Chemistry Unit 1.5 - Oxidation and Reduction

Notes on oxidation and reduction

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Chemistry Unit 1.5: Oxidation and Reduction
An ionic half equation is one that shows electron transfer. In a redox reaction, there are two ionic
half-equations, one represents oxidation and the other reduction.
If an ionic half equation shows electrons on the left-hand side it is oxidation, if it is on the right it is reduction.
Two ionic half-equations can be combined to give overall ionic equations by cancelling the electrons from
both half-equations.
Ionic half equations, when put together, form ionic equations. To construct a half equation, take an element,
compound or ion which has been oxidised or reduced, and write the change to it - e.g.
Next, the atoms need to be balanced. The only things allowed to be added to the equation are water and
hydrogen ions. For the example above:
Finally, the charges need to be balanced, by adding electrons:
This is the number assigned to an atom or ion to describe its relative state of oxidation or reduction. They are
useful as a concept to help identify redox and disproportionation, where one element experiences both
oxidation and reduction, reactions.
1. The oxidation number of the atoms of elements in the uncombined state is 0.
2. In neutrally charged molecules, such as H2O, the sum of the oxidation numbers is 0.
3. In ions such as NO3- the sum of the oxidation numbers is equal to the overall charge on the ion.
4. In any substance the more electronegative element has the negative oxidation number whilst the
less electronegative element has the positive oxidation number.
5. The oxidation number of fluorine is always -1.
6. The oxidation number of oxygen is -2 except in peroxides and when combined with fluorine.
7. The oxidation number of hydrogen is +1 except in metal hydrides.
8. The oxidation number of chlorine is -1 except when combined with oxygen or fluorine.


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