1. Oxidation and Reduction
Or, “Do you know where your electrons are?”

2. Definitions
Oxidation is the process of losing electrons (oxidation state becomes more positive)
Na  Na+ + 1e-
Reduction is the process of gaining electrons (oxidation state becomes more negative)
Cl + 1e-  Cl-

3. Losing Electrons Oxidation Gaining Electrons Reduction
Definitions
Losing Electrons Oxidation
goes
Gaining Electrons Reduction

4. Definitions
Oxidation Is Losing
Reduction Is Gaining

5. The number of electrons unequally shared in a covalent bond.
Oxidation state
Charge on an ion
Na+, Ca+2, O-2
The number of electrons unequally shared in a covalent bond.
H2O : H is +1, O is -2

6. Oxidation state assignment rules
Any element has oxidation number of zero
Oxygen has an oxidation number of -2, except in peroxides where it is -1
Hydrogen is +1 except in hydrides, where it is -1 – in HCl the H is +1, but in NaH it is -1
Nitrogen is -3 except with oxygen

7. Oxidation state assignment rules
Halogens are -1 except with oxygen or each other
All other oxidation numbers are assigned so that the sum of all the oxidation numbers equals the charge on the particle.
In examples not covered here the atom with greater electronegativity gets the negative charge.

8. Oxidation state assignment rules
NH3
H= +1, N= -3
NI3
N= -3, I = +1

9. Oxidation state assignment rules
NF3
N= +3, F= -1
H3O+
H= +1, O= -2

10. Oxidation state assignment rules
NO3-
O= -2, N= +5
Cr2O7-2
O= -2, Cr= +6

11. Redox reaction
Any reaction that results in a change of oxidation state for any reactant.
N2 + 3H2  2NH3
3Cu + 8HNO3  3Cu(NO3)2 + 2NO + 4H2O
-3, +1 +5 +2 +2

12. Redox Reaction 2Fe + 3CuSO4 3Cu + Fe2(SO4)3 0 +2 0 +3
Oxidizing agent – the reactant that is reduced
C + O2  CO2
Oxygen is reduced (0 to -2), so it is the oxidizing agent

13. Oxidizing and reducing agents
Reducing agent – the reactant that is oxidized
3H2 + 2Cr+3  6H+ + 2Cr
Hydrogen is oxidized (0 to +1), so it is the reducing agent
Example: Identify the oxidizing and reducing agents in the following reaction:
2HCl + Zn  ZnCl2 + H2
Zn – reducing agent
H+ – oxidizing agent

14. Redox and electronegativity
C + O2  CO2
Carbon is oxidized because it has lost some electron density to oxygen, which has greater electronegativity.
Oxygen is reduced because it gained some electron density from carbon

15. Balancing redox equations
Charge Balance
Redox is a transfer of electrons, so the number of electrons lost by the reducing agent = number of electrons gained by oxidizing agent
Total charge of reactants must = total charge of products
Cr+6 + Fe+2  Cr+3 + Fe+3
Even though the atoms are balanced, the charge is not.

16. Balancing redox equations
Oxidation number method:
Identify all changes in oxidation number
Cr Fe+2  Cr+3 + Fe+3

17. Balancing redox equations
Use coefficients to make the changes cancel
Cr Fe+2  Cr Fe+3
-3 +1x3 = +3 3 3

18. Balancing redox equations
Check charge balance
Cr+6 + 3Fe+2  Cr+3 + 3Fe+3 
HNO3 + H3AsO3    NO + H3AsO4 + H2O
Use least common multiple – 6
2HNO3 + 3H3AsO3    2NO + 3H3AsO4 + H2O

19. Balancing Redox Equations
Half reactions method
Every redox reaction consists of two half reactions
Fe + Cu+2  Fe+3 + Cu
oxidation
Fe  Fe+3 + 3e-
reduction
Cu+2 + 2e-  Cu
Oxidation and reduction reactions always happen in pairs

20. Balancing Redox Equations
Sum of appropriate numbers of half reactions yields a balanced equation – use coefficients to make
# electrons lost = # electrons gained
2(Fe  Fe+3 + 3e-) +
2(Cu+2 + 2e-  Cu) =
2Fe + 3Cu+2  2Fe+3 + 3Cu

21. Balancing Redox Equations
Atoms and electrons have to balance
If the electrons balance, the charge will also balance (but be sure to check it!)
Cu + HNO3Cu(NO3)2 + NO2 + H2O
Oxidation: Cu  Cu+2 + 2e-
Reduction: NO3- + 1e-  NO2

22. Balancing Redox Equations
Reduction half reaction must be balanced – in acid solution use 2H+ and H2O for each missing oxygen
2H+ + NO3- + 1e-  NO2 + H2O
Number of electrons in oxidation and reduction must be equal
Add half reactions to get balanced equation

23. Balancing Redox Equations
2(2H+ + NO3- + 1e-  NO2 + H2O)
Cu  Cu+2 + 2e-
4H++2NO3-+2e-+CuCu+2+2e-+2NO2+2H2O
Electrons cancel; addition of nitrates to each side (spectators) gives overall equation
4HNO3+CuCu(NO3)2+2NO2+2H2O

24. Balancing Redox Example #2
Zn + VO3-  Zn+2 + VO (in acid solution)
Half reactions:
Oxidation: Zn  Zn+2 + 2e-
VO3- V is +5, VO+2 V is +4
Reduction: VO3- + 1e-  VO+2

25. Balancing Redox Example #2
balance with H+ and H2O
4H+ + VO3- + 1e-  VO+2 + 2H2O
Balanced equation is sum of half reactions
4H++VO3-+ZnVO+2+2H2O+Zn+2

26. Balancing in Base Solution
Use 2OH- and H2O for each missing oxygen
Cr(OH)3 + ClO3-    CrO42- + Cl-
Oxidation
Cr(OH)3  CrO e-+ 3OH-
Hydroxides are added to balance hydrogens.
Balance oxygen (four missing on left) with 2OH-/H2O.

27. Balancing in Base Solution
8OH- + Cr(OH)3  CrO e-+ 3OH- + 4H2O
Cancel hydroxides on both sides.
5OH- + Cr(OH)3  CrO e- + 4H2O
Reduction:
ClO3- + 6e- Cl-
Balance oxygen (three missing on right) with 2OH-/H2O.

28. Balancing Redox in Base Solution
3H2O + ClO3- + 6e-  Cl- + 6OH-
Add equations and eliminate spectators
2[5OH- + Cr(OH)3  CrO e- + 4H2O]
10OH- + 2Cr(OH)3 + 3H2O + ClO3-  2CrO H2O + Cl- + 6OH- 4
4OH- + 2Cr(OH)3 + ClO3-  2CrO H2O + Cl- 5