1. Reduction - Oxidation
Chapters

2. Oxidation Numbers (States)
Positive, negative or neutral values assigned to an atom to keep track of the number of electrons lost or gained.
Charge

3. Oxidation Number Rules
Elements alone (not in a compound) = 0
Example: Cu, N2
Monatomic ion (single atom) = charge
Example: Na+, Cl-, Mg+2, O-2

4. Oxidation Number Rules
Compound sum of all atoms = 0
Example: H2O H + H + O = 0
Polyatomic ion
sum of all atoms = charge
Example: NO3- N + O + O + O = -1

5. Common Oxidation Numbers
Group 1 +1
Group 2  +2
Group 13  +3
Group 15  -3
Group 16  -2
Group 17  -1
Some exceptions to each above

7. Redox Reactions
Reduction – Oxidation, or redox, involves the transfer of electrons
Reduction – gain of electrons
Oxidation – loss of electrons

8. Redox Reactions LEO goes GER Lose Electrons Oxidation
Gain Electrons Reduction

9. Redox Reaction Mg + Cl2  MgCl2 Mg - lost electrons (oxidation)+2 -1
Mg + Cl2  MgCl2
Mg - lost electrons (oxidation)
Cl – gained electrons (reduction)

10. Redox Reaction 2Al + 3Ni(NO3)2  3Ni + 2Al(NO3)3+2 +5 -2 +3 +5 -2
2Al + 3Ni(NO3)2  3Ni + 2Al(NO3)3
Al - lost electrons (oxidation)
Ni – gained electrons (reduction)

12. Redox Reaction Zn + CuSO4  Cu + ZnSO4
One element loses electrons (oxidation)
One element gains electrons (reduction)
All other ions are spectators

13. Net Ionic Equation
Shows only the ions involved in the redox reaction, not spectator ions
Still shows conservation of mass and charge
Zn + CuSO4  Cu + ZnSO4
Zn + Cu+2  Cu + Zn+2

14. Net Ionic Example
Zn + 2HCl  H2 + ZnCl2
Zn + 2H+  H2 + Zn2+

16. Half Reactions
Only shows one element and how many electrons are gained or lost
Must maintain conservation of mass and charge

17. Half Reactions Zn + CuSO4  Cu + ZnSO4 Zn + Cu+2  Cu + Zn+2 Net Ionic
Zn  Zn e- Oxidation
Cu e-  Cu Reduction

18. Oxidation
Loss of Electrons
Examples:
Zn  Zn e-
2Cl-  Cl e-

19. Reduction
Gain of electrons
Examples:
Ag+ + e-  Ag
Cl e-  2Cl-

20. Balancing Reactions
The number of electrons lost must equal the number of electrons gained
Example:
2Na + ZnCl2  Zn + 2NaCl
Zn e-  Zn
2(Na Na + + e- )

21. Balancing Example Ti  Ti+4 + 4e- 2(Cu+2 + 2e-  Cu)
Ti + 2Cu+2  Ti Cu
Ti + 2CuCl2  TiCl Cu

23. Spontaneous Reactions
More active element does not want to be alone
Table J
Metal being oxidized must be ABOVE metal being reduced for spontaneous reactions to occur
Reversed for Nonmetals

24. Spontaneous Reactions
Examples:
Zn + CuSO4  Cu + ZnSO4
CaSO4 + Mg  Ca + MgSO4
Zn + 2HCl  H2 + ZnCl2
F NaI  I NaF
YES NO
YES
YES

26. Electrochemical Cells
any device that converts chemical energy into electrical energy or electrical energy into chemical energy
Two types
Voltaic (Chemical)
Electrolytic

27. Electrochemical Cells
Electrode – metal conductor in an electrical circuit that carries electrons to or from another substance
Cathode – electrode where reduction takes place
Anode – electrode where oxidation takes place

28. Voltaic Cell
Flow of electrons is spontaneous based on electronegativity and ionization energy
Chemical energy is converted to electrical energy
Examples: Batteries

29. Voltaic Cell

32. Electrochemical Cell Components
Salt Bridge
Allows for the passage of ions, not electrons
Switch
Device that opens(turns off) and closes(turns on) circuit

33. Voltaic Cell

34. Electrolysis
Process in which electrical energy is converted to chemical energy
Example:
2H2O  2H2 + O2

35. Electrolytic Cells Electrons are pushed by an outside power source
Electrical energy is converted to chemical energy
Examples: Electroplating, Electropolishing

36. Electrolytic Cell

37. Voltaic or Electrolytic?
Zn + NiCl2  Ni + ZnCl2
Cu + ZnSO4  Zn + CuSO4
2H2O  2H2 + O2
2NaCl  2Na + Cl2
Voltaic
Electrolytic
Voltaic
Electrolytic
Electrolytic