether it will be spontaneous or not.
Calculate the Gibbs free energy for the reaction.
Sn2+ + 2Fe3+  Sn4+ + 2Fe2+
Given: Sn4+ + 2e-  Sn2+ Eϴ = V

9. Electrochemical Cells
In an aqueous solution, if water is a stronger reducing agent than one species, it will be oxidized at the anode to H2 gas and OH- ions.
If water is a stronger oxidizing agent than the other species, it will be reduced at the cathode to H+ and O2.
Eg. AlF3
Al2+ + 3e-  Al Eϴ = V
H2O + e-  ½H2 + OH- Eϴ = V
½O2 + 2H+ + 2e-  H2O Eϴ = V
½F2 + e-  F Eϴ = V

10. Electrochemical Cells
Determine the products that will be formed during the electrolysis of zinc fluoride.
H2O + e-  ½H2 + OH- Eϴ = V
Zn2+ + 2e-  Zn Eϴ = V
H2O  ½O2 + 2H+ + 2e- Eϴ = V
F-  ½F2 + e- Eϴ = V

11. Important Note
The following can all affect the cell potential and/or species reacting:
Temperature
Rate of reaction
Concentration of reactants
For example in a solution of NaCl, water should be oxidized first. However, since the reaction is slow, in a concentrated solution of NaCl the Cl- ions will be oxidized.

12. Self Ionization of Water
Pure water is not pure H2O. Since water itself is a weak acid it undergoes a process called ‘self ionization’:
H2O(l)  H+ + OH-
The concentrations of H+ and OH- are very low (10-7M).

13. Electrochemical Cells
Concentrated sodium chloride.
Oxidation of H2O is slow, causing Cl- to be oxidized due to a phenomenon called ‘overvoltage’.
Dilute sodium chloride.
‘Overvoltage’ doesn’t play a part, therefore the reaction that occurs at the anode is the oxidation of OH-.

14. Electrochemical Cells
We can predict the amount of product formed if we have the following information:
Current (I)
Duration of electrolysis (t)
Charge on the ion (Z)
They relate to each other by the following equation:
Q = It
Where Q is the total charge in Coulombs.

15. Electrochemical Cells
Electrical charge is measured in Coulombs.
A Coulomb is defined as one amp for one second.
1C = 1A x 1s
96500 coulomb of charge (known as Faraday’s constant) is equal to one mole of electrons.
n = Q/F
Whatever, man. I don’t care about nothin’. Ω

16. Electrochemical Cells
Calculate the mass in grams of copper produced at the cathode when a current of 1.50 A is passed through a solution of aqueous copper (II) sulfate for s.
mol
5.78 g

17. Electrochemical Cells
Electrolysis can be used to ‘plate’ objects such as jewelry or metals that need protection from corrosion.
To silver plate an object an electrolytic cell is set up with impure silver as the anode and the object as the cathode.

18. Guidance
Electrolytic processes to be covered in theory should include the electrolysis of aqueous solutions (eg sodium chloride, copper (II) sulphate etc.) and water using both inert platinum or graphite electrodes and copper electrodes. Explanations should refer to E values, nature of the electrode and concentration of the electrolyte.
Delta G = -nFE is given in the data booklet in section 1.
Faraday’s constant = 96500C/mol is given in the data booklet in section 2.
The term ‘cells in series’ should be understood.