Oxidation and reduction are key concepts in electrochemistry. Chemistry involving oxidation and reduction is typically referred to as Redox chemistry. Oxidation involves the loss of electrons from a species and reduction involves the gain of electrons. Because we are never creating new charges out of nothing (aside from nuclear chemistry), we always have oxidation in conjunction with reduction.
Therefore in any Redox reaction, one species is always oxidized and another is reduced. We can't have one without the other. Finally, the number of electrons involved in this process must be balanced. That is, the number of electrons lost by one species needs to be the same as the number gained by the other. You can't simply build up a beaker full of electrons.
Introduction to ElectrochemistryIn order to identify Redox chemistry, chemists have developed a system by which we can think about the "number of electrons" a certain element "has". This is an idea that simply helps sort out the oxidation and reduction process.
In every chemical species, we can assign an oxidation number (or oxidation state) to each element. This can be thought of as the number of electrons that "belong" to that element compared to the number of valence electrons the element has. It is important to note that this is an idea that helps us to think about and classify the chemistry. Within a compound the electrons don't have any labels and don't belong to anyone one element. This is merely an accounting trick for us to use as chemists.
Assigning oxidation numbers is very straight forward for atomic ions. Take Na+ as an example. Normally sodium has one valence electron. As an ion it has zero valence electrons. We could then state that the oxidation number (or oxidation state) of Na+ is plus one. It is "missing" one electron. Similarly the oxidation state of Cl- is negative one. It has an "extra" electron. So for any atomic ion, the oxidation state is simply the charge. So in the compound NaCl, the oxidation number for Na is +1 and the oxidation number for Cl is -1.
What about polyatomic ions or compounds? For these we need to consider the oxidation state for each element in the chemical species. Here the oxidation number can be thought of as the charge that atom would have if the chemical species were broken up. It is as if we are thinking about all compounds as being ionic compounds.
For this we have a series of rules.
From there you can work out any element in any compound. General rule. Assign H and O first. The rest typically fall out form there.
What is the oxidation state of N in NH3? H is +1, the whole compound is neutral, so N must be -3 since there are three hydrogen at +1 each.
Redox NumbersIf something is oxidized then its oxidation number goes up. Oxidation is the loss of electrons.
If something is reduced then its oxidation number goes down. Reduction is the gain of electrons.
There are many ways to remember this. Pick one and remember it forever.
OIL RIG. Oxidation Is Loss. Reduction Is Gain.
LEO says GER. Lose Electrons Oxidation. Gain Electrons Reduction.
For example, let's look at the reaction.
\[\rm{2Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)}\]
In this example, the oxidation number of iron starts at zero and then goes to +3. The oxidation number of oxygen starts at zero and goes to -2. The iron is oxidized. The oxygen is reduced. Note: the name oxidation comes form the fact that reaction with oxygen leads to oxidation.
We can further put "names" on the compounds involved in the Redox reaction to identify their "actions". In this reaction, we would call oxygen the "oxidizing agent" since it is oxidizing the iron. The iron would be the "reducing agent" since it is reducing the oxygen. There is always oxidation with reduction, so we always have an oxidizing agent and a reducing agent. However, often in reactions we only identify one of them (the other is then identified by default).
Redox Agents
In the reaction CO + H2O --> CO2 + H2, what is being reduced?
(mouse over choices to get answer)
There are many methods to balance redox reactions, but it is best to pick a method and approach it systematically.
Here, one method (the half reaction method) will be presented. If another methods works better for you, then great. If not, learn this one and practice it. An important idea is that balancing Redox reaction is different in acidic versus basic conditions. This is because the reaction involves either H+ or OH- which will affect both the elements and the charge.
Here is the half reaction method in acidic solutions.
Step 1: Determine the oxidation numbers of the elements in the reaction and identify the oxidation and reduction reactions.
Step 2: Identify the oxidation and reduction half reactions and write the equations for each.
Step 3: For each half-reaction, balance all elements except for hydrogen and oxygen. Then balance the oxygen by adding H2O. Then balance the hydrogen using H+. Finally balance the charge in the half-reaction using electrons.
Step 4: Multiply one or both of the balanced half-reaction by whole numbers to equalize the number of the electrons in each half-reaction.
Step 5: Double check that all the elements and charge are balanced.
The half reaction method in basic conditions is nearly identical. The steps are the same up through step 5.
Step 6: We now have a balanced equation except it has H+ and we want to get to basic conditions. So now, add enough OH- to each side of the equations to neutralize away any H+. One side of the equation should now have the same number of H+ and OH-. These form water, H2O molecules. Now, cancel out any waters that appear on both sides of the equation.
Step 7: Double check that all elements and charge are balanced.
Balancing Redox Reactions in Acidic Solution© 2013 mccord/vandenbout/labrake