Redox

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Redox Reactions

Redox reactions are important in many areas of life (e.g. breathalysers). These reactions involve the loss and gain of electrons.

OIL RIG is a good mnemonic for remembering the difference between oxidation and reduction:
- Oxidation Is Loss of electrons.
- Reduction Is Gain of electrons.

An oxidising agent oxidises something else.
- It takes electrons from another compound.
- This means it gains electrons itself.
- So an oxidising agent is itself reduced.

A reducing agent reduces something else.
- It gives electrons to another compound.
- This means it loses electrons itself.
- So a reducing agent itself is oxidised.

Oxidation states are crucial tools for understanding redox reactions.

An oxidation state shows how many electrons an atom has gained or lost.
The concept of oxidation state is related to electronegativity.
- Electronegativities can be used to work out oxidation states.
There are a set of rules to assigning oxidation states - these will be unpacked in the next slide.

To assign oxidation states in a compound with multiple elements, we pretend every bond is ionic (even when they're definitely not)!
- We ask the question 'which element is going to take the electron pair?'
- The answer is, the most electronegative one!
This gives us our first rule:
- In a compound with fluorine, fluorine's oxidation state is always -1.

The oxidation state of oxygen is always -2.
- Unless you have a compound of oxygen and fluorine. The fluorine rule takes priority.
- Another exception is in a peroxide (e.g. Na₂O₂).

Hydrogen always has an oxidation state of +1.
- Except in metal hydrides e.g. NaH, where it is -1.

In a compound ion, the overall oxidation state is equal to the charge on the ion.
In a simple ion, the oxidation state is just the charge on the ion.
In a pure element, the oxidation state is zero. This includes atoms like He, but also molecules like H₂.
Oxidation states are commonly represented by Roman numerals e.g. Fe(III) sulfate means (Fe³⁺)₂(SO₄^2-)₃.

You will already be familiar with balanced equations. When the reaction is a redox reaction, we can make some further adjustments.

An ionic half equation shows either reduction or oxidation.
- An example is: O₂ + 4e^- $\rightarrow$ 2O^2-
- This is the ionic half equation for the reduction of O₂ to 2O^2-
If you have ionic half equations for both a reduction process and an oxidation process, you can add the two to generate a full balanced equation.

The ionic half equation for the oxidation of Ti $\rightarrow$ Ti⁴⁺ is:
- Ti $\rightarrow$ Ti⁴⁺ + 4e^-
The ionic half equation for the reduction of Cl₂ $\rightarrow$ to 2Cl^- is:
- Cl₂ + 2e^- $\rightarrow$ 2Cl^-
When adding the two equations together, you must balance the number of electrons on either side so that they cancel out.
- 2Cl₂ + 4e^- $\rightarrow$ 4Cl^-
- Ti $\rightarrow$ Ti⁴⁺ + 4e^-
So the full balanced equation is:
- Ti + 2Cl₂ $\rightarrow$ Ti⁴⁺ + 4Cl^-