Redox (II): Standard Electrode Potential Eo(2)

Edexcel A level Chemistry (2017)

Topic 14: Redox (II): Standard Electrode Potential Eo(2)

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

Standard Electrode Potential Eo

The e.m.f. generated when a half-cell is connected to the standard hydrogen electrode is called the standard electrode potential (Eº) of that half-cell. The units of Eº  are volts (v).

It is measured using solutions of 1M, any gases involved are at a pressure of 100kPa (what used to be 1 atmosphere) and measurements are made at 298K or 25℃.

The hydrogen electrode is a reference electrode and that means it is an electrode against which all other electrode potentials are measured.

Reference electrode also means that the hydrogen electrode is given the value 0.00v so that the emf measured when another half cell is connected up to it, is the  Eº of that half cell.

The Standard Hydrogen Electrode

This is what the setup looks like:

Hydrogen gas at 100kPa passes over a platinum electrode of very high surface area (it is platinised with platinum black, a form of platinum porous to hydrogen gas that creates a good electrical contact with the hydrogen ions in the solution).

The hydrogen is bubbled into 1M H⁺(aq) ions at 298K.

The equilibrium at the hydrogen electrode is

H₂(g)   ⇌   2H⁺(aq)   +   2e–

Connecting a copper/copper ion half-cell to the hydrogen electrode produces an emf of +0.34v.

The convention is always to place the standard hydrogen electrode as the left hand electrode and write the other half-cell as if the electrons generated from the hydrogen electrode turn metal ions into metal atoms.

In other words, we assume that in every case hydrogen will be oxidised and the right hand electrode reduced.

If this happens then the emf generated has a positive value (and we’ll come back to this idea later.)

This is exactly what happens in the case of the Cu/Cu²⁺ half-cell.

Since hydrogen is more reactive than copper, hydrogen gas molecules are oxidised and the copper ions are reduced.

H₂(g)   ⇌   2H⁺(aq)   +   2e–   and   Cu²⁺(aq)   +   2e– ⇌  Cu(s)

This is the set up:

But if we want to measure the Standard Electrode Potential of zinc, we must connect the Zn/Zn²⁺ half-cell to the hydrogen half cell as in the diagram below.

When we do this, we obtain an emf of –0.76v. 

The negative sign of the emf is significant.

The negative sign confirms what we already know that zinc is more reactive than hydrogen.

It tells us that hydrogen does not reduce zinc to zinc ions but conversely zinc reduces hydrogen ions to hydrogen.

The electrons in the cell are released from zinc not hydrogen.

The two redox equilibria move in the directions indicated in the diagram above.   

Standard electrode potentials are usually presented in tables like the one below.

You can see that each half cell is written in the table as if it were the right hand electrode of a cell with the hydrogen electrode as the left hand half cell.

Species at the top right of the table are the best reducing agents: 
e.g. lithium Eº  = –3.04v.

Species at the bottom left of the table are the best oxidising agents: 
e.g fluorine Eº  = +2.87v.

In my next post, I’ll discuss the use of these electrode potentials and what the values signify.