Chapter 20 Lecture Outline

Introduction to Electrochemistry

  1. REDOX Chemsitry. Introduction, to loss of electrons
    1. Oxidation; loss of electrons, increase charge, ie: Fe -> Fe2+
    2. Reduction; gain of electrons, reduce charge, ie O2 -> 2 FeO

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 (CD-ROM d:, 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+

Electrochemical 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 (CD-ROM d:, internet ©Saunders 1997)
      1. One with Zn electrode and ZnSO4
      2. One with CuSO4 and Cu electrode
      3. Add Wire
      4. Add Salt Bridge (KCl)

    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

Electrochemical Reactions and Standard Cell Potential

  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

Cell Potential and Nernst Equation

  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 (CD-ROM d:, internet ©Saunders 1997)
    4. Add more metal
    5. Equilibrium effects e-'s (pressure or concentration) this is voltage (potential).

  3. 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

  4. 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

  5. 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

  6. Solutions

Applications of Electrochemistry:

  1. Electrolysis demonstration and calculations (Mathcad, Acrobat)

  2. Batteries
    1. Dry Cell
      1. ( CD-ROM d:, internet, Saunders ©1997)
      2. 2 MnO2 (s) + 2 NH4Cl (s) + Zn (s) --> Mn2O3 + H2O (l) + Zn(NH3)2Cl2 (s)

    2. Alkaline Battery
      1. ( CD-ROM d:, internet, Saunders ©1997)
      2. Zn (s) + 2 MnO2 (s) <--> ZnO (s) + Mn2O3 (s)

    3. Lead Acid Battery ( CD-ROM d:, internet, Saunders ©1997)
      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)

    4. Ni-Cad Battery
      1. ( CD-ROM d:, internet, Saunders ©1997)
      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)

  3. Determining concentration, the pH electrode.
    1. Ag(s) | AgCl(s) | Cl1-(aq) || H1+(aq, outside) | H1+(aq, inside) || Cl1-(aq) | AgCl(s) | Ag(s)

    2. Show Cell (Overheads)

    3. H1+(aq, outside)|H1+(aq, inside)
      1. H+Gl-(s) <--> H+ (aq) + Gl- (s) inside of electrode (constant)
      2. H+Gl-(s) <--> H+ (aq) + Gl- (s) outside of electrode

    4. Nernst expression for pH electrode

    5. Demonstrate pH electrode in buffers and read voltage.
    6. Calculations

  4. Lithium Ion battery
    1. Figure of Cell

    2. Table of reduction potentials

    3. Lithium Metal Battary
      1. Anode (Ox) Li+ + e- <--> Li -3.045
      2. Cathode (red) Li+ + e- + Lix-1CoO2 <--> LixCO2 +1.0

    4. LiCO2/carbon cell
      1. Anode (Ox) Li+ + e- + Lix-1/graphite <--> Lix/graphite -3.0
        1. Cathode (red) Li+ + e- + Li1-xCoO2 <--> LixCO2 +0.5

      2. Advantages

  5. Electrochemistry experiment with EG&G 263
    1. Reduction of Hg to electrode and strip

    2. Reduction of Hg and Pb to electrode and strip

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