Redox (II): Standard Electrode Potential Eo (1)

Edexcel A level Chemistry (2017)

Topic 14: Redox (II): Standard Electrode Potential Eo 

Here are four further learning objectives:

14/3 To know what is meant by the term ‘standard electrode potential’, Eo

14/4 To know that the standard electrode potential, Eo, refers to conditions of:

i) 298 K temperature 


ii) 100 kPa (1 atm) pressure of gases 


iii) 1.00 mol.dm-3 concentration of ions

14/5 To know the features of the standard hydrogen electrode and understand why a reference electrode is necessary

14/6 To understand that different methods are used to measure standard electrode potentials of:

i) metals or non-metals in contact with their ions in aqueous solution 


ii) ions of the same element with different oxidation numbers

Background to Standard Electrode Potential Eo

All metals in an aqueous solution of their salt have a tendency to release electrons and go into solution as the metal ion.

The electrons released would form an electrical current.

The tendency of metals to lose electrons in solution we call the absolute electric potential.

But it cannot be measured.

The best we can do is put one metal up against another and see which of the two loses electrons so that the other metal gains them.

The result is a simple cell with a specific emf.

The Daniell Cell is a classical example of this arrangement.

A Daniell cell has a zinc metal rod in 1M zinc(II)sulphate solution and a copper wire in a 1M copper(II)sulfate solution.

A porous pot separates the two metals and their metal ion solutions.

The emf can be measured if a very high resistance voltmeter connects the two metals.

A high resistance voltmeter hardly draws down any current from the cell in order to work.

There are several possible practical set ups you can use as the images below reveal:

What happens in a Daniell Cell?

Since zinc is more reactive than copper it loses electrons more easily than copper.

Zinc atoms leave the zinc metal and go into the solution as zinc ions (Zn²⁺)

Electrons flow from the zinc electrode onto the copper electrode.

Copper ions (Cu²⁺) absorb these electrons and become copper atoms and these add to the copper electrode.

Oxidation takes place at the zinc electrode since zinc atoms lose electrons.

Zn(s)     ⟶      Zn²⁺(aq)   +    2e–

Reduction takes place at the copper electrode since copper ions in the solution gain electrons to become copper atoms on the electrode.

Cu²⁺(aq)   +    2e–     ⟶      Cu(s)

The difference in electrical potential between the copper/copper ion system and the zinc/zinc ion system is enough to produce a voltage of +1.10v from the cell.

If electrons flow from the zinc electrode to the copper electrode then the current flows in the opposite direction: copper to zinc.

Since current always flows from positive to negative then this make the copper electrode the positive terminal of the cell and the zinc electrode the negative terminal.

The salt bridge in the photo above is there to complete the electrical circuit.

As ions build up in the zinc sulphate solution the solution becomes more positively charged.

Equally as ions are removed from the copper salt solution that solution becomes more negatively charged.

So to maintain the electrical balance between both solutions, ions migrate from the salt bridge into the copper solution and from the zinc salt solution into the salt bridge.

This video here helps explain all this:

In an attempt to standardise this chemistry, a standard electrode has been developed against which all other electrodes potentials are measured.

In my next post, I’ll explain how standard electrode potentials are measured and how they can be used.