Class-3 by rey15315


Class III. Electrochemistry Background (II)
                Winter 2009

         Dr. Marc Madou
  Oxidants and reductants
  Battery
  Reference Electrodes
  Standard Reduction Potentials
  Thermodynamic Significance of
  How do Cell Potentials Change if
 We are Not at Standard State?
  Nernst-Equation
  Cyclic voltammetry
  Potentiometric sensors
  Amperometric sensors
Oxidants and Reductants
   oxidant   = oxidizing agent
    – reactant which oxidizes another reactant
      and which is itself reduced

   reductant   = reducing agent
    – reactant which reduces another reactant
      and which is itself oxidized
Oxidants and Reductants
 Identify  the oxidant and reductant in each of the
  following reactions:
  a) Karl Fischer reaction – for quantitation of
  I2 + SO2 + H2O = 2HI + SO3
  b) Hall Heroult process – production of Al:
  2Al2O3 + 3C = 4Al + 3CO2
  c) the Thermite reaction – used to produce
  liquid iron for welding
  2Al + Fe2O3 = 2Fel + Al2O3
Oxidants and Reductants
   Reactions occur pair wise: One cannot have
    oxidation without reduction
   Charge must be conserved: Number of electrons lost
    in oxidation must equal number of electrons gained
    in reduction
   Suppose we add a strip of Zinc metal to a solution of
   Zn - 2e- = Zn2+
        2+ + 2e- = Cu
                                                      Zn strip
   Cu

    Oxidants and Reductants
   It is the relative tendencies of oxidants and reductants to
    gain/lose electrons that determines the extent of a redox
   Strong oxidant + strong reductant  completion
   What if we could separate the oxidant from the reductant?
   We would have set up a constant flow of electrons = current
    = electricity!                                       1.1 V
                                                salt bridge
               Zn strip                 Zn                  Cu

        CuSO4                          ZnSO4           CuSO4
                                         1836 The Daniell
   Electrode
     – anode = electrode at which oxidation occurs
     – cathode = electrode at which reduction occurs
   Salt bridge = completes the electrical circuit
     – allows ion movement but doesn’t allow solutions to
     – salt in glass tube with vycor frits at both ends
   Since electrons flow from one electrode to the other in
    one direction, there is a potential difference between the
   This difference is called
     – The electromotive force (EMF)
     – Cell voltage

   Since all redox reactions occur pair wise, i.e.,
    reduction and oxidation always occur at the same
    time we cannot measure the cell potential for just one
    half cell reaction
    and this means we must establish a RELATIVE scale
    for cell potentials
Problem: True or False
   In the Daniell cell, zinc metal is reduced to zinc(II) at
    the cathode and copper is oxidized to copper(II) at
    the anode
   In the Daniell cell, zinc is the oxidant and copper is
    the reductant
    Reference Electrodes

   Electrodes with a potential
    independent of solution              H2(gas)
   Standard hydrogen electrode
     – 1 M H+(aq)+ 2e- = H2(g) (1 atm)
     – We define E0  0 V for this
        » where 0 stands for standard              HCl
             1 M all solutes                 Pt black
             1 atm all gases

             250C (298 K)
Reference Electrodes
Reference Electrodes

                  2H+(1M) + 2e-  H2(g,1atm
                             Eoredn = 0.0V
Reference Electrodes

                                      a Ag aCl 
                E  E o  0.0592log
                                       a AgCl
                E  E o  0.0592log aCl 
Reference Electrodes

                       0.244 V v. SHE
Reference Electrodes
Reference Electrodes
Standard Reduction Potentials

      Li+ + e- = Li            -3.0 V
      2H2O + 2e- = H2 + 2OH-   -0.83 V
      Zn2+ + 2e- = Zn          -0.76 V
      2H+ + 2e- = H2           0 V (SHE)
      Cu2+ + 2e- = Cu          0.34 V
      MnO4- +8H+ +5e- = Mn2+   1.51 V
Standard Reduction Potentials

 Always write
 the redox
 ractions as
 shown :
Standard Reduction Potentials
 Halfcell reactions are reversible, i.e.,
  depending on the experimental conditions
  any half reaction can be either an anode or a
  cathode reaction
 Changing the stoichiometry does NOT
  change the reduction potential (intensive
 Oxidation potentials can be obtained from
  reduction potentials by changing the sign
  Ecell = Eanode + Ecathode
Standard Reduction Potentials

