# Chapter 18 Lecture Outline

## 18.2 Balancing Oxidation-Reduction Equations

### Oxidation Numbers

1. Calculating Oxidation Numbers
1. Elements have an oxidation state of 0 in natural state
2. For monoatomic ions, is charge
3. Oxygen is usually 2-, unless peroxide, then 1-
4. Hydrogen is usually 1+, except metal hydrides (NaH) then 1-
5. Sum of oxidation states for a compound is 0
6. Sum of oxidation states for polyatomic ion is charge

2. Oxidation Number examples:
1. Na2SO3
1. Na 1+
2. S +4
3. O -2
2. PF3
1. P +3
2. F -1
3. CrO32-
1. Cr +4
2. O -2
4. Cr2O72-
1. Cr +6
2. O -2
5. (NH4)3PO4
1. N -3
2. H +1
3. P +5
4. O -2

### Balancing Redox Reactions

1. Overview
1. CHARGE MUST BALANCE!!!
2. Oxidation State Method (easy)
3. Half Reaction Method (hard but more useful)

2. Copper metal in silver nitrate solution (, internet ©Saunders 1997)

Cu(s) + Ag1+ (aq) --> Cu2+ (aq) + Ag (s)

1. Balance with Oxidation State Method
2. Balance with Half Reaction Method
1. Oxidation: Cu (s) --> Cu2+ (aq) + 2 e-
2. Reduction Ag1+ + 1 e- --> Ags
3. Balance electrons and Total: Cu(s) + 2 Ag1+ (aq) --> Cu2+ (aq) + Ag (s)

3. Electrolysis of water 2 H2O --> 2 H2 + O2
1. Oxidation State Method
2. Half Reaction Method
1. Oxidation 2 H2O --> O2 + 4 H+ + 4 e-
2. Reduction 4 H+ + 4 e- --> 2 H2
3. Total 2 H2O --> 2 H2 + O2

4. Oxidation of iron 4 Fe + O2 --> 2 Fe2O
1. Oxidation State Method
2. Half Reaction Method
1. Oxidation Fe --> Fe2+ + 2 e-
2. Reduction O2 + 4 e- --> 2 O-2
3. Balance electrons and Total 2 Fe2+ + O2- Fe2O

5. Cr2O7-2 + HSO3- --> Cr3+ + SO42-
1. First try to balance in acid solution (Add H+)
1. 1 : 1
2. 1 : 2
3. Give up
2. Oxidation State method in Acid
1. Find species that gains electrons
2. Find species that looses electrons
3. Balance electron gain and loss
4. Add H+ and H2O to balance
3. Half-Reaction method
1. Find Oxidized Half Reaction
1. Balance reaction, add H+ and H2O as needed
2. add electrons to balance charge
2. Find Reduction Half Reaction
1. Balance reaction, add H+ and H2O as needed
2. add electrons to balance charge
3. Set half reactions to same number of electrons
4. Combine everything and cancel
5. If in base add OH- on both sides to eliminate H+

## 18.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions

1. Definitions
1. Anode/Oxidation
2. Cathode/Reduction
3. Galvanic Cell, Spontaneous, [delta]G < 0
4. Electrolytic Cell, Not-spontaneous, [delta]G > 0

2. Water
1. Combustion 2 H2 + O2 -> 2 H2O
1. Oxidation 2 H2 -> 4 e- + 4 H+
2. Reduction O2 + 4 H+ + 4e- -> 2 H2O

2. Electrolysis 2 H2O -> 2 H2 + O2
1. Oxidation 2 H2O -> O2 + 4 H+ + 4 e-
2. Reduction 4 H+ + 4 e- -> 2 H2

3. Zinc and Copper Zn (s) + Cu2+ <-> Zn2+ + Cu (s)
1. Galvanic Cell Zn (s) + Cu2+ -> Zn2+ + Cu (s)
1. Oxidation Zn (s) -> Zn2+ + 2 e-
2. Reduction Cu2+ + 2 e- -> Cu (s)

