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Doc Brown's Chemistry KS4 science-chemistry GCSE/IGCSE/AS Revision Notes

Extra Electrochemistry Information Notes on Electrolysis, cells, batteries, fuel cells and industrial applications of electrolysis

This section starts with an introduction to electrolysis explaining which liquids conduct and why can some liquids conduct electricity and others cannot? - electrolytes, electrodes, what happens on electrode surfaces, the products of electrolysis. There are tables of electrode reactions, descriptions of experimental methods of electrolysis and a summary table of the electrolysis products from many common melts or aqueous solutions that undergo electrolysis when a d.c. electric current is passed through them. There is also a list of links to revision notes on the application of electrolysis to various industrial processes. Simple cells, batteries and fuel cells are also described. This page should also prove useful for coursework projects in electrolysis e.g. help with ideas, apparatus, experimental methods, products collection, electrode equations and calculations etc.

(c) doc bIndex for this page 1. Introduction to electrolysis * 2. Summary of practical experimental arrangements for laboratory (lab) electrolysis experiments and electrode reactions * 3a. Summary of the products of selected electrolysis processes * 3b. Summary of Industrial Electrolysis Process links * 4. Simple cells or batteries * 5. Fuel cells

(c) doc bAssociated LINKS to other pages on this site: Metal Extraction * Industrial chemistry * Metal reactivity (redox introduction) * Types of chemical reaction * Electrolysis calculations


1. Introduction to electrolysis - electrolytes and non-electrolytes

Electrolysis is the process of electrically inducing chemical changes in a conducting melt or solution e.g. splitting an ionic compound into the metal and non-metal.

    • What makes up a circuit in cells - batteries or electrolysis? What carries the current?
    • These materials carry an electric current via freely moving electrically charged particles, when a potential difference (voltage!) is applied across them, and they include:
    • All metals (molten or solid) and the non-metal carbon (graphite).
      • This conduction involves the movement of free or delocalised electrons (e- charged particles) and does not involve any chemical change.
    • Any molten or dissolved material in which the liquid contains free moving ions is called the electrolyte and can conduct an electrical current. (see non-electrolyte)
      • top sub-indexIons are charged particles e.g. Na+ sodium ion, or Cl- chloride ion, and their movement or flow constitutes an electric current, in other words the electrolyte consists of a stream of moving charged particles.
      • What does the complete electrical circuit for electrolysis consist of?
        • There are two ion currents in the electrolyte flowing in opposite directions 
          • positive cations e.g. sodium Na+ are attracted to the negative cathode electrode,
          • and negative anions e.g. chloride Cl- are attracted to the positive anode electrode,
          • and it is possible to demonstrate this flow using a coloured ion experiment.
          • remember no electrons flow in the solution, but they do flow in metal wires or carbon (graphite) electrodes of the external circuit.
          • Three 'connected' sub-notes: (i) The greater the concentration of the electrolyte ions, the lower the electrical resistance of the solution. This is because there are more ions present to carry the current e.g. if the voltage (V, volts) is kept constant, the current flowing (I, amps) will steadily increase as the concentration of the electrolyte is increased. (ii) If the electrolyte (ion) concentration is kept constant, the current will steadily increase with increase in voltage just like any other electrical circuit because the increase in electrical field effect from the increased p.d. (voltage) will force the ion flow at a greater rate. (iii) So, increase in ion concentration (salts, acids etc.) OR increase in voltage will increase the speed of electrolysis i.e. the electrode reactions, whether it involves gas formation or electroplating metals etc.
        • The circuit of 'charge flow' is completed by the electrons moving around the external circuit e.g. copper wire, metal or graphite electrode, from the positive to the negative electrode (this e- flow from +ve to -ve electrode perhaps doesn't make sense until you look at the electrode reactions later on).
      • The molten or dissolved materials are usually acids, alkalis or salts and their electrical conduction is usually accompanied by chemical changes e.g. decomposition.
  • The chemical changes occur at the electrodes which connect the electrolyte liquid containing ions with the external d.c. electrical supply.
    • If the current is switched off, the electrolysis process stops.
  • Non-electrolytes are liquids or solutions that do not contain ions, do not conduct electricity readily and cannot undergo the process of electrolysis e.g. ethanol (alcohol), sugar solution etc. and are usually covalent molecules.
  • How do we know that ions move to the electrodes when a d.c. current is applied?
    • The coloured ion experiment described below illustrates how you can show the slow movement of ions when a voltage is applies across a conducting solution of ions - the electrolyte.

