1. Electrochemical potentials (redox potentials)
Writing cell notations
Calculating emf values

2. Which reaction happens most readily?
Magnesium is more reactive & therefore happens more readily than copper.
Electrochemistry is a way to quantify this.

3. Need to consider everything in terms of equilibria
Build-up of electrons on the magnesium
Surrounded in the solution by a layer of positive ions.
A dynamic equilibrium will be established
The rate at which ions are leaving the surface equals the rate at which they are joining it again
Constant –ve charge on the magnesium
Constant number of magnesium ions in the solution around it.

4. How does copper differ from magnesium?
Copper is less reactive and so forms its ions less readily.
Any ions which do break away are more likely to reclaim their electrons and stick back on to the metal again.
You will still reach an equilibrium position, but there will be less –ve charge on the metal, and fewer ions in solution.

5. Comparing the two positions of equilibrium
By convention, all these equilibria are written with the electrons on the left-hand side of the equation.
Do NOT deviate from this
Electrode potentials is an attempt to attach some numbers to these differing positions of equilibrium
The position of the magnesium equilibrium . . .
 ...lies further to the left than that of the copper equilibrium.

6. Zn2+(aq) e– ⇌ Zn(s)
The difference in the negativeness of the metal (-ve potential) & the positiveness of the solution (+ve potential) or potential difference could be recorded by voltage.
However it cannot be measured directly.

7. Measuring standard electric potentials
If it can’t be measured directly....
.....how could it be done?

8. Using a reference analogy
Wanted to find out how tall a particular person was using an optical device.
Unfortunately, you can't see their feet because they are standing in long grass.
You can however measure their height relative to the convenient post.
You can usefully rank them in order and do some simple sums to work out exactly how much taller one is than another.
For example, C is 5 cm taller than A.

9. Reference used: S tandard H ydrogen E lectrode
100kPa
Temp = 298K
Equilibrium is set up between hydrogen molecules and hydrogen ions in solution.
2H+(aq) e– ⇌ H2(g) Eᶱ = 0
Assigned a potential of 0V
Primary standard
Potential to which all others compared
1 moldm-3 H+(aq)

10. Standard Conditions – Need to know!
Concentration
1.0 mol dm-3 (if ions involved in ½ equation)
Temperature
298 K
Pressure
100 kPa (if gases involved in ½ equation)
Current
Zero (use high resistance voltmeter)

11. SHE is always left hand electrode
Standard hydrogen electrode is attached to the electrode system you are investigating 
Salt bridge
Completes the circuit allowing ions to move from one cell to another
Often KNO3 or KCl
Allows soln. to remain separate
Must not react with either soln.
SHE is always left hand electrode
The whole of this set-up is described as a cell.
Each of the two beakers and their contents are described as half cells.

12. Voltmeter ideally wants to have an infinitely high resistance.
This is to avoid any flow of current through the circuit.
Low resistance in the circuit: electrons would flow from where there are a lot of them to where there are less.
If any current flows, the voltage measured drops.
For proper comparisons, it is important to measure the maximum possible voltage in any situation.
This is called the electromotive force or emf.
The emf of a cell measured under standard conditions is given the symbol E°cell.

13. Emf (measured voltage) Eᶱ cell = Eᶱright - Eᶱleft
Given that Eᶱleft is 0V
Eᶱcell = Eᶱright

14. RH electrode better at gaining e-
Eᶱcell +ve
RH electrode better at gaining e- +

15. RH electrode better at losing e-
Eᶱcell -ve
RH electrode better at losing e-

16. Pt(s) [ H2(g)] | H+(aq) || Cu2+(aq) | Cu(s)
Shorthand way of drawing a cell follow strict cell conventions
Pt(s) [ H2(g)] | H+(aq) || Cu2+(aq) | Cu(s)
reduced oxidised oxidised reduced
I = phase boundary
II = salt bridge
[ ] = gas flowing over Pt electrode
, = if you have >1 aqueous species in a ½ cell eg Fe2+, Fe3+

17. ROOR Ni(s) | Ni2+(aq) || Sn4+(aq), Sn2+(aq) | Pt(s)
K(s) | K+(aq) || Mg2+(aq) | Mg(s)

18. Cell notation
Task 6

19. Standard electrode potentials compare the position of the metal / metal ion equilibrium with the equilibrium involving hydrogen, under standard conditions

20. So......what do these values tell me?
The ones whose positions of equilibrium lie furthest to the left have the most negative E° values.
They form ions more readily - and leave more electrons behind on the metal, making it more negative.
Those which don't lose electrons as readily have positions of equilibrium further to the right.
Their E° values get progressively more positive.

21. The electrochemical series
We can arrange redox equilibria in order of their E° values.
Called the electrochemical series.
Usually the most negative E° values are placed at the top of the electrochemical series, and the most positive at the bottom.
However exam papers often present it ‘upside down’
So.....you need to understand it not just learn it!

22. Metals are more easily oxidised,
What do the E° values tell you about the oxidising or reducing ability in the series of: a) Ions b) Metals
Metals are more easily oxidised,
so are better reducing agents
Ions are more easily reduced,
so are better oxidising agents

23. Connecting two ½ cells together
Most +ve set up as RH electrode
Eᶱ cell = – (-0.76)
Eᶱ cell = V
Eᶱ cell = Eᶱright - Eᶱleft

24. Key Points Key equation, for any cell Eᶱ = Eᶱright - Eᶱleft
In the SHE – the standard electrode (i.e. H) is always the left hand electrode
In cells the more +ve potential is set up as the right hand electrode – way to remember....

25. The abc method of remembering
Metal rod
Left
Right Eᶱ
most Negative Eᶱ
most Positive Eᶱ
Electrode
e- flow Away from rod
Anode
Negative (-ve) electrode
Higher up ECS
e- flow Towards rod
Cathode
Positive (+ve) electrode
Lower down ECS
Redox
Oxidation
Reduction

26. Electrode potentials Task 7
A positive Ecell value is needed for the reaction to go in the forward direction shown by the cell equation.

27. Zn ⇌ Zn2+ + 2 e- oxidation Cu2+ + 2 e- ⇌ Cu reduction
- electrode
anode
oxidation
+ electrode
cathode
reduction
electron flow
At this electrode the metal loses electrons and so is oxidised to metal ions.
These electrons make the electrode negative.
At this electrode the metal ions gain electrons and so is reduced to metal atoms.
As electrons are used up, this makes the electrode positive. Zn Cu
Zn ⇌ Zn e-
oxidation
Cu e- ⇌ Cu
reduction