1. Advanced Higher Chemistry Unit 2(e)
Electrochemistry
Advanced Higher Chemistry
Unit 2(e)

2. Electrochemistry
The study of electron transfer and its relationship with electric currents.

3. Electrochemical cells
Are systems consisting of electrodes dipped into an electrolyte in which a chemical reaction uses or generates an electric current.
A voltaic cell is an electrochemical cell in which a spontaneous reaction generates an electric current.
An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise non-spontaneous reaction.

4. When a metal is placed in a solution of its ions, some of the metal atoms ionise:
M(s) M+(aq) + e-
M+(aq) + e M(s)
For some metals, e.g. magnesium, the equilibrium lies well to the left
The solution becomes slightly positive with respect to the metal
For other metals, e.g. copper, the equilibrium lies less to the left
The solution becomes less positive with respect to the metal

6. Potential difference
In both cases there is a slightly different charge between the metal and the solution
This is known as a potential difference
The potential difference in the case of the Mg is greater than in that of the Cu
However, it cannot be measured directly

7. Instead, measurements are made in terms of the difference compared to the standard hydrogen electrode.
This has a value of 0.00V.

8. Standard hydrogen electrode
298K

10. The potential difference between the two metals can be measured using a voltmeter.
The voltmeter would show the Mg as being the negative electrode and the H as being the positive electrode.
This is relative.

11. Ideally, the voltmeter would have an infinitely high resistance.
This would prevent it from drawing any current.
If current flows, the voltage drops.
To make viable comparisons between cells, it is important that the voltages referred to are the maximum values.

12. Emf Electromotive force, Eºcell
“The electric potential difference between the electrodes in a cell when no current is drawn”.
The maximum possible voltage in a situation – the driving force.
Value depends on concentration, temperature and the type of cell.

14. Cell conventions Zn(s)Zn2+(aq)Cu2+(aq)Cu(s)
The anode is always written on the left; the cathode on the right.
Terminals are at the extremes of the notation.
Single vertical lines represent phase boundary.
Double vertical line represents salt bridge.

15. Standard electrode (reduction) potentials
The cell emf is composed of contributions from the anode (oxidation reaction) and the cathode (reduction reaction).
These contributions are known as the oxidation potential and reduction potential.
Eºcell = oxidation potential + reduction potential

16. Reduction potential
Is a measure of the ability of a species to act as an oxidising agent.
The oxidation potential of a species equals the negative of the reduction potential.
It is an intensive property – it is independent of the amount of the species present in the reaction.

17. Those with the most positive reduction potentials have the greatest tendency to go left to right as written.
Those species being reduced are acting as oxidising agents.
The best oxidising agents are therefore those at the bottom of the table on the left.
The best reducing agents are at the top of the table, on the right.

18. Oxidising power
oxidation
reducing agent
reduction
oxidising agent
Reducing power
A half-reaction will be able to force any half reaction above it to go in reverse (i.e. as an oxidation)

19. Electrochemical series
Example
Imagine connecting a zinc half-cell to the standard hydrogen electrode.
The voltmeter reading is +0.76V.
The zinc is the anode – it is being oxidised (losing electrons).
So the standard reduction potential (for gain of electrons) is -0.76V.

20. Calculating cell emf Calculate the standard emf of a zinc-copper cell.
Ecell = oxidation potential + reduction potential
ANODE (oxidation): Zn
Reduction potential = -0.76V
Oxidation potential = +0.76V
CATHODE (reduction): Cu
Reduction potential = +0.34V
Ecell = = V

21. A positive emf will always be obtained if the reaction takes place as written i.e. the reaction is spontaneous.
This is related to the standard free energy change:
Gº = -nFE

22. Example calculation
Calculate the standard free energy change for the silver-nickel cell.
Solution
Write the ion-electron equations
Determine the oxidation and the reduction
Derive Eº
Combine to form the redox equation
Derive n
Calculate Gº

23. Fuel cells
Fuels are like batteries in that they use redox reactions to generate electricity.
However, unlike in a battery cell, in a fuel cell the reactants are continually fed into the cell.
The fuel is passed into the anode and oxygen to the cathode.