A simple experiment to show the movement of coloured ions

A simple experiment to show the movement of coloured ions (c) doc b

A rectangle of filter paper is soaked in an ammonia-ammonium chloride solution and mounted on a microscope slide. The paper is connected to a d.c. supply with clips. A 'line' of copper chromate solution is placed in the middle of the paper and the current switched on. The copper chromate is green-brown in solution but gradually it disappears and separates, in different directions, into a yellow and blue bands. The yellow band is due to negative chromate ions, CrO42--, moving towards the positive electrode. The blue band is due to positive copper ions, Cu2+, moving towards the negative electrode. All due to opposite charges attracting in the electric field produced by the potential difference (the voltage!).

top sub-index

  • Liquids that conduct must contain freely moving ions to carry the current and complete the circuit.
    • You can't do electrolysis with an ionic solid!, the ions are too tightly held by chemical bonds and can't flow from their ordered situation!
    • When an ionically bonded substances are melted or dissolved in water the ions are free to move about.
      • However some covalent substances dissolve in water and form ions.
      • e.g. hydrogen chloride HCl, dissolves in water to form 'ionic' hydrochloric acid H+Cl-(aq) 
  • The solution of ions (e.g. salts, acids etc.) or melt of ions (e.g. chlorides, oxides etc.) is called the electrolyte which forms part of the circuit. The circuit is completed by e.g. the external copper wiring and the (usually) inert electrodes like graphite (form of carbon) or platinum AND electrolysis can only happen when the current is switched on and the circuit complete.
  • Diagram showing the industrial electrolysis process to extract sodium metal from sodium chloride salt (c) doc bELECTROLYSIS SPLITS a COMPOUND:
    • When substances which are made of ions are dissolved in water, or melted material, they can be broken down (decomposed) into simpler substances by passing an electric current through them.
    • This process is called electrolysis.
    • Since it requires an 'input' of energy, it is an endothermic process.
  • During electrolysis in the electrolyte (solution or melt of free moving ions) ...
    • ... positive metal or hydrogen ions move to the negative electrode (cations attracted to cathode), e.g. in the diagram, sodium ions Na+ , move to the negative electrode (-ve),
    • and negatively charged ions move to the positive electrode (anions attracted to anode), e.g. in the diagram, chloride ions Cl-, move to the positive electrode (+ve). 
  • The diagram shows the industrial electrolysis process (in a Down's Process Cell) to extract sodium metal from sodium chloride (common salt).
  • During electrolysis, gases may be given off, or metals dissolve or are deposited at the electrodes.
    • Metals and hydrogen are formed at the negative electrode from positive ions by electron gain (reduction), e.g. in molten sodium chloride
      • sodium ions change to silvery grey liquid sodium
        • Na+ + e- ==> Na  (a reduction electrode reaction)
    • and non-metals e.g. oxygen, chlorine, bromine etc. are formed from negative ions changing on the positive electrode by electron loss (oxidation), e.g. in molten sodium chloride
      • chloride ions change to green chlorine gas
        • 2Cl- - 2e- ==> Cl2   or  2Cl- ==> Cl2 + 2e-  (an oxidation electrode reaction)
    • The electrons released by the oxidation at the positive anode, flow round through the anode and wire to the positive cathode and so bring about the reduction i.e. of the sodium ion.
  • In a chemical reaction, if an oxidation occurs, a reduction must also occur too (and vice versa) so these reactions 'overall' are called redox changes.
    • You need to be able to complete and balance electrode equations or recognise them and derive an overall equation for the electrolysis.
    • Below is list of some of the most common electrode equations you will encounter and the experimental arrangements are shown below them in section 2.

 top sub-index

2a. Summary of arrangements for laboratory electrolysis experiments

Examples of types of apparatus that can be used in schools and colleges

I'm happy to add other examples if they seem appropriate

Gas tests and this section followed by examples of electrode reactions to which the equation numbers refer.

