3.2.1 – Redox reactions
📌 Oxidation
- Oxidation can be defined one of three ways :
- The gain of oxygen
- The loss of hydrogen
- The loss of electrons
- It can also be defined as an increase in oxidation state
📌 Reduction
- Reduction can be defined one of three ways :
- The loss of oxygen
- The gain of hydrogen
- The gain of electrons
- It can also be defined as an decrease in oxidation state
📌 Redox reactions
- Redox reactions are ones in which both oxidation and reduction occur
- Redox reactions use changing oxidation states to analyse which reactants are oxidised during the reaction and which are reduced
- A few important things to remember are :
- Transition metals have variable oxidation states
- The oxidation state of elements in their standard state is 0
- The final oxidation state of a compound is 0 unless indicated otherwise by a charge on an ion
📌 Examples
Using the following examples we can understand how to identify which species are reduced and which are oxidised
EXAMPLE ONE : 4Fe(s) + 3O2(g) + 6H2O(l) → 4Fe(OH)3 (s)
Fe (s) has an oxidation state of 0, as does O2(g)
In H2O, H has an oxidation state of 1+ and O has an oxidation state of 2-
In Fe(OH)3 (s) we know that the OH ion has an oxidation state of 1- and there are 3x OH ions indicating a total 3- charge
Therefore to balance this, the Fe ion must have a charge of 3+ in order for the oxidation state of the compound to be 0
Given that the oxidation state of Fe is 0 in the left hand side (reactants) and 3+ in the right hand side (products), we can say it has had an increase in oxidation state (aka it has been oxidised)
Conversely, the O2(g) goes from 0 (reactants) to -2 (products)
We can say there is a decrease in oxidation state (aka it has been reduced)
EXAMPLE TWO : Fe2O3(s) + 2Al(s) → 2Fe(s) + Al2O3(s)
Al (s) has an oxidation state of 0, as does 2Fe(s)
In Fe2O3, O has an oxidation state of 2-
Given that there are 3x O, the overall charge of oxygen is 6
There are 2 Fe ions in the compound
Therefore 2x + (-6) = 0
The oxidation state of Fe is 3+
In Al2O3, O has an oxidation state of 2-
Given that there are 3x O, the overall charge of oxygen is 6
There are 2 Al ions in the compound
Therefore 2x + (-6) = 0
The oxidation state of Al is 3+
Given that the oxidation state of Fe is 3+ in the left hand side (reactants) and 0 in the right hand side (products), we can say it has had an decrease in oxidation state (aka it has been reduced)
Conversely, the Al(s) goes from 0 (reactants) to 3+ (products)
We can say there is a increase in oxidation state (aka it has been oxidised)
⭐️ always remember to balance equations before trying to identify oxidation/reduction as it will help later when finding out how many electrons are transferred during these reactions
📌 Oxidising and reducing agents
- Substances that aid in oxidation are known as ‘oxidising agents’. These substances are the ones that accept electrons during the reaction (ie the ones that are REDUCED)
- Similarly, substances that aid in reduction are known as ‘reducing agents’. These substances lose electrons during the reaction (ie the ones that are OXIDISED)
- It is important to note that if a compound is ‘reduced’, only certain ions in the compound can be actually classified as the reducing agent
- Take the example (example 2) above: Al is oxidised making it the reducing agent, but specifically Fe3+ is reduced. Thus Fe3+ is the reducing agent NOT the entire compound of Fe2O3
📌 Nomenclature
- IUPAC nomenclature dictates that the charge of an ion in a certain compound must be written in roman numerals
- For example, Fe can exist as both Fe2+and as Fe3+. In a compound such as iron oxide, we must then make it clear which Fe ion is present by writing it as iron (II) oxide OR iron (III) oxide
- A similar principle can be applied to other ions like Cu (exists as both +1 and +2) or Mn (exists as both +4 and +7)
- Although this nomenclature is mainly put in place to distinguish between ions that have variable oxidation states, it can be helpful to add the oxidation state to many compounds for easier readability