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

FUEL CELLS

e.g. the hydrogen-oxygen fuel cell

Doc Brown's Chemistry KS4 science–chemistry GCSE/IGCSE/A level Revision

ELECTROCHEMISTRY revision notes on electrolysis, cells, experimental methods, apparatus, batteries, fuel cells and industrial applications of electrolysis

11. FUEL CELLS

how do they work?  * See also 10. Simple cells (batteries)

The principles of fuel cells are explained with particular attention to the hydrogen–oxygen fuel cell. These revision notes on how fuels work and how fuel cells are constructed should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.

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11. FUEL CELLS e.g. the hydrogen-oxygen fuel electrical cell

  • Fuel cells have to be supplied by an external source of fuel (e.g. hydrogen) and an oxidant e.g. oxygen or air.

    • The hydrogen or any other fuel is oxidised electrochemically inside the fuel cell to produce a potential difference i.e. a voltage capable of producing a working current.

    • The overall chemical reaction in a hydrogen fuel electrochemical cell involves the oxidation of hydrogen by oxygen to produce only water.

    • Hydrogen fuel cells offer an alternative to rechargeable cells and batteries.

    • A fuel cell will produce a voltage until on of the reactants is used up.

  • Hydrogen gas can be used as fuel.

    • hazard warning symbol for a flammable materialIt burns with a pale blue flame in air reacting with oxygen to be oxidised to form water.

      • hydrogen + oxygen ==> water

        • 2H2(g) + O2(g) ==> 2H2O(l) 

        • Its a very exothermic reaction, releasing lots of heat energy when burnt, remember the 'squeaky pop' lit splint test for hydrogen!

        • But know think of it in terms of bond breaking and bond making.

        • H–H + H–H + O=O ==> H–O–H + H–O–H

        • As you can see from the diagram, theoretically, more energy is released by forming water, than was absorbed in breaking up the original hydrogen and oxygen molecules.

          • The energy level of the fuel + oxygen is higher than that of water.

          • Therefore the fuel + oxygen mixture has the potential to release potential chemical energy as heat energy OR, crucially, electrical energy, if the reaction can be made to go via oxidation and reduction reactions involving electron transfer.

          • See this page for more details of the full bond energy calculations

        • A simplified diagram of a hydrogen-oxygen fuel cell (c) doc bA fuel cell is a device that functions like an electrical cell or battery that you supply with a fuel (liquid or gas) and an oxidant (oxidising agent, usually oxygen gas) which react together in a redox reaction (oxidation + reduction) to release energy as a flow of electrons i.e. an electrical current capable of doing useful work.

        • The oxidation of the fuel and the reduction of the oxidant reactions take place on electrodes (see diagram on right, all fully explained further down the page).

        • Fuel cells essentially a device to release the chemical potential energy of combustible fuels as electrical energy at a much lower temperature and without the flame!

        • So, the idea of a hydrogen–oxygen fuel cell is to release the energy from hydrogen reacting with oxygen, not as heat as in normal combustion in air, but as useful electrical energy i.e. a practical electricity supply.

        • Fuel cells are more efficient than conventional power stations or batteries etc. because the electrical energy is directly generated from the chemical reaction between the oxidant and the fuel.

          • In fossil fuel power stations, motor vehicles, coal/gas fires etc. quite a percentage of energy is lost as waste heat.

          • With a fuel cell there are fewer stages in producing the useful energy, so there is less opportunity to lose potentially useful energy e.g. waste heat, friction from moving parts etc.

    • A hydrogen–oxygen fuel cell is a non–polluting clean fuel since the only combustion product is water.

    • It would be an ideal fuel on this basis e.g. for motor vehicles, but that's not the only factor to consider!

    • It would be ideal if the hydrogen fuel could be manufactured by electrolysis of water e.g. using solar cells.

    • hazard warning symbol for an explosive materialHydrogen can be used to power fuel cells.

    • It all sounds wonderful BUT, still technological problems to solve for large scale manufacture and distribution of 'clean' hydrogen gas or use in generating electricity AND its rather an inflammable explosive gas!