                            Li+ + e- = Li            -3.0 V
Problem:                    2H2O + 2e- = H2 + 2OH-   -0.83 V
                            Zn2+ + 2e- = Zn          -0.76 V
   Calculate the cell      2H+ + 2e- = H2           0V
    potential for the        (SHE)
    Daniell cell.
                            Cu2+ + 2e- = Cu          0.34 V
                            MnO4- +8H+ +5e- = Mn2+   1.51 V
Standard Reduction Potentials
Standard Reduction Potentials

  Zn --> Zn2+ + 2e-   Cu2+ + 2e- -->Cu
  oxidation           reduction
Standard Reduction Potentials

   Anode reaction appears leftmost while cathode
    reaction appears rightmost
   All redox forms of reagents present should be listed.
    Phase and concentration specified in brackets, e.g.,
    ZnSO4(aq, 1 M)
   A single vertical line (|) is used to indicate a change of
    phase (s to l to g)
   A double vertical line (||) indicates a salt bridge
   A comma should be used to separate 2 components
    in the same phase
Thermodynamic Significance
of Potentials
   We usually operate electrochemical cells at
    constant P and T
   Recall,
      – G = H - T S (change in Gibbs free energy)
      – H = E + (PV)
     So, GT,P=welec = -qE = -(nF)E
      – since q = n F
      – Recall, F is Faraday’s constant 96,485 C/mole
Thermodynamic Significance of
 The  maximum electrical work done by an
  electrochemical cell equals the product of the
  charge flowing and the potential difference
  across which it flows. The work done on the cell
  – W = -E x Q, where E is the Electromotive Force of
    the Cell (EMF), and Q is the charge flowing: Q = n x
    NA x e
  – where n is the number of moles of electrons
    transferred per mole of reaction, NA is Avogadro's
    Number (6.02 x 1023), and e is the charge on an
    electron (-1.6 x 10-19 C).
 Note:   NA x e = F (one Faraday). Thus: W = -
Thermodynamic Significance
of Potentials
   Recall sign of G provides information on
    G negative  spontaneous reaction
    G positive  non-spontaneous reaction
   So, since G = - nFE
     E positive  spontaneous reaction
    E negative  non-spontaneous reaction

       a         b
     A + ne = B
     reac tan t product
     Ox + ne = Red
Thermodynamic Significance
of Potentials
   Since half-cell potentials are measured relative to
    SHE, they reflect spontaneity of redox reactions
    relative to SHE
   More positive potentials  more potent oxidants
    (oxidants want to be reduced)
   More negative potentials  more potent reductants
    (reductants don’t want to be reduced; they
    spontaneously oxidize)
Thermodynamic Significance
of Potentials
   Galvanic
    – Chemical energy  electrical energy
    – Spontaneous
      (so Ecell is positive)

       » Primary (non-rechargeable)
           Le Clanche (dry cell)

       » Secondary (rechargeable)
           Lead storage battery

       »Hydrogen-Oxygen Fuel Cell
Thermodynamic Significance
of Potentials
   Electrolytic
     – Electrical energy  chemical energy
     – Non-spontaneous
       (Ecell is negative)

    – Lead storage battery when recharging
    – Electrolysis of water
Thermodynamic Significance
of Potentials
Thermodynamic Significance
of Potentials
Thermodynamic Significance
of Potentials
Thermodynamic Significance
of Potentials
Thermodynamic Significance
of Potentials
Thermodynamic Significance of
    Arrange the following in order of increasing oxidizing
      – MnO4- in acidic media
      – Sn2+
      – Co3+
     Co3+ + e- = Co2+                       1.82 V
     MnO4- + 4H+ + 3e- = MnO2 + 2H2O        1.70 V
     MnO4- + 8H+ + 5e- = Mn2+ + 4H2O        1.51 V
     Sn2+ + 2e- = Sn                        -0.14 V

     So, Co3+ > MnO4- > Sn2+
Thermodynamic Significance
of Potentials-Problems
   A galvanic cell consists of a Mg electrode in a 1.0 M
    Mg(NO3)2 solution and a Ag electrode in a 1.0 M AgNO3
    solution. Calculate the standard state cell potential and
    diagram the cell.

