Electrochemistry Lecture

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					    Harnessing the Power
      of Voltaic Cells

      Batteries and Corrosion

                  Commercial Voltaic Cells
    Voltaic Cells are convenient energy sources
    Batteries is a self-contained group of voltaic cells arranged in series.
      Advantage: Portable
      Disadvantage: Very Expensive (.80€ / Kwatt-h)
                        Need cells in series to provide power

     The Processes occurring during the
     discharge and recharge of a lead-acid
     battery. When the lead-acid battery
     is discharging (top) it behaves like a
     voltaic cell: the anode is negative
     (electrode-1) and the cathode is
     positive (electrode-2). When it is
     recharging (bottom), it behaves like
     an electrolytic cell; the anode is
     positive (electrode-2) and the
     cathode is negative (electrode-1).

                   Dry Cell or LeClanche Cell
    Dry Cells
    Invented in the 1860’s the common dry cell or LeClanche cell, has become a familiar household
    item. An active zinc anode in the form of a can house a mixture of MnO 2 and an acidic electrolytic
    paste, consisting of NH4Cl, ZnCl2, H2O and starch powdered graphite improves conductivity. The
    inactive cathode is a graphite rod.

                             Anode (oxidation)
                             Zn(s) g Zn2+(aq) = 2e-
                             Cathode (reduction). The cathodic half-reaction is complex and even today, is still
                             being studied. MnO2(s) is reduced to Mn2O3(s) through a series of steps that may
                             involve the presence of Mn2+ and an acid-base reaction between NH4+ and OH- :

                             2MnO2 (s) + 2NH4+(aq) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O (l)
                             The ammonia, some of which may be gaseous, forms a complex ion with Zn2+,
                             which crystallize in contact Cl- ion:

                             Zn2+(aq) + 2NH3 (aq) + 2Cl-(aq) g Zn(NH3)2Cl2(s)

                             Overall Cell reaction:
                             2MnO2 (s) + 2NH4Cl(aq) + Zn(s) g Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s)   Ecell = 1.5 V

                             Uses: common household items, such as portable radios, toys, flashlights,
                             Advantage; Inexpensive, safe, available in many sizes
                             Disadvantages: At high current drain, NH3(g) builds up causing drop in voltage,
                             short shelf life because zinc anode reacts with the acidic NH4+ ions.

               Dry Cell or LeClanche Cell
    Invented by George Leclanche, a French Chemist.

    Acid version:        Zinc inner case that acts as the anode and a carbon
    rod in contact with a moist paste of solid MnO2 , solid NH4Cl, and
    carbon that acts as the cathode. As battery wear down, Conc. of Zn+2
    and NH3 (aq) increases thereby decreasing the voltage.
    Half reactions:      E°Cell = 1.5 V
    Anode:             Zn(s) g Zn+2(aq) + 2e-
    Cathode:    2NH4+(aq) + MnO2(s) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O(l)

       Inexpensive, safe, many sizes

       High current drain, NH3(g) build
       up, short shelf life

                             Alkaline Battery
    Alkaline Battery
    The alkaline battery is an improved dry cell. The half-reactions are similar, but the
    electrolyte is a basic KOH paste, which eliminates the buildup of gases and maintains the
    Zn electrode.
                               Anode (oxidation)
                               Zn(s) + 2OH- (aq) g ZnO(s) + H2O (l) + 2e-
                               Cathode (reduction).
                               2MnO2 (s) + 2H2O (l) + 2e- g Mn(OH)2(s) + 2OH-(aq)

                               Overall Cell reaction:
                               2MnO2 (s) + H2O (l) + Zn(s) g ZnO(s) + Mn(OH)2(s)          Ecell
                               = 1.5 V

                               Uses: Same as for dry cell.
                               Advantages: No voltage drop and longer shell life than
                               dry cell because of alkaline electrolyte; sale ,amu sizes.
                               Disadvantages; More expensive than common dry cell.

                         Alkaline Battery
    Leclanche Battery: Alkaline Version
    In alkaline version; solid NH4Cl is replaced with KOH or NaOH. This
    makes cell last longer mainly because the zinc anode corrodes less
    rapidly under basic conditions versus acidic conditions.
    Half reactions: E°Cell = 1.5 V
    Anode:          Zn(s) + 2OH-(aq) g ZnO(s) + H2O(l) + 2e-
    Cathode:        MnO2 (s) + H2O(l) + 2e- g MnO3 (s) + 2OH-(aq)
    Nernst equation: E = E° - [(0.592/n)log Q], Q is constant !!

     No voltage drop, longer shelf life.

