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Revision notes on chemical equilibrium - how simple cells, batteries & fuel cells work Advanced Level Theoretical-Physical Chemistry

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

Doc Brown's A Level Chemistry Revision Notes

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

7.5 Electrochemical cells (batteries) and fuel cell systems

Redox reactions take place in electrochemical cells ('batteries') where electrons are transferred from the reducing agent (e.g. 'fuel') to the oxidising agent (e.g. O2, MnO2)  indirectly via an external circuit. A potential difference is created that can drive an electric current to do work. Electrochemical cells have very important commercial applications as a portable supply of electricity to power electronic devices such as mobile phones, tablets and laptops. On a larger scale, they can provide energy to power a vehicle. So, 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? How does a lithium ion battery work? The advantages and disadvantages of electrochemical cells - primary cells, secondary cells, fuel cells, rechargeable cells and non-rechargeable cells etc.

Chemical Equilibrium Notes Part 7 Index

cell8

 

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

Cells ('batteries') can be:

(i) non-rechargeable (irreversible) where the chemicals are used up

(ii) rechargeable, where the chemistry generating the electrical current can be reversed, so charging up the battery again, AND the chemicals are not used up,

 (iii) or fuel cells, which must be continuously fed chemicals e.g. fuel plus oxidant to maintain the chemistry producing the electrical current. Fuel cells used to generate an electric current do not need to be electrically recharged.

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.

    • The Leclanché cell is a primary cell or battery invented and patented by the French scientist Georges Leclanché. The battery contains a conducting solution (electrolyte paste) of ammonium chloride and water, a cathode (positive terminal) of carbon, a depolarizer of manganese dioxide (oxidizer), and an anode (negative terminal) of zinc (reductant). How it works is described and explained below.

    • 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.

    • The 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 - it can be confusing !

    • 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)

  • Secondary 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 lithium cell (the rechargeable 'Lion battery')

      • A popularly used commercial battery that powers many a computer!

      • The chemicals involved in one type are lithium metal, graphite and a lithium-cobalt oxide.

      • The discharging simplified electrode reactions in a lithium cell to generate electricity are as follows:

      • At the positive electrode (reduction): Li+ + CoO2 + e–  ==>  Li+[CoO2]–

      • Negative electrode via graphite (oxidation): Li ==> Li+ + e–

      • When the battery is being recharged, these reactions are reversed.

  • 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.

  • The advantages and disadvantages of using different types of electrochemical cells

    • Batteries - commercial electrochemical cells, have provided us with an extremely useful source of power, both in the home and industry.

    • They are convenient to carry around in phones, computers, torches and fitted into a variety of vehicles.

    • You can take power to remote places without the need for expensive cabling, and rechargeable batteries via solar cells can be invaluable in this context e.g. powering a telephone box in some off the beaten track location in some way out location in the countryside.

    • The orbiting space station and spacecraft use fuels cells to provide power and collect the water produced for domestic use! You can't get a more remote location than outer space or the moon!

    • BUT, there are drawbacks to all this electrochemical cell technology and some of the 'pros and cons' are discussed below.

    • Non-rechargeable cells

      • Non-rechargeable batteries have a long history and cheap to produce. They can be made in all sorts of sizes, with a range of voltages and applied to a wide range of devices from fire alarms to torches.

      • BUT, all the chemicals are 'used up' and the voltage-power declines as they come towards the end of their working life.

      • They can't be recharged and are often just thrown away without any attempt at recycling which is waste of both energy and materials.

    • Rechargeable cells

      • Using rechargeable cells obviously reduces the total number of cells used and so far less will be added to the 'scrapheap' of used 'batteries'.

      • Batteries, such as those in cars are constantly being recharged when the vehicles is moving and provide power for lights, starter motor, electric windows and ignition and any other automated device.

      • Rechargeable cells can be used in conjunction with solar panels, these work in daylight, but storing some of the electrical energy enables it to be used at night when the solar panels are inactive.

      • However, the cadmium in nickel-cadmium rechargeable cells is toxic and must be carefully disposed of and the cadmium recovered, hopefully to be recycled.

      • The lead and lead(II) sulfate in lead-acid car batteries can be recovered and recycled in new batteries and they are a better design these days and last longer now in 2017+ than they did back in the 1960s as I recall !

      • The design of lithium-ion cells has improved over the years and are now widely used in phones, computers and even larger machines like disability scooters. These Li-ion batteries are quite costly, but their useful lifetime is increasing as is the amount of electrical energy they can store efficiently e.g. several hours use on your computer/phone from a relatively compact source!

    • Fuel cells

      • The main big advantage of fuel cells of the hydrogen-oxygen variety is that the only waste product is harmless water.

      • Theoretically, fuel cells can replace the power generation from conventional petrol and diesel engines, thus reducing emission of greenhouse gases and harmful nitrogen/sulfur oxides and hydrocarbon/carbon particulate pollution.

      • However, most industrial hydrogen (e.g. for the Haber synthesis of ammonia) is actually made from fossil fuel hydrocarbon molecules!

      • Hydrogen is also highly flammable-explosive with issues over safe storage and distribution.

      • One possibility is large-scale solar powered electrolysis of water and collect the gases, but such an enterprise seems a long way off and the health and safety issues remain!

Chemical Equilibrium Notes Part 7 Index

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