    • See fuel survey and alternative fuels like biofuels

  • Fuel cells are 'battery systems' in which two reactants can be continuously fed in.

    • They must undergo an oxidation–reduction (redox) reaction.

    • The consequent redox chemistry produces a working current e.g. hydrogen, ethanol and hydrocarbons can be oxidised with oxygen in a fuel cell, rather than conventional combustion.

    • Fuel cells have the big advantage of not requiring charging up, as long as you keep on supplying the fuel and oxidant.

    • The motor vehicle industry is looking at fuel cell cars, but lately electric cars with better quality modern battery technology and convenient charging facilities seem to be more in the news.

  • Plus points in favour of using fuel cells in cars and other road vehicles – advantages

    • Fuel cells do not produce the usual pollutants like sulfur oxides (acid rain), nitrogen oxides and carbon monoxide (harmful gases from traffic pollution).

    • Unless an organic fuel like a hydrocarbon or an alcohol is used, there will be no greenhouses gases like carbon dioxide, because hydrogen combustion only produces harmless water.

      • Neither do you produce harmful and polluting gases like carbon monoxide, nitrogen oxides, sulfur oxides and hydrocarbon particulates.

    • Fuel cells could be used in countries with little oil and make them less dependant on costly imported oil.

    • Hydrogen-oxygen fuel cells create water and since the hydrogen will have to be made from abundant water supplies in the long run, it is effectively a renewable resource, unlike fossil fuels like petrol or diesel.

    • Fuel cells are much more efficient than fossil fuel power stations or costly chemical batteries etc. because the electrical energy is generated directly from the chemical reaction.

      • In fossil fuel power stations, motor vehicles, coal/gas fires etc. too much energy is lost as waste heat.

      • Fossil fuels are a finite supply, whereas if we can get hydrogen from water, and the air has lots of oxygen, there is effectively, if not practical at the moment, an infinite supply of fuel.

      • With a fuel cell there are fewer stages in producing the useful electrical energy, so there is less chance of  losing useful energy e.g. waste heat, friction from moving parts etc.

      • Maybe in the long–run we can get rid of polluting fossil fuelled road vehicles and batteries containing harmful chemicals e.g. toxic metal compounds of cadmium.

      • Although current electric cars are improving in the design, the batteries are very costly can only be recharged a limiting number of times before requiring replacement. Batteries also store less energy than fuel cells (which still have to be 're-fuelled').

    • .AND fuel cells have found use in spacecraft (e.g. space shuttle), satellites and orbiting space stations.

      • Fuel cells can provide a convenient source of electrical power in the space industry.

      • Fuel cells can be made reasonably lightweight and compact saving both space and weight.

      • Hydrogen and oxygen may be used from the fuel tanks of spacecraft.

      • Fuel cells do not have moving parts which could go wrong.

      • The final product is water, which could be used as drinking water, reducing the initial payload in the spacecraft.

      • Since water is the only product from a hydrogen-oxygen fuel cell, there are no waste pollution products to deal with.

  • Issues over the use of fuel cells in transport systems or as a major energy resource – disadvantages

    • Fuel cells cannot be used for large–scale energy production, so conventional fossil fuel or nuclear power stations still have an important future.

    • Hydrogen is a gas and requires a much larger storage volume compared to fossil fuels like petrol.

    • Safe storage is an issue, especially as it would be stored under high pressure to decrease the storage space required.

      • This immediately makes leaks and accidents more likely to happen because hydrogen is a highly flammable explosive gas – too easily ignited, remember the 'squeaky pop' lit splint test for hydrogen!

    • There is no efficient means of mass producing hydrogen.

      • Efficient large scale technology is not yet developed to produce hydrogen on a large scale eg from electrolysis using solar power electricity – photovoltaic power system, wind turbines or hydroelectric power.

      • Although water is cheap and plentiful, it requires expensive electrical energy to electrolyse water to split it into hydrogen and oxygen.

