Simple half-equations

Oxidation and reduction are best defined within organic chemistry in terms of the electrons transferred.

• Oxidation is where electrons are lost. The species is then said to have been oxidised.
• Reduction is where electrons are gained. The species is then said to have been reduced.

Processes in which electrons have been lost or gained are called half-equations. For example:

• Oxidation: Na ? Na⁺ + e
• Reduction: Cl₂ + 2e ? 2Cl^–

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Oxidation numbers

The oxidation number of an atom is the charge that would exist if its bonding was completely ionic.

In a simple ion the oxidation number of the atom is the charge of the ion. For example:

• Na⁺ has an oxidation number of +1
• Ca²⁺ has an oxidation number of +2
• Br^– has an oxidation number of -1
• O^2- has an oxidation number of -2

In a molecule or compound the sum of oxidation numbers on the atom amounts to zero. So, for SO₃ the oxidation number for:

• S = +6
• O = -2

Therefore: +6 + 3(-2) = 0

In complex ions the sum of oxidation numbers on the atoms amounts to the charge of the ion overall. So, for SO₄^2- the oxidation number of:

• S = +6
• O = -2

Therefore: +6 + 4(-2) = -2

In elements which are in their standard state the oxidation number of each atom amounts to zero. For example, Na and O₂ have an oxidation number of zero.

A lot of atoms can exist in a number of oxidation states, like Cl or N. The oxidation number of these atoms can be calculated by making the assumption that the oxidation number of the other atom is fixed. Generally, O at -2.

The following use a specific figure for the oxidation state in their compounds:

• All group I: +1
• All group II: +2
• Aluminium: +3
• Fluorine: -1
• Hydrogen: +1 (unless bonded to a metal, silicon or boron where it uses -1)
• Oxygen: -2 (unless bonded to a group I metal, group II metal, or hydrogen where it generally uses -1, or fluorine where it uses +2)

The oxidation numbers of other atoms in their compounds vary.

The oxidation numbers of an atom changes during oxidation or reduction.

• In oxidation electrons are lost so the oxidation number increases.
• In reduction electrons are gained so the oxidation number decreases.

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More complex half-equations

Half-equations become more complicated when complex ions or molecules are involved. In these cases, oxidation numbers become a particularly useful tool.

Half-equations can be balanced using two methods.

Method 1: you can see immediately whether an oxidation or reduction process is taking place.

• First of all, identify which atom is being oxidised or reduced. Balance out the reaction to ensure that the atom is of the same amount both sides.

? Atoms of O can be balanced by adding water.

? Atoms of H can be balanced by adding H⁺.

• Insert how many electrons are being either lost (on the right) or gained (on the left).

? number of electrons gained/lost = change in oxidation number x number of atoms changing oxidation number

Method 2: this is an easier method to use and does not include oxidation numbers.

• First of all, identify which atom is being oxidised or reduced. Balance out the reaction to ensure that the atom is of the same amount both sides.

? Atoms of O can be balanced by adding water.

? Atoms of H can be balanced by adding H⁺.

• Add in the necessary number of electrons so that the charge on both sides is equal.

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Redox reactions

In half-equations you are considering the loss and gain of electrons. However, electrons can only be transferred from one species to another; they cannot be created or destroyed. Therefore, it follows that oxidation and reduction reactions occur simultaneously. This is known as a redox reaction.

You can derive a redox reaction by combining a oxidation half-equation and a reduction half-equation so that the electrons gained equals the electrons lost.

Oxidising agents and reducing agentAn oxidising agent is the species which accepts the electrons in an oxidation reaction. For example, Al³⁺ and Cl₂.

A reducing agent is the species which donates the electrons in a reducing reaction. For example, O^2- and Na

Therefore, a redox reaction can described as one in which electrons are transferred from a reducing agent to an oxidising agent.

Disproportionation

Some substance behave as both oxidising and reducing agents and readily undergo oxidation and reduction reactions. In fact, they are capable of doing both simultaneously, a process known as disproportionation. Examples include H₂O₂ and ClO^–.