Reduction - Oxidation Chapters 20 + 21

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

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

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

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

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

Redox Reactions LEO goes GER Lose Electrons Oxidation Gain Electrons Reduction

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

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)

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

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

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

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

Half Reactions Zn + CuSO4  Cu + ZnSO4 Zn + Cu+2  Cu + Zn+2 Net Ionic Zn  Zn+2 + 2e- Oxidation Cu2+ + 2e-  Cu Reduction

Oxidation Loss of Electrons Examples: Zn  Zn+2 + 2e- 2Cl-  Cl2 + 2e-

Reduction Gain of electrons Examples: Ag+ + e-  Ag Cl2 + 2e-  2Cl-

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

Balancing Example Ti  Ti+4 + 4e- 2(Cu+2 + 2e-  Cu) Ti + 2Cu+2  Ti+4 + 2Cu Ti + 2CuCl2  TiCl4 + 2Cu

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

Spontaneous Reactions Examples: Zn + CuSO4  Cu + ZnSO4 CaSO4 + Mg  Ca + MgSO4 Zn + 2HCl  H2 + ZnCl2 F2 + 2NaI  I2 + 2NaF YES NO YES YES

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

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

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

Voltaic Cell

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

Voltaic Cell

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

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

Electrolytic Cell

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