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Harnessing the Power of Voltaic Cells Batteries and Corrosion

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)

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Harnessing the Power of Voltaic Cells Batteries and Corrosion

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  1. Harnessing the Power of Voltaic CellsBatteries and Corrosion

  2. 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).

  3. 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 MnO2 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.

  4. 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) Advantage: Inexpensive, safe, many sizes Disadvantage: High current drain, NH3(g) build up, short shelf life

  5. 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.

  6. 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 !! Advantage: No voltage drop, longer shelf life. Disadvantage: More expensive

  7. 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) Advantage: Small, large potential, silver is nontoxic. Disadvantage: Mercury is toxic, silver is expensive.

  8. 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- PbSO4 (s) +2 e- E° = 0.356 Cathode (Hg): PbO2 (s) + SO42- + 4H+ + 2e- PbSO4 (s) + 2 H2O E° = 1.685V Net: PbO2 (s) + Pb(s) + 2H2SO4  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)

  9. 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) Cd(OH)2 (s) + 2e- Cathode: 2Ni(OH) (s) + 2H2O(l) + 2e- Ni(OH)2 (s) + 2 OH-(aq)

  10. Fuel Cells

  11. 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) 4H2O(l) + 4e- Cathode: O2 (g) + 2H2O(l) + 4e- 4OH-(aq) Advantage: Clean, portable and product is water. Efficient (75%) contrast to 20-25% car, 35-40% from coal electrical plant Disadvantage: Cannot store electrical energy, needs continuous flow of reactant, Electrodes are short lived and expensive.

  12. Corrosion 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

  13. Conditions for Corrosion Conditions for Iron Oxidation: Iron will oxidize in acidic medium SO2g H2SO4g 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)

  14. 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 process. 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 object. 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).

  15. Corrosion Prevention

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