      • Electrolysis of acidified water is expensive because electricity is expensive and much is still made by burning fossil fuels.

      • Not only that, most electricity in the world is generated from burning fossil fuels!

      • Most hydrogen used in industry is made from fossil fuel hydrocarbons, which won't last forever. (see making hydrogen for the Haber Synthesis of ammonia).

      • Apart from the fact that the electrode catalysts are costly, disposing of used cells has its problems e.g. may contain toxic metal compounds.

  • More on hydrogen's potential use in fuel and energy applications includes powering vehicles, running turbines or fuel cells to produce electricity, and generating heat and electricity for buildings and very convenient for remote and compact situations like the space shuttle.

    • Fuel cells were developed in the 1960s as part of the USA NASA's space exploration programme to provide electrical power.

    • Fuel cells were/are used in lunar landing vehicles, space stations orbiting the Earth etc.

    • They are more practical and robust than solar cell panels and definitely safer than small nuclear power units.

    • When hydrogen is the fuel, the product of its oxidation is water, so this is potentially a clean non–polluting and non–greenhouse gas? fuel.

    • Most fuel cells use hydrogen, but alcohols and hydrocarbons can be used.

    • A fuel cell works like a battery but does not run down or need recharging as long as the 'fuel' supply is there.

  • It will produce electricity and heat as long as fuel (hydrogen) is supplied.

  • Notes based on the DIAGRAMS and the CHEMISTRY of a hydrogen–oxygen fuel cell

    • A typical acid fuel cell (above) consists of two electrodes consisting of a negative electrode (or anode) and a positive electrode (or cathode) which are sandwiched around an electrolyte (conducting salt/acid/alkali solution of free ions).

    • Hydrogen (or other fuel) is fed to the (–) anode, and oxygen is fed to the (+) cathode.

    • The platinum catalyst activates the hydrogen atoms/molecules to separate into protons (H+) and electrons (e), which take different paths to the (+) cathode.

    • The electrons go through an external circuit, creating a flow of electricity e.g. to light a bulb.

    • The protons (H+) migrate through the electrolyte and pass through the semi–permeable membrane to the cathode, where they and electrons reduce the oxygen to water.

    • In both acid or alkaline hydrogen–oxygen fuel cells, oxygen is the oxidising agent (oxidant, gets reduced) and hydrogen (fuel, gets oxidised) is the reducing agent.

      • The hydrogen is fed into the anode compartment and the oxygen into the cathode compartment.

      • The oxidation of the fuel (hydrogen) at the anode electrode provides electrons that flow round (via some device e.g. bulb, motor etc.) to the cathode electrode to effect the reduction of the oxygen.

      • So you have simultaneous oxidation and a reduction, therefore overall a redox reaction.

    • In an alkaline hydrogen–oxygen fuel cell, the construction is quite similar, but it is hydroxide ions (OH) that migrate in the potassium hydroxide solution electrolyte.

    • Each cell only produces a small voltage (typically 0.4 to 1.0V) so many cells have to be put together in series to give a bigger working voltage.

    • Note on reverse reaction water ==> hydrogen + oxygen

      • If there is spare electricity from another source available, you can run the fuel cell in reverse and electrolyse the water to make hydrogen and oxygen (acting as an electrolyser).

      • The two gases are stored, and when extra electricity or heat needed, the fuel cell can then be re–run using the stored gaseous fuel.

      • This is called a regenerative fuel cell system.

      • You can use solar energy from external panels on the space shuttle to do this, and use the fuel during the 'darkness of night'.

 


A simplified diagram of a hydrogen-oxygen fuel cell (c) doc bChemical descriptions of  hydrogen–oxygen fuel cell

In the electrode equations I have included the 'subscripted' state symbols (aq), (g) and (l) where appropriate.

Example 1.

This acid hydrogen–oxygen fuel cell uses costly platinum electrodes and an acid electrolyte such as phosphoric acid, H3PO4  (diagram on right for an acid hydrogen–oxygen fuel cell lighting a bulb)

The electrode equations for an acid hydrogen–oxygen fuel cell are ...