Electrolysis Cell to collect cathode and anode gases (c) doc bMethod 1a

Aqueous solutions with inert electrodes (carbon or platinum)

Graphite electrodes 'upwardly' dipped into an solution of the electrolyte. The cell can be made from plastic pipe and a big rubber bung with two holes in it. A more elaborate format is to use a Hoffman Voltammeter (see below) using platinum electrodes and accurately calibrated collecting tubes like burettes.

a crucible and carbon electrode system for electrolysing a molten salt (c) doc bMethod 2

Molten salts with carbon electrodes

Carbon (graphite) electrodes dipped into molten salt which has been strongly heated in a crucible. It is difficult to collect the gases at the electrodes! The salts may be very high melting, so sometimes a small amount of another salt impurity is added to lower the melting point.

Diagram of an electrolysis cell for purifying copper of copper cathode plating (c) doc bMethod 3

Metal electrodes dipped in aqueous salt solutions

 M = metal, usually dipped in the metal sulphate solution.

For electroplating in general: The (-) is made the metal/conducting surface to be coated, and the (+) is made of the plating metal which dissolves and replaces any deposit formed on the (-) electrode

For more details see

Electroplating details and

Extraction and purification of copper

Brine - concentrated sodium chloride solution (brine) with carbon (graphite) gives equal volumes of hydrogen gas (-) and green chlorine gas (+) with sodium hydroxide left in solution, via equations 2 and 3. However in very dilute solution, reaction 8 occurs too giving oxygen gas as well as chlorine gas. Molten lead(II) bromide gives silver beads of lead (-) and brown fumes of bromine (+), illustrated above, see equations 2 and 10. Copper(II) sulphate with copper electrodes, the copper deposits (-) and the copper dissolves (+), equations 4 and 5. The blue colour of the Cu2+ ions stays constant because Cu deposited = Cu dissolved. Both involve a 2 electron transfer so it means mass of Cu deposited (eq 4) = mass of Cu dissolving (eq 5). Compare with carbon (graphite) electrodes.
Copper(II) sulphate with carbon (graphite) electrodes, the copper deposits (-) and oxygen gas (+), equations 4 and 8. Molten sodium chloride gives silvery sodium (+) and pale green chlorine gas (-), see equations 2 and 11. See also extraction of sodium in introduction. Compare with brine solution. top sub-index
Copper(II) chloride with carbon (graphite) electrodes, the copper deposits (-) and chlorine gas (+), equations 2 and 4. Compare with copper electrodes. Molten anhydrous zinc chloride gives zinc (+) and chlorine (-), equations 1 and 2  
Diagram of the Hofmann Voltameter to collect gases from electrolysis of acid or salt solutions (c) doc bMethod 1b

Dilute sulphuric acid (= acidified water) gives hydrogen and oxygen gases (2 : 1 volume ratio) via equations 3 and 8 This is one way of showing water is a compound i.e. by splitting into two gaseous elements.

Molten anhydrous calcium chloride gives calcium (+) and chlorine (-), equations 13 and 2.

Chemical Tests for some of the Gases formed by these electrolysis experiments.

  • hydrogen - colourless gas that gives a squeaky pop when ignited with a lit splint.
  • oxygen - colourless gas that relights a glowing splint.
  • chlorine - green gas that bleaches damp litmus paper. Best done with blue litmus which changes to red before being bleached.

For more details and tests see the qualitative chemical analysis page

Concentrated hydrochloric acid gives equal volumes of hydrogen (-) and chlorine (+) via equations 2 and 3. However in very dilute solution, equation 8 occurs too, giving oxygen as well as chlorine.    
For most sulphate salts of reactive metals e.g. sodium/magnesium sulphate the electrolysis products of the aqueous salt solution are hydrogen at the negative (-) cathode electrode (Eq 3) and oxygen at the positive (+) anode (Eq 8) with inert electrodes such as carbon or platinum.    

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2b. Summary of ELECTRODE REACTIONS - half-cell electrode equations






(-) negative cathode electrode where reduction of the attracted positive cations is by electron gain to form metal atoms or hydrogen [from Mn+ or H+, n = numerical positive charge]. The electrons come from the positive anode (see below).

(+) positive anode electrode where the oxidation of the atom or anion is by electron loss. Non-metallic negative anions are attracted and may be oxidised to the free element. Metal atoms of a metal electrode can also be oxidised to form positive metal ions which pass into the liquid electrolyte. The released electrons move round in the external part of the circuit to produce the negative charge on the cathode electrode.

