Chapter 18 Lecture Outline
18.1 Pulling the Plug on the Power Grid
18.2 Balancing Oxidation-Reduction Equations
Oxidation Numbers
- Calculating Oxidation Numbers
- Elements have an oxidation state of 0 in natural state
- For monoatomic ions, is charge
- Oxygen is usually 2-, unless peroxide, then 1-
- Hydrogen is usually 1+, except metal hydrides (NaH) then 1-
- Sum of oxidation states for a compound is 0
- Sum of oxidation states for polyatomic ion is charge
- Oxidation Number examples:
- Na2SO3
- Na 1+
- S +4
- O -2
- PF3
- P +3
- F -1
- CrO32-
- Cr +4
- O -2
- Cr2O72-
- Cr +6
- O -2
- (NH4)3PO4
- N -3
- H +1
- P +5
- O -2
- Overview
- CHARGE MUST BALANCE!!!
- Oxidation State Method (easy)
- Half Reaction Method (hard but more useful)
- Copper metal in silver nitrate solution (, internet ©Saunders 1997)
Cu(s) + Ag1+ (aq) --> Cu2+ (aq) + Ag (s)
- Balance with Oxidation State Method
- Balance with Half Reaction Method
- Oxidation: Cu (s) --> Cu2+ (aq) + 2 e-
- Reduction Ag1+ + 1 e- --> Ags
- Balance electrons and Total: Cu(s) + 2 Ag1+ (aq) --> Cu2+ (aq) + Ag (s)
- Electrolysis of water 2 H2O --> 2 H2 + O2
- Oxidation State Method
- Half Reaction Method
- Oxidation 2 H2O --> O2 + 4 H+ + 4 e-
- Reduction 4 H+ + 4 e- --> 2 H2
- Total 2 H2O --> 2 H2 + O2
- Oxidation of iron 4 Fe + O2 --> 2 Fe2O
- Oxidation State Method
- Half Reaction Method
- Oxidation Fe --> Fe2+ + 2 e-
- Reduction O2 + 4 e- --> 2 O-2
- Balance electrons and Total 2 Fe2+ + O2- Fe2O
- Cr2O7-2 + HSO3- --> Cr3+ + SO42-
- First try to balance in acid solution (Add H+)
- 1 : 1
- 1 : 2
- Give up
- Oxidation State method in Acid
- Find species that gains electrons
- Find species that looses electrons
- Balance electron gain and loss
- Add H+ and H2O to balance
- Half-Reaction method
- Find Oxidized Half Reaction
- Balance reaction, add H+ and H2O as needed
- add electrons to balance charge
- Find Reduction Half Reaction
- Balance reaction, add H+ and H2O as needed
- add electrons to balance charge
- Set half reactions to same number of electrons
- Combine everything and cancel
- If in base add OH- on both sides to eliminate H+
18.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions
- Definitions
- Anode/Oxidation
- Cathode/Reduction
- Galvanic Cell, Spontaneous, [delta]G < 0
- Electrolytic Cell, Not-spontaneous, [delta]G > 0
- Water
- Combustion 2 H2 + O2 -> 2 H2O
- Oxidation 2 H2 -> 4 e- + 4 H+
- Reduction O2 + 4 H+ + 4e- -> 2 H2O
- Electrolysis 2 H2O -> 2 H2 + O2
- Oxidation 2 H2O -> O2 + 4 H+ + 4 e-
- Reduction 4 H+ + 4 e- -> 2 H2
- Zinc and Copper Zn (s) + Cu2+ <-> Zn2+ + Cu (s)
- Galvanic Cell Zn (s) + Cu2+ -> Zn2+ + Cu (s)
- Oxidation Zn (s) -> Zn2+ + 2 e-
- Reduction Cu2+ + 2 e- -> Cu (s)
- Electrolytic Cell Zn2+ + Cu (s) -> Zn (s) + Cu2+
- Oxidation Cu (s) -> Cu2+ + 2 e-
- Reduction Zn2+ + 2 e- -> Zn (s)
- Half Reaction and Half Cell for Zn (s) + Cu2+ <-> Zn2+ + Cu (s)
- 1 Beeker
- CuSO4
- Zn electrode
- 2 Beekers (internet ©Saunders 1997)
- One with Zn electrode and ZnSO4
- One with CuSO4 and Cu electrode
- Add Wire
- Add Salt Bridge (KCl)
