Chapter 18 Outline

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.

18.1 Pulling the Plug on the Power Grid


18.2 Balancing Oxidation-Reduction Equations

  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 reactions with the 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.

  5. Balancing reactions with 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.

  6. Balancing reactions with if the reaction is in a basic solution.
    1. Change all the H1+ (aq) to H2O.
    2. Change all the H2O to OH1- (aq)

  7. Check that the reaction is balanced for elements.
  8. Check that the reaction is balanced for charge.


18.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions


18.4 Standard Electrode Potentials


18.5 Cell Potential, Free Energy, and the Equilibrium Constant


18.6 Cell potential and Concentration


18.7 Batteries: Using Chemistry to Generate Electricity


18.8 Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity


18.9 Corrosion: Undesirable Redox Reactions


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