Oxidising and reducing agents

Table of Contents

1. Oxidising agents and reducing agents
2. Redox reactions
   1. Oxidation numbers
3. Oxidising agents and reducing agents
   1. Oxidants and reductants
   2. Common oxidants and reductants
4. Half-equations
   1. Writing half-equations
   2. Combining half-equations
5. Reduction potentials
   1. Source of reduction potentials
   2. Using standard electrode potentials
   3. Predicting reaction tendencies
6. Redox reactions in action

1. Additional resources:
   1. Glossary
   2. Print content

Source of reduction potentials

Consider an oxidising agent (oxidant), such as chlorine gas. We have seen that as an oxidant it has the ability to 'soak up' electrons. There is a 'force of combination' that causes electrons to join with chlorine atoms.

Reduction potentials are a way of comparing the relative strength of these forces. For example, fluorine is a stronger oxidising agent than chlorine. This means that the force by which fluorine combines with electrons is stronger than the force for chlorine.

So what is this force?

Because redox reactions involve the transfer of electrons, the force is described as the Electro (electrons) Motive (motion) Force, or EMF.

The EMF is measured as a voltage, which is why the reduction potentials are given as voltages.

To understand the measurement of reduction potentials better, you can learn about electrochemical cells by working through the section on Reduction potentials and electrochemical cells. The concept of electrochemical cells is dealt with in more detail in a later topic, but here the idea is introduced in the context of electrode potentials.

Reduction potentials and electrochemical cells

Don't worry if the idea of electrochemical cells is a little confusing at first. When you encounter more examples later in the course the ideas should become clearer.

An electrochemical cell is one in which the reduction and oxidation processes within a redox reaction are occurring in two separate places, and with the two sides of the system connected to allow for the movement of electrons from one 'half-cell' to the other. You will have encountered a somewhat similar set-up when studying electrolysis in Stage 2 Chemistry. The diagram below shows an electrochemical 'half-cell'.

In theory, if the copper ions reacted to produce copper metal, the following reaction would occur.
Cu²⁺ + 2 e^- → Cu(s)

If this were the case, it may be possible to measure the force of the movement of these electrons. However, this will never happen because there is no complete circuit so movement of electrons is not possible.

In effect one half-cell will never do anything unless it is connected to another half cell which will complete the circuit. This is the same as saying that a reduction reaction cannot happen without an opposite oxidation process occurring.

If we connect up another half-cell, the electrons can flow around the circuit and the EMF of the cell can be measured using a voltmeter. The relevant half-equations are shown in the diagram below.

Click on the image to enlarge

However, the EMF of this cell is now a combination of the two half-cells, the copper and the zinc. So how do we get a measurement of the EMF for just one cell?

The answer is to decide on a cell that will be a reference point for all other half-cells. This reference half-cell can then be used to measure the EMF (or cell potential) of any other cell. This is what chemists do, and the cell that has been chosen to be a reference point is the hydrogen half-cell.

The hydrogen half-cell is set up by having hydrogen gas in contact with a solution that has a hydrogen ion (H⁺) concentration of 1.00 mol L^-1.

This cell is given a value of zero, so if any other half-cell is connected to the hydrogen half-cell, the EMF measured is only dependant on the half-cell that is being measured. In this way the hydrogen half-cell is used as a reference point.

The example below represents the copper half-cell linked to the hydrogen half-cell reference. The salt bridge, which contains ions that are not involved in the reactions, allows an electric current to flow between the two solutions. The reading on the voltmeter will give the electrode potential for this cell.

Click on the image to enlarge

Standard conditions

If we are comparing the relative values of electrode potentials, the values must be measured under the same conditions. These conditions are defined as a concentration (of the ions involved in the redox reaction) of 1.00 mol L^-1, the temperature should be 25°C and if any gasses are involved, the pressure will be 1.00 atm (101.3 kPa).

Summary

The result of all this is a list of half-equations, each with its own standard reduction potential. These values have been calculated using the hydrogen half-cell as a reference point.

This list can be used to predict how the various substances will behave in a vast range of redox reactions. So get hold of your copy of the list, and be prepared to put it to use in the next section.