Oxidation Numbers and Ionic Equations

 

Reduction

The gain of electrons.

 

Oxidation

The loss of electrons.

 

OIL RIG

Oxidation is Loss           Reduction is Gain

 

 

Reduction and oxidation can be represented with half-equations. For example:

 

Na -> Na⁺ + e^-              Oxidation

Cl^- + e^- -> Cl                 Reduction

 

 

The Halogens are sometimes described as Oxidising Agents. This is because they gain an electron to form a full outer shell, which causes other species in a reaction to lose electrons. The loss of electrons is oxidisation.

 

The more reactive the halogen the better the oxidising agent. On the other hand the less reactive the halogen the better the reducing agent.

 

 

Oxidation Numbers

These are numbers which can be used to indicate what elements have been oxidised or reduced.

 

Rules

·         All elements have an oxidation number of 0 and all non-charged compounds also have a combined oxidation number of 0.

·         In compounds the oxidation numbers are equal the ionic number formed e.g. Cl in NaCl has an oxidation number of –1 and Na has an oxidation number of +1. This means overall the compound has an oxidation number of 0.

 

Example:

2Mg + O₂ -> 2MgO

  0        0        +2  -2

 

The Magnesium is oxidised as the oxidation number is increased indicating a loss of electrons.

The Oxygen is reduced as the oxidation number is decreased indication a gain of electrons.

 

 

Example of an Ionic Equation

When writing ionic equations you first write the symbol equation and then consider the oxidation numbers. The oxidation numbers which don’t change in the reaction are known as ‘spectator ions’ and can be ignored when writing the ionic equation.

 

Symbol equation:           Cl₂ + 2KBr^- -> 2KCl^- + Br₂

                                     0      +1 –2       +1 –2     0

Ionic equation:               Cl₂ + 2Br^- -> 2Cl^- + Br₂

 

 

NOTE: Cl^-₂ doesn’t exist, you must write 2Cl^-