Redox Reactions

Redox reactions:

Redox reactions also known as oxidation and reduction reactions are reactions in which one species is reduced and another is oxidized at the same time because when there is a reaction if one species is oxidized another species must be reduced. Therefore, this means the oxidation state of the species involved must change and therefore means that electrons must be transferred between species.

How do redox reactions work?

Using an example of a chemical reaction we can see if it is a redox reaction for an example:

2Fe2O3 (s) + 3C (s) —> 4Fe (s) + 3CO2 (g)

For a chemical reaction to be a redox reaction there must be a transfer of electrons you can work this out by finding the oxidation states:

Working out oxidation states reminder (worked example)

2Fe2O3 (s) + 3C (s) —> 4Fe (s) + 3CO2 (g)

For 2Fe2O3 (fe2) has the oxidation state of +3 and O3 has the oxidation state of -2 (because oxygen is in group 6 it needs to gain two electrons when reacted therefor it has a -2 charge but because its (O3) there’s 3 atoms of oxygen giving it an overall charge of -6 which means the Fe needs a oxidation number of -6 to cancel it out but since there's two of them (fe2) you would divide the +6 charge into 2 so it would have a +3 charge.

3C and 4Fe both have an oxidation state is 0 because there on their own (combined) they have not been oxidized or reduced.

For 3CO2 the oxygen would have an oxidation state of -2 (because it’s in group 6 and needs to gai two electrons) and carbon would have an oxidation state of +4 (because its in group 4 and the overall compound has no charge and therefor it would lose the 4 electrons (+4) to cancel out the 2 oxygens which both have a -2 charge.)

Once you have worked out the oxidation numbers of the species in the reaction you can look at the individual elements to see if they have been oxidized or reduced.

Because there is a change in the oxidation numbers of carbon and iron from the reactants to the products this is a redox reaction and there has been an electron transfer between the iron and carbon species.

The carbon when from a charge of 0 to a charge of +4 therefore its lost electrons hence why it has a positive charge therefor this show that it has been oxidized and would also be the reducing agent.

The iron went from a +3 charge to a charge of 0 therefor its gained electrons this shows that it has been reduced and would also be the oxidizing agent in this reaction.

As part of the OCR B specification, you must be able to understand and apply this knowledge for redox reactions of s-, p- and d-block elements and their compounds.

Balancing redox reactions using oxidation states:

First work out the oxidation numbers of each of the species in the reaction making sure you remember their charges.
Then work out which species have been oxidised and which ones have been reduced.
Work out the change that the species have.
Notice that if electrons gained does not equal electrons lost you need to balance the redox equation.
Use coefficients to make sure the charges are balanced and cancel out each other.

To understand this more, we can use an example:

Fe2O3 + CO → Fe + CO2

Find the oxidation numbers of the species:

2. Find which species have been oxidized and which have been reduced (carbon went from a +2 to +4 so it’s been oxidized (-2)) (iron went from a +3 to a 0 so it’s been reduced) (oxygen is neutral overall)

3. Notice that electrons gained do not equal electrons lost so it needs balancing.

4. Use coefficients to make them balance each other. Finding a multiple is a good way to start the -2 and the +3 need to be balanced and they are both multiples of six. So the oxidation number -2 can be multiplied by 3 and the oxidation number +3 can be multiplied by 2.

5. Fe2O3 + 3CO --> 2Fe + 3CO2

Disproportionation reactions:

A disproportionation reaction is when one species in a reaction undergoes both oxidation and reduction.

This can happen in many reactions and we can use the hydrogen peroxide reaction as an example because in this reaction the oxygen undergoes disproportionation, and this reaction is commonly used to show disproportionation.

Once you have worked out the oxidation states, we can see that on the reactants side, the oxygen has an oxidation state of -1 and on the product side the oxygen has an oxidation state of 0 and -2.

Therefor the oxygen when from a -1 charge to a -2 charge meaning it gained oxygen and has therefor been oxidized. However, it has also gone from a -1 to a 0 charge meaning it has also lost electrons in this reaction meaning tis also been reduced.