     More expensive

                  Mercury Button Battery
    Mercury and Silver batteries are similar.
    Like the alkaline dry cell, both of these batteries use zinc in a basic
    medium as the anode. The solid reactants are each compressed with
    KOH, and moist paper acts as a salt bridge.
    Half reactions: E°Cell = 1.6 V
    Anode:        Zn(s) + 2OH-(aq) g ZnO(s) + H2O(l) + 2e-
    Cathode (Hg): HgO (s) + 2H2O(l) + 2e- g Hg(s) + 2OH-(aq)
    Cathode (Ag): Ag2O (s) + H2O(l) + 2e- g 2Ag(s) + 2OH-(aq)

      Small, large potential,
      silver is nontoxic.

      Mercury is toxic, silver is

    Lead Storage Battery
    Lead-Acid Battery. A typical 12-V lead-acid battery has six cells
    connected in series, each of which delivers about 2 V. Each cell
    contains two lead grids packed with the electrode material: the
    anode is spongy Pb, and the cathode is powered PbO2. The grids
    are immersed in an electrolyte solution of 4.5 M H2SO4. Fiberglass
    sheets between the grids prevents shorting by accidental physical
    contact. When the cell discharges, it generates electrical energy as a
    voltaic cell.
     Half reactions: E°Cell = 2.0 V
     Anode: Pb(s) + SO42- g PbSO4 (s) +2 e-                              E° = 0.356
     Cathode (Hg): PbO2 (s) + SO42- + 4H+ + 2e- g
                                              PbSO4 (s) + 2 H2O          E° = 1.685V
     Net: PbO2   (s) +   Pb(s) + 2H2SO4 g PbSO4 (s) + 2 H2O                  E°Cell = 2.0 V

     Note hat both half-reaction produce Pb2+ ion, one through
     oxidation of Pb, the other through reduction of PbO2. At both
     electrodes, the Pb2+ react with SO42- to form PbSO4(s)

                  Nickel-Cadmium Battery
    Battery for the Technological Age
    Rechargeable, lightweight “ni-cad” are used for variety of cordless appliances.
    Main advantage is that the oxidizing and reducing agent can be regenerated
    easily when recharged. These produce constant potential.
    Half reactions: E°Cell = 1.4 V
    Anode:            Cd(s) + 2OH-(aq) g Cd(OH)2 (s) + 2e-
    Cathode:          2Ni(OH) (s) + 2H2O(l) + 2e- g Ni(OH)2 (s) + 2 OH-(aq)

     Fuel Cells

                         Fuel Cells; Batteries
     Fuel Cell also an electrochemical device for converting
     chemical energy into electricity.
     In contrast to storage battery, fuel cell does not need to involve a
     reversible reaction since the reactant are supplied to the cell as needed
     from an external source. This technology has been used in the Gemini,
     Apollo and Space Shuttle program.
     Half reactions: E°Cell = 0.9 V
     Anode:           2H2 (g) + 4OH-(aq) g 4H2O(l) + 4e-
     Cathode:         O2 (g) + 2H2O(l) + 4e- g   4OH-(aq)
     Clean, portable and product is water.
     Efficient (75%) contrast to 20-25% car,
     35-40% from coal electrical plant

     Cannot store electrical energy, needs
     continuous flow of reactant, Electrodes
     are short lived and expensive.
     Not all spontaneous redox reaction are beneficial.
     Natural redox process that oxidizes metal to their oxides and
     sulfides runs billions of dollars annually. Rust for example is
     not the direct product from reaction between iron and oxygen
     but arises through a complex electrochemical process.

      Rust: Fe2O3 • X H2O
      Anode: Fe(s) g Fe+2 + 2e-       E° = 0.44 V
      Cathode:        O2 (g) + 4H+ + 4e- g 2H2O (l)   E° = 1.23 V
      Net: Fe+2 will further oxidized to Fe2O3 • X H2O

                      Conditions for Corrosion
     Conditions for Iron Oxidation:
     Iron will oxidize in acidic medium
        SO2 g H2SO4 g H+ + HSO4+
     Anions improve conductivity for oxidation.
        Cl- from seawater or NaCl (snow melting) enhances rusting

     Conditions for Prevention:
     Iron will not rust in dry air; moisture must be present
     Iron will not rust in air-free water; oxygen must be present
     Iron rusts most rapidly in ionic solution and low pH (high H+)
     The loss of iron and deposit of rust occur at different placm on object
     Iron rust faster in contact with a less active metal (Cu)
     Iron rust slower in contact with a more active metal (Zn)

                        Iron Corrosion; Chemistry
     Most common and economically destructive
     form of corrosion is the rusting of iron. Rust
     is not a direct product of the reaction
     between iron and oxygen but arises through
     complex electrochemical process. The
     features of a voltaic cell can help explain this

     Iron will not rust in dry air; moisture must be present.
     Iron will not rust in air-free water; oxygen must be present
     Iron rusts most rapidly in ionic solutions and at low pH (High H+)

     The loss of iron and the
     depositing of rust often occur at
     different places on the same
     Iron rust faster in contact with a
     less active metal (such as Cu)
     and more slowly in contact with
     a more active metal (such as Zn).

     Corrosion Prevention