2. Electrolytic Cell Zn2+ + Cu (s) -> Zn (s) + Cu2+
1. Oxidation Cu (s) -> Cu2+ + 2 e-
2. Reduction Zn2+ + 2 e- -> Zn (s)

4. Half Reaction and Half Cell for Zn (s) + Cu2+ <-> Zn2+ + Cu (s)
1. 1 Beeker
1. CuSO4
2. Zn electrode

2. 2 Beekers (internet ©Saunders 1997)
1. One with Zn electrode and ZnSO4
2. One with CuSO4 and Cu electrode

3. Cell Notation for galvanic cell
1. Zn|Zn2+ (1 M)||Cu2+ (1 M)|Cu
2. Oxidation, Anode on left
3. Reduction, Cathode on right
4. Electrons flow from left to right
5. Rewrite for electrolysis cell

## 18.4 Standard Electrode Potentials

1. Units and Concepts

 Property Unit Unit Abbreviation Potential (electron pressure) Voltage V Current (electron flow) Amps A Charge (electron amount) Coulomb C

2. Introduce table of standard potentials
3. Calculating E°cell & Standard cell notation
1. Zn | Zn2+ (1 M) || Fe2+ (1 M) | Fe
1. Draw cell
2. Oxidation, Anode
3. Reduction, Cathode
4. Use Reduction potentials: Ecell = Ecathode - Eanode
5. Solutions

2. Work in groups
1. Fe | Fe2+ (1 M) || H1+ (1 M), H2 (1 atm) | Pt
2. Pt | H1+ (1 M), H2 (1 atm) ||Cu2+ (1 M) | Cu
3. Fe | Fe2+ (1 M) || Cu2+ (1 M) | Cu
4. Cu | Cu2+ (1 M) || Fe2+ (1 M) | Fe

## 18.5 Cell Potential, Free Energy, and the Equilibrium Constant

1. 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

2. Discuss Equilibrium for the reaction

Zn | Zn2+ (1 M) || Fe2+ (1 M) | Fe

1. Draw cell and write out half reactions
2. Cell potential
3. Change Ion Concentration (internet ©Saunders 1997)
5. Equilibrium effects e-'s (pressure or concentration) this is voltage (potential).

## 18.6 Cell potential and Concentration

1. Nernst Equation
1. E° = Std Cell potential
2. R = 8.314 J Mol-1 K-1
3. T = Kelvin
4. n = number of electrons in balanced redox
5. F = 96485 coulomb equiv-1 (converts charge to moles)
6. ln = natural log
7. Q = equilibrium expression for balanced redox reaction

2. 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

3. Calculate concentration ratio (Q) based upon cell potential for the cell

Zn|Zn2+ || Fe2+ | Fe

1. Ecell = 0.3538 V
2. Ecell = 0.3530 V
3. Ecell = 0.3400 V
4. Ecell = 0.3550 V
5. Ecell = 0.3600 V

4. Solutions

## 18.7 Batteries: Using Chemistry to Generate Electricity

1. Batteries
1. Dry Cell
2. 2 MnO2 (s) + 2 NH4Cl (s) + Zn (s) --> Mn2O3 + H2O (l) + Zn(NH3)2Cl2 (s)

2. Alkaline Battery
2. Zn (s) + 2 MnO2 (s) <--> ZnO (s) + Mn2O3 (s)

1. Discharge: Pb (s) + PbO2 (s) + 2 H2SO4 (aq) --> 2 PbSO4 (s) + 2 H2O (l)
2. Charging: 2 PbSO4 (s) + 2 H2O (l) --> Pb (s) + PbO2 (s) + 2 H2SO4 (aq)

2. Discharge: Cd (s) + 2 NiO(OH) (s) + 2 H2O (l) --> Cd(OH)2 (s) + 2 Ni(OH)2 (s)
3. 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

1. Electrolysis demonstration and (calculations)