1. oxidation electrode half–equation, hydrogen atoms/molecules lose electrons

2H2(g)  ==>  4H+(aq) + 4e      (half cell equation for the negative anode electrode*)

or     2H2(g)  –  4e  ==>  4H+(aq)

electrode equation 1. has been doubled up to balance the electrons involved and so that 1. + 2. equals the overall 'normal' molecular equation for the hydrogen fuel consumption.

2. reduction electrode half–equation, oxygen atoms/molecules gain electrons

O2(g) + 4H+(aq) + 4e ==> 2H2O(l)     (half cell equation for the positive cathode electrode*)

Adding the two electrode equations 1 + 2 gives the overall equation, which is the same as the normal combustion equation, in a way its a sort of 'cooler' combustion without the flame!

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.


Example 2.

You can also run a hydrogen–oxygen fuel cell with an alkaline electrolyte [e.g. potassium hydroxide solution, KOH(aq)], but still with costly electrodes

In the case of the alkaline hydrogen–oxygen fuel cell the different electrode equations are ...

1. oxidation electrode half–equation, hydrogen atoms/molecules lose electrons, to give hydrogen ions which immediately combine with the hydroxide ions to form water.

2H2(g) + 4OH(aq) – 4e  ==>  4H2O(l)    (oxidation at the anode)

or    2H2(g) + 4OH(aq)  ==>  4H2O(l) + 4e

2. reduction electrode half–equation, oxygen atoms/molecules gain electrons and in combination with water molecules are reduced to hydroxide ions.

O2(g) + 2H2O(l) + 4e–  ==>  4OH(aq)       (reduction at the cathode)

Again, the overall equation is same as the normal combustion equation by adding equations 1. and 2. together.

2H2(g) + O2(g)  ==>  2H2O(l)

Example 3 for more advanced level chemistry students

The diagram on the right shows the same arrangement for a fuel cell, as for hydrogen, but you can use an organic fuel like an alcohol e.g. ethanol (C2H5OH) or methanol (CH3OH), which burn exothermally if ignited, but here made to release their chemical potential energy as electrical energy.

The overall cell reactions are ...

C2H5OH + 3O2  ==> 2CO2 + 3H2O + energy

CH3OH + 3/2O2 ==> CO2 + 2H2O + energy

For more details click on the Redox 3 link below


See also 10. Simple cells (batteries)


Other notes on ADVANCED chemistry pages:

Alkaline hydrogen–oxygen fuel cell is described in (c) doc b Equilibria Part 7 Redox Chemistry

and (c) doc b organic fuel cells are described in Redox Chemistry Part 3.

 


ELECTROCHEMISTRY INDEX:  1. INTRODUCTION to electrolysis - electrolytes, non-electrolytes, electrode equations, apparatus 2. Electrolysis of acidified water (dilute sulfuric acid) and some sulfate salts and alkalis 3. Electrolysis of sodium chloride solution (brine) and bromides and iodides 4. Electrolysis of copper(II) sulfate solution and electroplating with other metals e.g. silver 5. Electrolysis of molten lead(II) bromide (and other molten ionic compounds) 6. Electrolysis of copper(II) chloride solution 7. Electrolysis of hydrochloric acid 8. Summary of electrode equations and products 9. Summary of electrolysis products from various electrolytes 10. Simple cells (batteries) 11. Fuel Cells e.g. the hydrogen - oxygen fuel cell 12. The electrolysis of molten aluminium oxide - extraction of aluminium from bauxite ore & anodising aluminium to thicken and strengthen the protective oxide layer 13. The extraction of sodium from molten sodium chloride using the 'Down's Cell' 14. The purification of copper by electrolysis 15. The purification of zinc by electrolysis 16. Electroplating coating conducting surfaces with a metal layer 17. Electrolysis of brine (NaCl) for the production of chlorine, hydrogen & sodium hydroxide AND 18. Electrolysis calculations


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