Chapter 20 Outline

Electrochemistry and Redox Reactions

1. Oxidation-Reduction Reactions. Oxidation-Reduction reactions (redox reactions) involve changing the oxidation state. An example is changing from Cu-I to Cu-II, or from Fe to Fe-III. This is a very important class of reactions, but they are a bit difficult to understand.
2. Oxidation State. The first part of figuring out these reactions is to determine the oxidation state (charge) of each element. We have done some of this when balancing equations and when discussing nomenclature. The basic idea is to start with the most obvious parts and then balance out the charge. The most obvious parts are the elements closest to the edge (left or right) of the periodic table. Keep in mind the following:
1. Oxygen is usually -2 (except in peroxides ie H2O2, where it is -1)
2. Hydrogen is usually +1 (except with metals like NaH where it is -1)
3. Halogens are usually -1
4. Alkali metals are usually +1
5. Alkali earth metals are usually +2
3. Redox reactions
1. Reduction is gaining electrons ie: Fe-III to Fe-II
2. Oxidation is loosing electrons ie: Fe-II to Fe-III (you can remember this because when iron is rusts it is oxidized from Fe to Fe-II and Fe-III.)
3. Just to keep things confusing then, an oxidizing agent is something that causes oxidation (like oxygen causes rust), and a reducing agent is something that causes reduction. In the process the oxidizing reagent is reduced (the other thing is oxidized). Or the reducing agent is oxidized (the other thing is reduced).

4. Balancing redox reactions. This can be tricky because you need to take into account the transfer of electrons. SOMETIMES this is simple and you can come up with the answer by the "oxidation state method" discussed on page 154 and 155.
5. For some more complex redox reactions the oxidation state method will not work and the half-reaction method is useful. This is nicely outlined on page 156, and an excelent set of examples are given on page 157-159. It is important to note that these reactions are balanced differently for acidic and basic solutions.
6. Redox Titrations. Same idea here as acid-base titrations, except now we are looking at redox reactions instead of neutralization reactions.

Steps to balancing a redox reaction

1. The Oxidation Number Method
1. Determine the oxidation numbers for all elements in the reactants and the products.
2. Identify which species is being oxidized and which species is being reduced.
3. Determine how many electrons are lost by each atom of the oxidized species.
4. Determine how many electrons are gained by each atom of the reduced species.
5. Balance the number of electrons gained and lost.
6. Balance the reaction is balanced for elements. Add H₂O and H¹⁺ (aq) as necessary.
7. Check that the reaction is balanced for charge.
2. The Half Reaction Method
1. Identify which compound is oxidized.
1. Write the oxidation half reaction.
2. Balance the half reaction for elements.
3. Add H₂O and H¹⁺ (aq) as necessary.
4. Balance the half reaction for charge by adding electrons as products.
2. Identify which compound is reduced
1. Write the reduction half reaction
2. Balance the half reaction for elements.
3. Add H₂O and H¹⁺ (aq) as necessary.
4. Balance the half reaction for charge by adding electrons as reactants.
3. Balance the number of electrons from the oxidation and the reduction.
4. Add the two reactions together.
5. Check that the reaction is balanced for elements.
6. Check that the reaction is balanced for charge.
3. If the reaction is in a basic solution.
1. Change all the H¹⁺ (aq) to H₂O.
2. Change all the H₂O to OH^1- (aq)
4. Check that the reaction is balanced for elements.
5. Check that the reaction is balanced for charge.

Please send any comments, corrections, or suggestions to svanbram@science.widener.edu.

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Last Updated Friday, May 25, 2001 2:11:17 PM