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					    Electrochemistry

Thermodynamics at the electrode
              Learning objectives
   You will be able to:
   Identify main components of an electrochemical cell
   Write shorthand description of electrochemical cell
   Calculate cell voltage using standard reduction potentials
   Apply Nernst equation to determine free energy change
   Apply Nernst equation to determine pH
   Calculate K from electrode potentials
   Calculate amount of material deposited in electrolysis
       Energy in or energy out
 Galvanic (or voltaic) cell relies on
  spontaneous process to generate a potential
  capable of performing work – energy out
 Electrolytic cell performs chemical reactions
  through application of a potential – energy in
              Redox Review
 Oxidation is...
   Loss of electrons
 Reduction is...
   Gain of electrons
 Oxidizing agents oxidize and are reduced
 Reducing agents reduce and are oxidized
 Redox at the heart of the matter
 Zn displaces Cu from CuSO4(aq)
 In direct contact the enthalpy of reaction is
  dispersed as heat, and no useful work is done
 Redox process:
   Zn is the reducing agent
   Cu2+ is the oxidizing agent
                           2
            Zn ( s )  Zn (aq)  2e
                2
            Cu (aq)  2e  Cu ( s)
        Separating the combatants




   Each metal in touch with a solution of its own ions
   External circuit carries electrons transferred during the redox process
   A “salt bridge” containing neutral ions completes the internal circuit.
   With no current flowing, a potential develops – the potential for work
   Unlike the reaction in the beaker, the energy released by the reaction
    in the cell can perform useful work – like lighting a bulb
Labelling the parts
        Odes to a galvanic cell
 Cathode                     Anode
   Where reduction occurs      Where oxidation occurs
   Where electrons are         Where electrons are
    consumed                     generated
   Where positive ions         Where negative ions
    migrate to                   migrate to
   Has positive sign           Has negative sign
    The role of inert electrodes
 Not all cells start with elements as the redox
  agents
 Consider the cell
           Fe( s )  2 Fe 3 (aq)  3Fe 2 (aq)

 Fe can be the anode but Fe3+ cannot be the
  cathode.
 Use the Fe3+ ions in solution as the
  “cathode” with an inert metal such as Pt
Anode     Catho
           de




Oxidati   Reduct
 on        ion
                     Cell notation
   Anode on left, cathode on right
   Electrons flow from left to right
   Oxidation on left, reduction on right
   Single vertical = electrode/electrolyte boundary
   Double vertical = salt bridge

       Anode:                                          Cathode:
      Zn →Zn2+                                         Cu2+ + 2e
        + 2e                                             →Cu
Vertical │denotes different phase
 Fe(s)│Fe2+(aq)║Fe3+(aq),Fe2+(aq)│Pt(s)

 Cu(s)│Cu2+(aq)║Cl2(g)│Cl-(aq)│C(s)
 Connections: cell potential and free
              energy
 The cell in open circuit generates an
  electromotive force (emf) or potential or
  voltage. This is the potential to perform
  work
 Energy is charge moving under applied
  voltage

           1J  1C 1V
    Relating free energy and cell
               potential
 The Faraday:
            F = 96 485 C/mol e

           G  nFE
   Standard conditions (1 M, 1 atm, 25°C)

          G  nFE
  Standard Reduction Potentials
 The total cell potential is the sum of the potentials
  for the two half reactions at each electrode
                   Ecell = Ecath + Ean
 From the cell voltage we cannot determine the
  values of either – we must know one to get the
  other
 Enter the standard hydrogen electrode (SHE)
 All potentials are referenced to the SHE (=0 V)
          Unpacking the SHE
 The SHE consists of a Pt electrode in contact with
  H2(g) at 1 atm in a solution of 1 M H+(aq).
 The voltage of this half-cell is defined to be 0 V
 An experimental cell containing the SHE half-cell
  with other half-cell gives voltages which are the
  standard potentials for those half-cells
                   Ecell = 0 + Ehalf-cell
        Zinc half-cell with SHE
 Cell measures 0.76 V
 Standard potential for Zn(s) = Zn2+(aq) + 2e = 0.76
  V
       Where there is no SHE
 In this cell there is no SHE and the
  measured voltage is 1.10 V
                   2            2
           Zn Zn (aq) Cu (aq) Cu
                  2             2
      Zn ( s)  Cu (aq)  Zn (aq)  Cu ( s)

