A simplified diagram of a hydrogen-oxygen fuel cell (c) doc bDoc Brown's Chemistry

Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 7.5

7.5 Electrochemical cells (batteries) & fuel cell systems

How do electrochemical cells like simple batteries work? How does a zinc–carbon battery work? How does a NiCad or alkaline battery work? How does a lead–acid battery work? How does a fuel cell work?

GCSE/IGCSE reversible reactions & chemical equilibrium notes

and GCSE/IGCSE Notes on Electrochemistry

cell8Part 7 sub–index: 7.1 Half cell equilibria, electrode potential * 7.2 Simple cells notation and construction * 7.3 The hydrogen electrode and standard conditions * 7.4 Half–cell potentials, Electrochemical Series and using Eθcell for reaction feasibility * 7.5 Electrochemical cells ('batteries') and fuel cell systems * 7.6 Electrolysis and the electrochemical series * 7.7 Exemplar Questions, Appendix 1. The Nernst Equation, Appendix 2 Free Energy, Cell Emf and K

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibria and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 8. Phase equilibria–vapour pressure, boiling points and intermolecular forces


7.5 Electrochemical cells ('batteries') and fuel cell systems

PRIMARY CELLS are not rechargeable and are discarded (hopefully by safe recycling systems!)  after they run down when all the chemicals are used up ie no more chemical potential energy available

SECONDARY CELLS can be recharged after they have run down ie the discharge reactions producing the electricity are reversed to built up the store of chemical potential energy

FUEL CELLS produce electricity directly from gaseous of liquid fuels such as hydrogen or hydrocarbons with only safe waste products of water or carbon dioxide.

Important notation convention: Note that the +ve and ve electrode charges in battery cells/fuel cells are reversed compared to the electrodes in the process of electrolysis in electrolytic cells. This is because battery cells and electrolysis cells operate in 'opposite directions' to each other in terms of oxidation & reduction and electron flow. Therefore, we have, a ve cathode in electrolysis and a +ve cathode in battery/fuel cells AND a +ve anode in electrolysis and ve anode in battery/fuel cells. But, as in electrolysis, you still get reduction at the cathode and oxidation at the anode, so watch out for the () and (+) electrode signs!

  • Primary Cells

    • The first primary cells were galvanic cells in which the reactants are sealed in when manufactured and ready for immediate use i.e. the chemicals are capable of spontaneously reacting and the redox changes released energy as an electron flow (rather than heat energy). They cannot be recharged, and when they run down, that is the chemical reactants are completely depleted, they stop working and are discarded!

      • The copper–zinc Danielle Cell was one of the first useful batteries see 7.2 Simple cells notation and construction, though porous pots were used rather than beakers and a salt–bridge of filter paper + electrolyte.

    • The common ones such as the zinc–carbon batteries are used in torches, radios, cameras, flashlights, cameras etc.

    • Hopefully recycling of the materials will be increasingly possible as well as being worthwhile from the point of view of conserving valuable resources and minimising environmental pollution from poisonous metals or their compounds.

    • cell8 Dry cell zinc–carbon battery, 1.5V falling to 0.8V as reaction products build up.

      • In the zinc–carbon cell a rod of carbon cathode (+ convention) is set into a paste of zinc and ammonium chloride (weakly acid electrolyte) and fine particles of manganese(IV) oxide and carbon contained in a zinc anode (– convention) 'compartment'. Although called a 'dry' cell, the paste must contain water, which is thickened with e.g. starch.

      • Zn(s)|ZnCl2(aq),NH4Cl(aq)||MnO2(s)|MnO(OH)(s)|Cgraphite

        • Note the standard conventions in common use

          • the | notation indicates a phase boundary, solutes in the same phase are separated by a comma

            • the || notation 'divides' the two half cells and

              • the oxidation state increases 'towards' it

      • () anode discharging reaction (i) Zn(s) + 4NH3(aq) ==> [Zn(NH3)4]2+(aq) + 2e

      • (+) cathode discharging reaction (ii) MnO2(s) + NH4+(aq) + e ==> MnO(OH)(s) + NH3(aq)

      • overall working cell reaction (iii) Zn(s) + 4NH3(aq) + 2MnO2(s) + 2NH4+(aq)

        • ==> [Zn(NH3)4]2+(aq) + 2MnO(OH)(s) + 2NH3(aq)

          • from (iii) = {(i) + 2 x (ii)}

      • oxidation state changes: (i) oxidation Zn(0) ==> Zn(+2), (ii) reduction 2Mn(IV) ==> 2Mn(III) to balance

        • Some alternative equations you see in textbooks – they amount to the same chemical changes in the end eg via the formation of manganese(III) oxide

          • anode: Zn(s) + 2NH3(aq) ==> [Zn(NH3)2]2+(aq) + 2e

          • cathode:2MnO2(s) + 2H+(aq) + 2e ==> Mn2O3(s) + H2O(l)

          • overall: Zn(s) + 2MnO2(s) + 2NH4+(aq) ==>  [Zn(NH3)2]2+(aq) + Mn2O3(s) + H2O(l)

      • Advantages: Low cost and non–toxic materials.

