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