         Zn ( s )  Zn 2 (aq)  2e, E o  0.76V
           2
       Cu (aq)  2e  Cu ( s), E  0.34V
                                      o
  Standard reduction potentials
 Any half reaction can be written in two ways:
 Oxidation:
              M = M+ + e (+V)
 Reduction:
               M+ + e = M (-V)
 Listed potentials are standard reduction
  potentials
       Applying standard reduction
                potentials
 Consider the reaction
           Zn ( s )  2 Ag  (aq)  Zn 2 (aq)  2 Ag ( s )
 What is the cell potential?
 The half reactions are:
    
  Ag (aq)  e  Ag ( s)     Zn ( s )  Zn 2 (aq)  2e
 E° = 0.80 V – (-0.76 V) = 1.56 V
 NOTE: Although there are 2 moles of Ag
  reduced for each mole of Zn oxidized, we do not
  multiply the potential by 2.
          Extensive v intensive
 Free energy is extensive property so need to
  multiply by no of moles involved

             G  nFE
 But to convert to E we need to divide by no of
  electrons involved
                                  
                   E   G
                     
                                      nF
 E is an intensive property
        The Nernst equation
 Working in nonstandard conditions
                       
          G  G  RT ln Q
                           
            nFE  nFE  RT ln Q
                   
           E  E  RT               ln Q
                               nF
               
         E  E  0.0592 log Q
                       n
    Electrode potentials and pH
 For the cell reaction
                                                
      H 2 ( g )  2 H (aq)  2e
 The Nernst equation
          EH                E   
                                              
                                                0.06V   
                                                         log 
                                                              H 
                                                                  2
                                                                    
               2 2 H
                        
                                 H 2 2 H 
                                                  n          pH 2 
                                                                   


          EH
                2 2 H
                               
                                   0.06V
                                     n
                                         log H              2




 Half-cell potential is proportional to pH
The pH meter is an electrochemical cell

 Overall cell potential is proportional to pH
           Ecell  0.06V  pH   Eref

                     Ecell  Eref
            pH 
                        0.06V
 In practice, a hydrogen electrode is
  impractical
   Calomel reference electrodes
 The potential of the calomel electrode is known vs
  the SHE. This is used as the reference electrode
  in the measurement of pH
                                          
    Hg 2Cl2 ( s)  2e  2 Hg (l )  2Cl
 The other electrode in a pH probe is a glass
  electrode which has a Ag wire coated with AgCl
  dipped in HCl(aq). A thin membrane separates
  the HCl from the test solution
    Cell potentials and equilibrium
                G  nFE
   Lest we forget…   
                    G   RT ln K

                    
 So then
            nFE   RT ln K
 and E   RT             2.303RT
                    ln K          log10 K
                 nF           nF
Cell potential a convenient way to
            measure K
  Many pathways to one ending
 Measurement of K from different
  experiments
   Concentration data
                       c
                         C  D
                            d


                         A B
                             a   b



   Thermochemical data      G    RT ln K

                                     
   Electrochemical data   nFE   RT ln K
                    Batteries
 The most important application of galvanic
  cells
 Several factors influence the choice of
  materials
     Voltage
     Weight
     Capacity
     Current density
     Rechargeability
          Running in reverse
 Recharging a battery requires to run the
  process in reverse by applying a voltage
 In principle any reaction can be reversed
 In practice it will depend upon many factors
 Reversibility depends on kinetics and not
  thermodynamics
 Cell reactions that involve minimal structural
  rearrangement will be the easiest to reverse
           Lithium batteries
 Lightweight (Molar mass Li = 6.94 g)
 High voltage
 Reversible process
     Fuel cells – a battery with a
              difference
 Reactants are not contained within a sealed
  container but are supplied from outside
  sources
       anode : 2 H 2 ( g )  4OH  (aq)  4 H 2O(l )  4e

