3.5.3 Redox Equilibria - Redox equations
The oxidation state is the apparent valency of an atom within a compound. It is considered as if the element were bonded ionically, to allow the apparent number of electrons gained or lost to be assessed.
To find the oxidation state of an element within a species some very simple rules are applied:
- 1 The sum of all the oxidation states in a species must add up to the charge on the species.
- 2 The oxidation number of an uncombined element is always zero (0).
- 3 Oxygen in compounds and ions has an oxidation state of -2, unless in peroxides, or superoxides.
- 4 Reactive metals oxidation state is the same as the group number, eg Na = +1, Mg = +2, etc..
- 5 When two atoms are attached together the more electronegative takes the negative oxidation state.
Oxidation and reduction
Both of these processes involve a transfer of electrons.
- Oxidation is loss of electrons
- Reduction is gain of electrons
If an element gets oxidised its oxidation state increases, becomes more positive. If an element gets reduced its oxidation state decreases, i.e. gets more negative.
Although an oxidation reaction MUST also involve a reduction reaction (the electrons must have both a source and a target), the overall process can be broken down into so-called 'half equations' showing just the oxidation process and just the reduction process.
Species tend to always behave in the same way, so a knowledge of the half-equations allows us to predict the actual redox (reduction + oxidation) processes.
For example: The reaction:
Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)
is a redox reaction. We know this because the oxidation state of magnesium changes from zero (the element) to +2 in MgCl2. It has increased its oxidation state (more positive) therefore it has been oxidised. We can write this as an oxidation half-equation:
Mg Mg2+ + 2e
The hydrogen changes oxidation state from +1 in HCl to zero in the element. It has therefore been reduced (less positive oxidation state). This can be written as a reduction half-equation.
2H+(aq) + 2e H2(g)
Transition metals are commonly involved in redox treactions because they have variable oxidation states. To change from one state to another requires the loss or gain of electrons. Hence, transition metal ions can be reducing agents (in low oxidation states) or oxidising agents (in high oxidation states)
Some common transition metal redox systems:
- Fe2+(aq) Fe3+(aq) + 1e
- MnO4-(aq) + 8H+(aq) + 5e Mn2+(aq) + 4H2O(l)
Iron(II) salts are able to behave as reducing agents, providing electrons to reduce strong oxidising agents. Manganate(VII) ions are good oxidising agents, absorbing electrons to become Mn(II) ions.
Example: The reaction between iron(II) salts and potassium manganate(VII)
Fe2+(aq) Fe3+(aq) + 1e
MnO4-(aq) + 8H+(aq) + 5e Mn2+(aq) + 4H2O(l)
To combine the half-equations first the electrons must be equalised:
5Fe3+(aq) + 5e