      • Disadvantages: Cannot be recycled, can leak (weak acid electrolyte reacts with zinc), short shelf–life, unstable voltage and current (as battery 'runs down') and low power.

    • TOP and LINKSThe dry cell alkaline battery, 1.5–1.9V depending on constituents.

      • In the alkaline dry cell the electrolyte is the strong base sodium/potassium hydroxide contained in 'typically' zinc anode (–) compartment and a cathode of manganese(IV) oxide. Metals like cadmium or aluminium can be used as the anode, and copper, iron, lead, mercury, nickel and silver oxide can be used as cathode materials.

      • Zn(s)|ZnO(s)|OH(aq)||MnO2(s)|Mn(OH)2(s)|Cgraphite

        • Note the standard conventions in common use

          • the | notation indicates a phase boundary

            • the || notation 'divides' the two half cells and

              • the oxidation state increases 'towards' it

      • (–) anode discharging reaction (i) Zn(s) + 2OH(aq) ==> ZnO(s) + H2O(l) + 2e

      • (+) cathode discharging reaction (ii) MnO2(s) + 2H2O(l) + 2e ==> Mn(OH)2(s) + 2OH(aq)

      • overall cell reaction (iii) Zn(s) + MnO2(s) + H2O(l) ==> ZnO(s) + Mn(OH)2(s)

        • from (iii) = (i) + (ii)

      • oxidation state changes: (i) oxidation Zn(0) ==> Zn(+2), (ii) reduction Mn(IV) ==> Mn(II)

        • Some alternative equations you see in textbooks – they amount to similar chemical changes in the end eg via the formation of manganese(III) oxide rather than manganese(II) hydroxide (I'm afraid different sources quote different chemistry!)

          • anode: Zn(s) + 2OH(aq) ==> ZnO(s) + H2O(l) + 2e

          • cathode: 2MnO2(s) + H2O(l) + 2e ==> Mn2O3(s) + 2OH(aq)

          • overall: Zn(s) + 2MnO2(s) ==> ZnO(s) + 2Mn2O3(s)

      • Advantages: Low cost and non–toxic materials. The alkaline electrolyte does not readily react with zinc (compare Zn–C cell above) giving a much longer shelf–life (5 years) and the current and voltage are steady (handy in smoke alarms!) due to the strong base/alkali electrolyte having a smaller resistance the ammonium chloride–carbon paste.

      • Disadvantages: Cannot be recycled, more expensive due to extra sealing and low power.

  • Fuels cells are a development of primary cells but with one significant difference from their predecessors, the chemical potential energy source or 'fuel' can be continually fed in to give the cell a long active life.

    • The hydrogen–oxygen fuel cell

    • A simplified diagram of a hydrogen-oxygen fuel cell (c) doc b equation

      It uses costly platinum electrodes and an acid electrolyte such as phosphoric acid, H3PO4

      1. oxidation 2H2(g) ==> 4H+(aq) + 4e  (at negative anode electrode*)
      2. reduction O2(g) + 4H+(aq) + 4e ==> 2H2O(l) (at positive cathode electrode*)
      3 = 1 + 2 redox 2H2(g) + O2(g) ==> 2H2O(l)

      * Note the +ve and ve electrode charges are reversed compared to electrolysis, because the system is operating in the opposite direction. But, as in electrolysis, you still get reduction at the cathode and oxidation at the anode.

    • The hydrogen–oxygen cell with an alkaline electrolyte is known as the 'alkali fuel cell' and is used in NASA's space shuttle craft.

    • (–) anode reaction (i) 2H2(g) + 4OH(aq) ==> 4H2O(l) + 4e

    • (+) cathode reaction (ii) O2(g) + 4H+(aq) + 4e ==> 2H2O(l)

    • overall cell reaction: 2H2(g) + O2(g) ==> 2H2O(l)

    • The electrolyte is the alkali potassium hydroxide solution, KOH(aq).