   Consider the following cell:
    Ag(s)/AgNO3(aq, 1 M)//CuSO4(aq, 1 M)/Cu(s)
    a) what is the anode reaction?
    b) what is the cathode reaction?
    c) what is the net number of electrons involved?
    d) what is the net reaction?
    e) what is the cell potential at standard state?
    f) is the cell galvanic or electrolytic?
    Thermodynamic Significance
    of Potentials -Problems
   Is the following redox reaction spontaneous?
    Mg2+ + 2Ag = Mg + 2Ag+        given:
    Ag+ + e- = Ag                    +0.80 V
    Mg2+ + 2e- = Mg                  -2.37 V
Thermodynamic Significance
of Potentials
   Using a table of standard reduction potentials, any species
    on the left of a given half reaction will react spontaneously
    with any species appearing on the right of any half reaction
    that appears below it when reduction potentials are listed
    from highest and most positive to lowest and most
Thermodynamic Significance
of Potentials -Problems
 What would the cell potential be for the following cell?
  Ag(s)/AgNO3(aq, 1 M)//CuSO4(aq, 0.5 M)/Cu(s)
 This represents a set of non-standard state
  conditions so we need derive an equation relating the
  standard state to the non-standard state or the
  Nernst Equation
  Standard state:
    –   Temperature 250C (K = 273.15 + 0C)
    –   Pressure 1 atm
    –   Concentrations of all solutes 1 M
    –   0 (not) is used to indicate at standard state

    –   Example: E0 = cell potential at standard state
Change if We are Not at
Standard State?
 Forthe reaction:
  aA + bB = cC + dD
 G = G0 + 2.303 RT log Q
  where Q is the reaction quotient:

        Q   a b
                c   d

 Where   c is the activity for product C
Change if We are Not at
Standard State?
 Since  G = - nFE then
  E = E0 - 2.303 (RT/nF) log Q
 At standard state,
  E = E0 - (0.0591 V/n) log Q
  This is called the Nernst equation
 Apply the Nernst Equation to a pH sensor: pH=-
 What is the cell potential for the following
  electrochemical cell? What type of cell is it?
  Ni(s) | Ni2+ (aq, 0.1 M) || Co2+ (aq, 2.5 M) | Co(s)
Nernst Equation

    G  G o  RT ln Q
Nernst Equation
   The Nernst equation underlies the operating
    principle of potentiometric sensing electrodes and
    reference electrodes
   Electrolysis vs. battery is determined by Eo sign
Two-electrode and three-eletrode cells,
potentiostat, galvanostat
    Electrolytic cell (example):
      – Au cathode (inert surface for
          e.g. Ni deposition)
      – Graphite anode (not attacked
          by Cl2)
    Two electrode cells (anode,
     cathode, working and reference or
     counter electrode) e.g. for
     potentiometric measurements
     (voltage measurements) (A)
    Three electrode cells (working,
     reference and counter electrode)
     e.g. for amperometric
     measurements (current
Cyclic voltammetry: activation control

   At equilibrium the exchange
    current density is given by:
                                     (1  )F e                         F e
                            kT                                 kT
        ie  i  k c zF          e       RT          i  k a zF      e     RT
                               h                                    h

 The
      reaction polarization is then
    given by:
              
           i i i

   The measurable current density
    is then given by:
                  (1 )F            F 
      i  ie (e      RT
                               e     RT
   For large enough overpotential:

        a  blog(i)
                                                  (Tafel law)
Cyclic voltammetry: diffusion control
                                            Since C x=0 i l - i
   From activation control to                          
                                                   C 0
                                                            i l we get :
    diffusion control:  C
            dC C     0
                       x        x 0
                                                            nFc
            dX                
                                             i  il (1  e    RT

   Concentration difference leads
    to another overpotential i.e.
    concentration polarization:
                  RT C x=0
          c        ln
                  nF    C 0

   Using Faraday’s law we may
    write also: 0
                  C  Cx0
      i  nFAD0
   At a certain potential C x=0=0
    and then:nFAD C  0
          I 
            l             0
    Cyclic voltammetry and potentiometric
    and amperometric sensors

   Scan the voltage at a given
    speed (e.g. from + 1 V vs
    SCE to -0.1 V vs SCE and
    back at 100 mV/s) and
    register the current
   Potentiometric: the voltage
    between the sensing
    electrode and a reference
    electrode is registered
   Amperometric: the current at
    a fixed voltage in the diffusion
    plateau is registered

Cyclic voltammetry (also polarography)
and potentiometric and amperometric
1.   Calculate the potential of a battery with a Zn bar in a 0.5 M Zn
     2+ solution and Cu bar in a 2 M Cu 2+ solution.

2.   Show in a cyclic voltammogram the transition from kinetic
     control to diffusion control and why does it really happen ?
3.   Derive how the capacitive charging of a metal electrode
     depends on potential sweep rate.
4.   What do you expect will be the influence of miniaturization on a
     potentiometric sensor and on an amperometric sensor?

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