So, before each electrode equation is a (-) for a negative cathode electrode = a reduction reaction equation or a (+) for a positive anode electrode = an oxidation reaction equation

The electrode equations are shown on the left with examples of industrial processes where this electrode reaction happens below on the right. Unless otherwise stated, the electrodes are inert i.e. they do not chemically change e.g. platinum or carbon-graphite.

 PLEASE NOTE - all electrode equations are a summary-simplification of what happens on an electrode surface in electrolysis. There may be e.g. two equations which are totally equivalent to each other to describe WHAT IS ACTUALLY FORMED e.g. the formation of hydrogen or oxygen and in some cases other products may be formed too.


a reduction electrode reaction

(-) Na+(l) + e- ==> Na(l) (sodium metal)

sodium ion reduced to sodium metal atoms: typical of electrolysis of molten chloride salts to make chlorine and the metal

an oxidation electrode reaction

(+) 2Cl-(l/aq) - 2e- ==> Cl2(g)

 or  2Cl- ==> Cl2 + 2e-

Note that you can write these anode oxidation reactions either way round

chloride ion oxidised to chlorine gas molecules: electrolysis of molten chloride salts(l) or their concentrated aqueous solution(aq) or conc. hydrochloric acid(aq) to make chlorine

a reduction electrode reaction

(-) 2H+(aq) + 2e- ==> H2(g) (hydrogen gas)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

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

All three equations amount to the same overall change i.e. the formation of hydrogen gas molecules and as far as I know any is acceptable in an exam?

top sub-indexhydrogen ion or water reduced to hydrogen gas molecules: electrolysis of many salt or acid solutions to make hydrogen

a reduction electrode reaction

(-) Cu2+(aq) + 2e- ==> Cu(s) (copper deposit)

copper(II) ion reduced to copper atoms: deposition of copper in its electrolytic purification or electroplating using copper(II) sulphate solution, electrode can be copper or other metal to be plated

an oxidation electrode reaction

(+) Cu(s) - 2e- ==> Cu2+(aq) (copper dissolves)

or  Cu(s) ==> Cu(s) + 2e-

copper atoms oxidised to copper(II) ions: dissolving of copper in its electrolytic purification or electroplating (must have positive copper anode)

a reduction electrode reaction

(-) Al3+(l) + 3e- ==> Al(l) (aluminium)

aluminium ions reduced to aluminium atoms: extraction of aluminium in the electrolysis of its molten oxide ore(l) 

an oxidation electrode reaction

(+) 2O2-(l) - 4e- ==> O2(g) (oxygen gas)

or  2O2-(l)  ==> O2(g) + 4e-

oxide ion oxidised to oxygen gas molecules: electrolysis of molten oxides e.g. anode reaction in the extraction of aluminium from molten bauxite.

an oxidation electrode reaction

(i) (+) 4OH-(aq) - 4e- ==> 2H2O(l) + O2(g) (oxygen gas)

or  4OH-(aq) ==> 2H2O(l) + O2(g) + 4e-

(ii) (+) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g) (oxygen gas)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e-

Both equations amount to the same overall change i.e. the formation of oxygen gas molecules and as far as I know either is acceptable in an exam?

There are two equations that describe the formation of oxygen in the electrolysis of water.

hydroxide ions or water molecules are oxidised to oxygen gas molecules: electrolysis of many salt solutions such as sulphates, sulphuric acid etc. gives oxygen (chlorides ==> chlorine in concentrated solution, but can also give oxygen in diluted solution)


 a reduction electrode reaction

(-) Pb2+(l) + 2e- ==> Pb(l) (lead deposit)

lead(II) ions reduced to lead atoms: electrolysis of molten lead(II) bromide(l) 

an oxidation electrode reaction

(+) 2Br-(l/aq) - 2e- ==> Br2(g/l) (bromine)

or  2Br- ==> Br2 + 2e-

bromide ions oxidised to gas/liquid bromine molecules: electrolysis of molten bromide salts(l) or their concentrated aqueous solution(aq) or conc. hydrobromic acid(aq) to make bromine

a reduction electrode reaction

(-) Zn2+(aq) + 2e- ==> Zn(s) (zinc deposit)

zinc ions reduced to zinc atoms: galvanising steel (the electrode) by electroplating from aqueous zinc sulphate solution, (or from molten zinc chloride?)

a reduction electrode reaction

(-) Ag+(aq) + e- ==> Ag(s) (silver deposit)

silver ions reduced to silver atoms: silver electroplating from silver salt solution(aq), electrode can be other metal

a reduction electrode reaction

(-) Ca2+(l) + 2e- ==> Ca(s) (calcium metal)

calcium ions reduced to calcium atoms e.g. in molten calcium chloride or bromide etc.