    • In both acid or alkaline hydrogen–oxygen fuel cells the oxidation state changes are

      • (i) oxidation H(0) ==> H(+1), (ii) reduction O(0) ==> O(–2)

    • Advantages: Can run on conventional fuels without the need of expensive metals except for the catalyst

    • Disadvantages: Quite costly at the moment eg the platinum catalyst

    • Organic fuel cells are described in Advanced Redox Chemistry Part III (Organic reactions)

  • TOP and LINKSSecondary Cells (electrical 'accumulators')

    • Secondary cells are galvanic cells that must be charged before they can be used and rechargeable many times. In the charging process, the spontaneous–feasible cell reaction that produces electrical energy is reversed, so building up the chemical potential of the cell system.

    • cell9 Lead–acid storage battery, 2 V. (usually 6 in series to give 12V supply).

      • The electrodes are initially hard lead–antimony alloy plates coated in a paste of lead(II) sulphate encased in dilute sulphuric acid. During the first charging some of the lead(II) sulphate is reduced lead(0) on one of the electrodes (this will acts as the (–) anode in discharging). Simultaneously in charging, lead(II) sulphate is oxidised to lead(IV) oxide on the other electrode which acts as the cathode (+) in discharging.

      • Pb(s)|PbSO4(s)|H+(aq),HSO4(aq)||PbO2(s)|PbSO4(s)|Pb(s)

        • Note the standard conventions in common use

          • the | notation indicates a phase boundary, solutes in the same phase are separated by a comma

            • the || notation 'divides' the two half cells and

              • the oxidation state increases 'towards' it

      • (–) anode discharging reaction (i) Pb(s) + HSO4(aq) ==> PbSO4(s) + H+(aq) + 2e

      • (+) cathode discharging reaction (ii) PbO2(s) + 3H+(aq) + HSO4(aq) + 2e ==> PbSO4(s) + 2H2O(l)

      • working cell reaction (iii) PbO2(s) + 2H+(aq) + 2HSO4(aq) + Pb(s) ==> 2PbSO4(s) + 2H2O(l)

      • oxidation state changes: (i) oxidation Pb(0) ==> Pb(II) : (ii) reduction Pb(IV) ==> Pb(II)

      • The charging reactions will be the opposite of (i) and (ii)

      • Advantages: Inexpensive, high power density (can car starter motor as well as lights), long shelf life, readily recharges, so has a long working life of many years.

      • Disadvantages: Lead needs to be recycled to avoid environmental contamination, sometimes generates hydrogen gas at the cathode when charging (explosive in air + spark) – though batteries seem to be made of a high standard these days in completely sealed units that last many years.

      • Uses: Car batteries.

    • The NiCad Cell, 1.25 V.

      • diagram?

      • Cd(s)|Cd(OH)2(s)|KOH(aq)||Ni(OH)3(s)|Ni(OH)2(s)|Ni(s)

        • Note the standard conventions in common use

          • the | notation indicates a phase boundary

            • the || notation 'divides' the two half cells and

              • the oxidation state increases 'towards' it

      • (–) anode discharging reaction (i) Cd(s) + 2OH(aq) ==> Cd(OH)2(s) + 2e

      • (+) cathode discharging reaction (ii) 2Ni(OH)3(s) + 2e ==> 2Ni(OH)2(s) + 2OH(aq)

      • overall cell reaction (iii) Cd(s) + 2Ni(OH)3(s) ==> Cd(OH)2(s) + 2Ni(OH)2(s)

      • oxidation state changes: (i) oxidation Cd(0) ==> Cd(II), (ii) reduction Ni(III) ==> Ni(II)

      • The charging reactions will be the opposite of (i) and (ii)

      • Advantages:

      • Disadvantages: Cadmium is a toxic metal.

      • Uses: Portable computers

  • The voltage and power available from a battery or cell

    • The voltage depends primarily on the materials used in the chemical process generating the electrical energy.

    • Since the voltage is small from an individual cell, (typically 0.4 to 2V), several cells can be assembled in parallel to increase the voltage.

    • The power primarily depends on the amount of material and how fast the chemicals can react.

      • For a single cell the voltage will depend on the half–cell potentials of chemicals employed, but the current flow depends on the bulk reaction rate of the chemicals.

TOP and LINKS

WHAT NEXT?

Part 7 sub–index: 7.1 Half cell equilibria, electrode potential * 7.2 Simple cells notation and construction * 7.3 The hydrogen electrode and standard conditions * 7.4 Half–cell potentials, Electrochemical Series and using Eθcell for reaction feasibility * 7.5 Electrochemical cells ('batteries') and fuel cell systems * 7.6 Electrolysis and the electrochemical series * 7.7 Exemplar Questions, Appendix 1. The Nernst Equation, Appendix 2 Free Energy, Cell Emf and K

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibria and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 8. Phase equilibria–vapour pressure, boiling points and intermolecular forces

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