1. Oxidation-reduction reactions & electrochemistry
Chapters 19&20 Holt, Rinehart and Winston
NGSS: PS1.A; PS1.B; HS-PS1-7; HS-PS1-8

2. Redox reactions
Chemical reactions in which electrons are transferred from one atom to another are called oxidation-reduction reactions.
Oxidation is loss of electrons, reduction is gain of electrons. OILRIG or LEO the lion says GER.
The substance that oxidizes another substance by accepting it electrons is called an oxidizing agent (it is reduced). The substance that reduces another substance by losing electrons is called a reducing agent (it is oxidized).

3. Redox and Electrochemistry videos
Redox reactions (Bozeman) 11:40 Redox rxns (Crash Course chemistry) 11:12 Redox reactions (Brightstorm science) 3:32 Oxidation number (Brightstorm science) 5:12
Half-reactions (Brightstorm science) 3:30
How to balance in acidic solution (Khan Academy) 12:03
How to balance in basic solution (Khan academy)  
14:01

4. Rules for determining oxidation numbers
The oxidation number of an uncombined atom is zero. This also applies to elements that exist in polyatomic states.
The oxidation number of a monatomic ion is equal to the charge on the ion.
The oxidation number of the more electronegative atom in a molecule or a complex ion is the same as the charge it would have it if were an ion.
The most electronegative element, fluorine, always has an oxidation number of –1 when it is bonded to another element.

5. Rules for determining oxidation numbers, CONt.
The oxidation number of oxygen in compounds is always –2, except in peroxides, such as H2O2, where it is –1. When it is bonded to fluorine, the only element more electronegative than oxygen, the oxidation number of oxygen is +2.
The oxidation number of hydrogen in most of its compounds is +1. The exception is when it is bonded to less electronegative metals to form hydrides, such as LiH, NaH, CaH2, and AlH3. In these compounds, H has an oxidation number of –1.

6. Rules for determining oxidation numbers, CONt
7. The metals of groups 1A and 2A and aluminum in group 3A form compounds in which the metal atom always has a positive oxidation number equal to the number of its valence electrons (+1, +2, and +3 respectively).
8. The sum of the oxidation numbers in a neutral compound is zero.
9. The sum of the oxidation numbers of the atoms in a polyatomic ion is equal to the charge on the ion.

7. Balancing redox reactions
In acid solution
Identify the two half reactions
Balance each ½ reaction
All elements except H and O
Balance O by adding H2O
Balance H by adding H+
Balance the charge by adding e- to the side with a greater + charge.
Multiply each ½ reaction by an integer so that the e- are balanced.
Add the two-half reactions.
Check the numbers of atoms and the charge.

8. Cu (s) + NO3- (aq)  Cu2+ (aq) + NO2 (g)
Examples
Cu (s) + NO3- (aq)  Cu2+ (aq) + NO2 (g)
Mn2+ (aq) + NaBiO3 (s)  Bi3+ (aq) + MnO4- (aq)

9. Examples Cu (s)  Cu2+ (aq) + 2e-
Cu (s) + NO3- (aq)  Cu2+ (aq) + NO2 (g)
Cu (s)  Cu2+ (aq) + 2e-
2(e- + 2H+ + NO3- (aq)  NO2 (g) + H2O)
Cu (s) 4H+ + 2NO3- (aq)  Cu2+ (aq) + 2NO2 (g) + 2H2O

10. Examples Mn2+ (aq) + NaBiO3 (s)  Bi3+ (aq) + MnO4- (aq)
2(4H2O + Mn2+ (aq)  MnO4- (aq) + 8H+ + 5e-)
5(2e- + 6H+ + NaBiO3 (s)  Bi3+ (aq) + Na+ + 3H2O)
2Mn2+ (aq) + 14H+ + 5NaBiO3  2MnO4- + 5Bi3+ + 5Na+ + 7H2O

11. NO2- (aq) + Al (s)  NH3 (aq) + Al(OH)4- (aq)
In basic solution
Follow the same rules for acid solution but use OH- and H2O instead of H+ and H2O. Neutralize H+ with OH- (initially balance as if in acid solution).
Example
NO2- (aq) + Al (s)  NH3 (aq) + Al(OH)4- (aq)

12. NO2- (aq) + Al (s)  NH3 (aq) + Al(OH)4- (aq)
In basic solution
Follow the same rules for acid solution but use OH- and H2O instead of H+ and H2O. Neutralize H+ with OH- (initially balance as if in acid solution).
Example
NO2- (aq) + Al (s)  NH3 (aq) + Al(OH)4- (aq)
7H+ + 7OH- + NO2- (aq)  NH3 (aq) + 2H2O + 7OH-
(think of as H+ and OH-)
2(4OH- + Al (s)  Al(OH)4- (aq) +3e-)
NO2- (aq) + 2Al(s) + 5H2O + OH-  NH3 + 2Al(OH)4-

13. Disproportionation
The process whereby a substance can act as an oxidizing agent and reducing agent.
2H2O2  2H2O + O2

14. Electrochemistry
An Electrochemical cell is an apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction.
Voltaic (aka galvanic) cells are types of electrochemical cells that convert chemical energy to electrical energy by a spontaneous redox reaction.
Oxidation takes place at the anode.
Reduction takes place at the cathode