Half Equations

Learning outcomes:

Demonstrate knowledge, understanding, and application of:

Rules for calculating oxidation number.
Balancing half equations by adding electrons, H+ ions, and water.
Combining half equations

Let’s first recap some key definitions:

Oxidation = Loss of electrons or increase in oxidation number
Reduction = Gain of electrons or decrease in oxidation number
Oxidising agent = Gain of electrons – themselves reduced
Reducing agent = Loss of electrons – themselves oxidised.

***Make sure you have fully understood Oxidation and Reduction before revising this section***

Oxidising and reducing agents:

In a redox reaction, there will always be an oxidizing agent and reducing agent.

For example: Let’s have a look at what is happening to the reaction between Ag+ ions and Cu as shown below:

Silver’s oxidation number goes from +1 to 0 therefore Ag+ has been reduced. = Oxidising agent. Therefore, Silver is the oxidising agent.
Copper’s oxidation number goes from 0 to +2 so, therefore, Cu has been oxidised = reducing agent

Let's first have a look at some examples of balancing simple redox equations by adding electrons

Oxidation and reduction in word equations:

In an exam, they may give you a word equation from which you need to deduce which element is oxidised and reduced.

Let’s go through an example:

Study Tip:

Reduction is the decrease in oxidation number. (e.g. 0 to -1)

Oxidation is the increase in oxidation number. ( e.g. -1 to 0)

Combining half equations together:

So now we have both of our half equations balanced, we can now combine them to make one full equation. For this, we need to compare both of the half equations and cancel out similar species.

As you can see, both half equations have 2 electrons present on both sides of the equation. We can cancel these out like this…

Now the spectator species have been canceled out, all we need to do now is to combine all the reactants together and all the products together in one equation like this..