Combustion is an example of a type of reaction called oxidation. In an oxidation reaction, a substance gains oxygen. Metals and non-metals can take part in oxidation reactions. Metals react with oxygen in the air to produce metal oxides. For example, magnesium reacts with oxygen to produce magnesium oxide when it is heated in air:. So oxidation reactions need not involve oxygen. This redox reaction is actually the sum of two separate half-reactions a reduction half-reaction and an oxidation half-reaction.
Example 1. In the following redox reaction, which species is being oxidized? Which one is being reduced? Al s is being oxidized. In the following image, we see methane combusting to release energy. Combustion reaction of methane : This is an example of a combustion reaction, a redox process.
In some redox reactions, substances can be both oxidized and reduced. These are known as disproportionation reactions. One real-life example of such a process is the reaction of hydrogen peroxide, H 2 O 2 , when it is poured over a wound.
At first, this might look like a simple decomposition reaction, because hydrogen peroxide breaks down to produce oxygen and water:. The key to this reaction lies in the oxidation states of oxygen, however. Notice that oxygen is present in the reactant and both products. In H 2 O 2 , oxygen has an oxidation state of In H 2 O, its oxidation state is -2, and it has been reduced. In O 2 however, its oxidation state is 0, and it has been oxidized. Oxygen has been both oxidized and reduced in the reaction, making this a disproportionation reaction.
The general form for this reaction is as follows:. Every balanced redox reaction is composed of two half-reactions: the oxidation half-reaction, and the reduction half-reaction. For example, look at the following redox reaction between iron and copper:. In this reaction, iron is oxidized, and copper is reduced or, iron is the reducing agent , and copper is the oxidizing agent. We can split this reaction into two half-reactions. The oxidation half-reaction looks as follows:. This shows the oxidation of iron and the loss of two electrons.
Notice that this equation is balanced in both mass and charge: we have one atom of iron on each side of the equation mass is balanced , and the net charge on each side of the equation is equal to zero charge is balanced. This half-reaction explicitly shows the copper II ion gaining two electrons. Note again that the equation is balanced in mass and charge.
Now that we have our two balanced half-reactions, we can combine them to get the full redox reaction:. Adding the two halves of a redox reaction : These two halves of the reaction can be added like any other chemical equation. Once the equations are added, the electrons on each side cancel out. Note that the two electrons on each side of the equation cancel out. This is very important, because the final balanced equation for any redox reaction should never contain any electrons.
First, we need to split this reaction into its two half-reactions. We need to balance this equation by mass. The equation is already balanced in nitrogens, but not oxygens. You can balance oxygen by adding the appropriate number of water molecules:. Now the equation is balanced in oxygens, but not hydrogens. The equation is now balanced in mass, but not charge. To balance the charge, we will add two electrons to the right side of the equation:. The equation is now balanced for mass, and we need only balance for charge.
We therefore need to add 6 electrons to the right side of the equation to balance the charges:. Lastly, in order to get our full balanced redox equation, we need to add our half-reactions so that all the electrons cancel out. For this reaction, we can multiply the first half-reaction by Although this example seems intimidating, balancing redox reactions in acidic solution becomes much easier with careful practice.
For instance:. The half-reaction above is balanced for mass in acidic solution.
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