- Cell Notation for galvanic cell
- Zn|Zn2+ (1 M)||Cu2+ (1 M)|Cu
- Oxidation, Anode on left
- Reduction, Cathode on right
- Electrons flow from left to right
- Rewrite for electrolysis cell
18.4 Standard Electrode Potentials
- Units and Concepts
Property |
Unit |
Unit Abbreviation |
Potential (electron pressure) |
Voltage |
V |
Current (electron flow) |
Amps |
A |
Charge (electron amount) |
Coulomb |
C |
- Introduce table of standard potentials
- Calculating E°cell & Standard cell notation
- Zn | Zn2+ (1 M) || Fe2+ (1 M) | Fe
- Draw cell
- Oxidation, Anode
- Reduction, Cathode
- Use Reduction potentials: Ecell = Ecathode - Eanode
- Solutions
- Work in groups
- Fe | Fe2+ (1 M) || H1+ (1 M), H2 (1 atm) | Pt
- Pt | H1+ (1 M), H2 (1 atm) ||Cu2+ (1 M) | Cu
- Fe | Fe2+ (1 M) || Cu2+ (1 M) | Cu
- Cu | Cu2+ (1 M) || Fe2+ (1 M) | Fe
18.5 Cell Potential, Free Energy, and the Equilibrium Constant
- Relate Cell potential to free energy
[delta]G = -nFE
Free Energy |
Cell Potential |
Reaction |
[delta]G > 0 |
E < 0 |
Not Spontaneous |
[delta]G = 0 |
E = 0 |
Equlibrium |
[delta]G < 0 |
E > 0 |
Spontaneous |
- Discuss Equilibrium for the reaction
Zn | Zn2+ (1 M) || Fe2+ (1 M) | Fe
- Draw cell and write out half reactions
- Cell potential
- Change Ion Concentration (internet ©Saunders 1997)
- Add more metal
- Equilibrium effects e-'s (pressure or concentration) this is voltage (potential).
18.6 Cell potential and Concentration
- Nernst Equation
- E° = Std Cell potential
- R = 8.314 J Mol-1 K-1
- T = Kelvin
- n = number of electrons in balanced redox
- F = 96485 coulomb equiv-1 (converts charge to moles)
- ln = natural log
- Q = equilibrium expression for balanced redox reaction
- In Groups
[Zn2+] |
[Fe2+] |
Ecell (Volt) |
1.00 |
1.00 |
0.3538 |
1.00 |
2.00 |
0.3627 |
2.00 |
1.00 |
0.3449 |
0.100 |
0.100 |
0.3538 |
1.00 |
0.010 |
0.2946 |
0.010 |
1.00 |
0.41296 |
- Calculate concentration ratio (Q) based upon cell potential for the cell
Zn|Zn2+ || Fe2+ | Fe
- Ecell = 0.3538 V
- Ecell = 0.3530 V
- Ecell = 0.3400 V
- Ecell = 0.3550 V
- Ecell = 0.3600 V
- Solutions
18.7 Batteries: Using Chemistry to Generate Electricity
- Batteries
- Dry Cell
- (
internet, Saunders ©1997)
- 2 MnO2 (s) + 2 NH4Cl (s) + Zn (s) --> Mn2O3 + H2O (l) + Zn(NH3)2Cl2 (s)
- Alkaline Battery
- (
internet, Saunders ©1997)
- Zn (s) + 2 MnO2 (s) <--> ZnO (s) + Mn2O3 (s)
- Lead Acid Battery (
internet, Saunders ©1997)
- Discharge: Pb (s) + PbO2 (s) + 2 H2SO4 (aq) --> 2 PbSO4 (s) + 2 H2O (l)
- Charging: 2 PbSO4 (s) + 2 H2O (l) --> Pb (s) + PbO2 (s) + 2 H2SO4 (aq)
- Ni-Cad Battery
- ( internet, Saunders ©1997)
- Discharge: Cd (s) + 2 NiO(OH) (s) + 2 H2O (l) --> Cd(OH)2 (s) + 2
Ni(OH)2 (s)
- Charging: Cd(OH)2 (s) + 2 Ni(OH)2 (s) --> Cd (s) + 2 NiO(OH) (s) + 2
H2O (l)
18.8 Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity
- Electrolysis demonstration and (calculations)
18.9 Corrosion: Undesirable Redox Reactions
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Widener University
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Last Updated Friday, May 25, 2001 1:59:43 PM