      cathode : O2 ( g )  2 H 2O(l )  4e  4OH  (aq)

      overall : 2 H 2 ( g )  O2 ( g )  2 H 2O(l )
     Store up not treasures on earth
         where moth and rust…
 An electrochemical mechanism for corrosion of iron. The metal and a
  surface water droplet constitute a tiny galvanic cell in which iron is
  oxidized to Fe2+ in a region of the surface (anode region) remote from
  atmospheric O2, and O2 is reduced near the edge of the droplet at
  another region of the surface (cathode region). Electrons flow from
  anode to cathode through the metal, while ions flow through the water
  droplet. Dissolved O2 oxidizes Fe2+ further to Fe3+ before it is deposited
  as rust (Fe2O3·H2O).
               Mechanisms
 Why does salt enhance rusting?
   Improves conductivity of electrolyte


 Standard reduction potentials indicate which
  metals will “rust”
 Aluminium should corrode readily. It
  doesn’t. Is thermodynamics wrong?
   No, the Al2O3 provides an impenetrable barrier
 No greater gift than to give up your
         life for your friend
 A layer of zinc protects iron from oxidation, even when the
  zinc layer becomes scratched. The zinc (anode), iron
  (cathode), and water droplet (electrolyte) constitute a tiny
  galvanic cell. Oxygen is reduced at the cathode, and zinc is
  oxidized at the anode, thus protecting the iron from
  oxidation.
                         Electrolysis
 Electrolysis of a molten salt using inert electrodes
 Signs of electrodes:
    In electrolysis, anode is positive because electrons are removed
     from it by the battery
    In a galvanic cell, the anode is negative because is supplies
     electrons to the external circuit



  Anode : 2Cl  (l )  Cl2 ( g )  2e

Cathode : 2 Na  (l )  2e  2 Na(l )

                                   
 Overall : 2 Na (l )  2Cl (l )  2 Na (l )  Cl2 ( g )
  Electrolysis in aqueous solutions – a
            choice of process
  There are (potentially)
   competing processes
   in the electrolysis of an
   aqueous solution
     Cathode Cathode : 2 Na  (l )  2e  2 Na (l )... E   2.71V

Cathode : 2 H 2O(l )  2e  H 2 ( g )  2OH  (aq)... E   0.83V
     Anode
              Anode : 2Cl  (l )  Cl2 ( g )  2e... E   1.36V
             Anode : 2 H 2O(l )  O2 ( g )  4 H   4e... E   1.23V
     Thermodynamics or kinetics?
  On the basis of thermodynamics we choose
   the processes which are favoured
   energetically
         Anode : 2 H 2O(l )  O2 ( g )  4 H   4e... E   1.23V

Cathode : 2 H 2O(l )  2e  H 2 ( g )  2OH  (aq)... E   0.83V

  But…chlorine is evolved at the anode
     The role of overpotentials
 Thermodynamic quantities prevail only at
  equilibrium – no current flowing
 When current flows, kinetic considerations
  come into play
 Overpotential represents the additional
  voltage that must be applied to drive the
  process
 In the NaCl(aq) solution the overpotential for
  evolution of oxygen is greater than that for
  chlorine, and so chlorine is evolved
  preferentially
 Overpotential will depend on the electrolyte
  and electrode. By suitable choices,
  overpotentials can be minimized but are never
  eliminated
 The limiting process in electrolysis is usually
  diffusion of the ions in the electrolyte (but not
  always)
 Driving the cell at the least current will give
  rise to the smallest overpotential
                Electrolysis of water
  In aqueous solutions of
   most salts or acids or
   bases the products will
   be O2 and H2
Cathode : 2 H 2O(l )  2e  H 2 ( g )  2OH  (aq)... E   0.83V
                                        
Anode : 2 H 2O(l )  O2 ( g )  4 H  4e... E   1.23V
Quantitative aspects of electrolysis
 Quantitative analysis
  uses the current
  flowing as a measure
  of the amount of
  material
 Charge = current x
  time
 Moles =
  charge/Faraday

				
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posted:8/12/2012
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