Oxidation-reduction reactions are the final type of reactions that we will concern ourselves with. Oxidation-reduction reactions are also called redox reactions, and involve the transfer of electons. Below are two examples of redox reactions.
In example 1 the reaction is obviously ionic and is considered a redox reaction because there is a transfer of electrons, but in example 2, where is the transfer of electrons? To understand this we must first look at the oxidation states of each element. The oxidation states are numbers that allow us to keep track of the electrons in each element. To do this there are more rules. These rules are called the rules for assigning oxidations states.
Now lets look at example 2. The oxidation number of carbon in CO2 is +4 and the oxidation number of carbon in CO is +2. Carbon's oxidation number decreases by 2. Also the oxidation number of hydrgen in H2 is 0 and the oxidation number of hydrogen in H2O is +1. Hydrogen is increasing by 1. The transfer of electrons is why this is an redox reaction.
In redox reactions there are a few terms that we must review. The first is oxidation. Oxidation is the increase in oxidation numbers. Reduction is exactly the opposite. It is the reduction in oxidation numbers. The oxidizing agent is the element that is reduced and the reducing agent is the element that is oxidized.
In example 2, CO2 is being reduced. It is also the oxidizing agent. H2 is being oxidized and is the reducing agent.
It is important to balance redox reactions because if a reaction is producing 2 electrons then that same reaction must be using 2 electrons. For example, in example 2, 2 electrons are being produced by the carbon and only 1 of them are being used up the by the hydrogen. This is not even so this is why redox reactions must be balanced.
To balance redox reactions there are a few steps to follow.
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Unit 2 - Section 2
Unit 2 - Section 4