# Oxidation and Reduction

## Redox reactions

Now that we know how to determine the oxidation number of a compound, we will go on to look at how to use this knowledge in reactions.

### Oxidation and reduction

By looking at how the oxidation number of an element changes during a reaction, we can easily see whether that element is being oxidised (lost electrons) or reduced (gained electrons).

If the oxidation number of a species becomes more positive, the species has been oxidised and if the oxidation number of a species becomes more negative, the species has been reduced.

#### Tip:

The word species is used in chemistry to indicate either a compound, a molecule, an ion, an atom or an element.

We will use the reaction between magnesium and chlorine as an example.

The chemical equation for this reaction is:

$\text{Mg (aq)} + \text{Cl}_{2}\text{(aq)} \rightarrow \text{MgCl}_{2}\text{(aq)}$

As a reactant, magnesium has an oxidation number of zero, but as part of the product magnesium chloride, the element has an oxidation number of $$\text{+2}$$. Magnesium has lost two electrons and has therefore been oxidised (note how the oxidation number becomes more positive). This can be written as a half-reaction. The half-reaction for this change is:

$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^{-}$

As a reactant, chlorine has an oxidation number of zero, but as part of the product magnesium chloride, the element has an oxidation number of $$-\text{1}$$. Each chlorine atom has gained an electron and the element has therefore been reduced (note how the oxidation number becomes more negative). The half-reaction for this change is:

$\text{Cl}_{2} + 2e^{-} \rightarrow 2\text{Cl}^{-}$

### Definition: Half-reaction

A half reaction is either the oxidation or reduction reaction part of a redox reaction.

In the two half-reactions for a redox reaction the number of electrons donated is exactly the same as the number of electrons accepted. We will use this to help us balance redox reactions.

Two further terms that we use in redox reactions and that you may see are reducing agents and oxidising agents.

An element that is oxidised is called a reducing agent, while an element that is reduced is called an oxidising agent.

You can remember this by thinking of the fact that when a compound is oxidised, it causes another compound to be reduced (the electrons have to go somewhere and they go to the compound being reduced).

### Definition: Redox reaction

A redox reaction is one involving oxidation and reduction, where there is always a change in the oxidation numbers of the elements involved. Redox reactions involve the transfer of electrons from one compound to another.

## Optional Experiment: Redox reaction – displacement reaction

### Aim

To investigate the redox reaction between copper sulfate and zinc.

### Materials

• A few granules of zinc
• $$\text{15}$$ $$\text{ml}$$ copper (II) sulfate solution (blue colour)
• glass beaker

### Method

Add the zinc granules to the copper sulfate solution and observe what happens. What happens to the zinc granules? What happens to the colour of the solution?

### Results

• Zinc becomes covered in a layer that looks like copper.

• The blue copper sulfate solution becomes clearer.

$$\text{Cu}^{2+}$$ ions from the $$\text{CuSO}_{4}$$ solution are reduced to form copper metal. This is what you saw on the zinc crystals. The reduction of the copper ions (in other words, their removal from the copper sulfate solution), also explains the change in colour of the solution (copper ions in solution are blue). The equation for this reaction is:

$\text{Cu}^{2+}\text{(aq)} + 2e^{-} \rightarrow \text{Cu (s)}$

Zinc is oxidised to form $$\text{Zn}^{2+}$$ ions which are clear in the solution. The equation for this reaction is:

$\text{Zn (s)} \rightarrow \text{Zn}^{2+}\text{(aq)} + 2e^{-}$

The overall reaction is:

$\text{Cu}^{2+}\text{(aq)} + \text{Zn (s)} \rightarrow \text{Cu (s)} + \text{Zn}^{2+}\text{(aq)}$

### Conclusion

A redox reaction has taken place. $$\text{Cu}^{2+}$$ ions are reduced and the zinc is oxidised. This is a displacement reaction as the zinc replace the copper ions to form zinc sulfate.

## Optional Experiment: Redox reaction – synthesis reaction

### Aim

To investigate the redox reaction that occurs when magnesium is burnt in air.

### Materials

A strip of magnesium; bunsen burner; tongs; glass beaker.

### Method

#### Warning:

Do not look directly at the flame.

1. Light the bunsen burner and use a pair of tongs to hold the magnesium ribbon in the flame.

2. Hold the lit piece of magnesium over a beaker. What do you observe?

### Results

The magnesium burns with a bright white flame. When the magnesium is held over a beaker, a fine powder is observed in the beaker. This is magnesium oxide.

The overall reaction is:

$2\text{Mg (s)} + \text{O}_{2}\text{(g)} \rightarrow 2\text{MgO (s)}$

### Conclusion

A redox reaction has taken place. Magnesium is oxidised and the oxygen is reduced. This is a synthesis reaction as we have made magnesium oxide from magnesium and oxygen.

## Optional Experiment: Redox reaction – decomposition reaction

### Aim

To investigate the decomposition of hydrogen peroxide.

### Materials

Dilute hydrogen peroxide (about $$\text{3}\%$$); manganese dioxide; test tubes; a water bowl; stopper and delivery tube, Bunsen burner

#### Warning:

Hydrogen peroxide can cause chemical burns. Work carefully with it.

### Method

1. Put a small amount (about $$\text{5}$$ $$\text{ml}$$) of hydrogen peroxide in a test tube.

2. Set up the apparatus as shown above.

3. Very carefully add a small amount (about $$\text{0.5}$$ $$\text{g}$$) of manganese dioxide to the test tube containing hydrogen peroxide.

### Results

You should observe a gas bubbling up into the second test tube. This reaction happens quite rapidly.

The overall reaction is:

$2\text{H}_{2}\text{O}_{2}\text{(aq)} \rightarrow 2\text{H}_{2}\text{O (l)} + \text{O}_{2}\text{(g)}$

### Conclusion

A redox reaction has taken place. $$\text{H}_{2}\text{O}_{2}$$ is both oxidised and reduced in this decomposition reaction.