Electrolysis calculations (c) doc b(c) doc bsection 13. of the (c) doc b(c) doc bChemical Calculations pages index

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3a. Summary of the products of electrolysis of various electrolytes

What are the products of the electrolysis of molten aluminium oxide, aqueous copper sulphate solution, aqueous sodium chloride solution (brine), hydrochloric acid, sulphuric acid, molten lead(II) bromide, molten calcium chloride?

Electrolyte cathode product cathode equation anode product anode equation comments
molten aluminium oxide


molten aluminium Al3+(l) + 3e- ==> Al(l) oxygen gas

2O2-(l) - 4e- ==> O2(g)

or  2O2-(l)  ==> O2(g) + 4e-

Method 2 in principle: The industrial method for the extraction of aluminium its ore
aqueous copper(II) sulfate


copper deposit any conducting electrode e.g. carbon rod, any metal including copper itself

Cu2+(aq) + 2e- ==> Cu(s)

oxygen gas inert electrode like carbon (graphite rod) or platinum

(i) 4OH-(aq) - 4e- ==> 2H2O(l) + O2(g)

or  4OH-(aq) ==> 2H2O(l) + O2(g) + 4e-

(ii) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e-

Method 1a or Adapt Method 3 - e.g. use carbon rods: The blue colour of the copper ion will fade as the copper ions are converted to the copper deposit on the cathode
aqueous copper (II) sulphate


copper deposit any conducting electrode e.g. carbon rod, any metal including copper itself

Cu2+(aq) + 2e- ==> Cu(s)

copper(II) ions - the copper anode dissolves copper anode

Cu(s) - 2e- ==> Cu2+(aq)

or  Cu(s) ==> Cu(s) + 2e-

See Method 3: This is the basis of the method of electroplating any conducting solid with a layer of copper. When using both copper cathode and anode, the blue colour of the copper ion does not decrease because copper deposited at the (-) cathode = the copper dissolving at the (+) anode.
molten sodium chloride


molten sodium Na+(l) + e- ==> Na(l) chlorine gas

2Cl-(l) - 2e- ==> Cl2(g)

 or  2Cl-(l) ==> Cl2(g) + 2e-

Method 2 in principle: This a method used to manufacture sodium and chlorine. See Down's Cell
aqueous sodium chloride solution (brine)



2H+(aq) + 2e- ==> H2(g)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

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

chlorine gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2(g) + 2e-

Method 1a: This is the process by which hydrogen, chlorine and sodium hydroxide are manufactured
hydrochloric acid


hydrogen gas 2H+(aq) + 2e- ==> H2(g)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

For theoretical calculations of gas volumes formed see Section 14 of Chemical Calculations

chlorine gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2 + 2e- 

Method 1a and method 1b: All acids give hydrogen at the cathode.

Theoretically the gas volume ratio is H2:Cl2 is 1:1, BUT, chlorine is slightly so there seems less chlorine formed than actually was.

sulphuric acid

sulfuric acid


hydrogen gas 2H+(aq) + 2e- ==> H2(g)

For theoretical calculations of gas volumes formed see Section 14 of Chemical Calculations

oxygen gas

(i) 4OH-(aq) - 4e- ==> 2H2O(l) + O2(g)

or  4OH-(aq) ==> 2H2O(l) + O2(g) + 4e-

(ii) (+) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e- 

Method 1a and method 1b: All acids give hydrogen at the cathode. Whereas hydrochloric acid gives chlorine at the anode, the sulfate ion does nothing and instead oxygen is formed. This is the classic 'electrolysis of water'.

Theoretically the gas volume ratio is H2:O2 is 2:1 which you see with the Hofmann Voltameter

molten lead(II) bromide


molten lead Pb2+(l) + 2e- ==> Pb(l) bromine vapour

2Br-(l) - 2e- ==> Br2(g)

 or  2Br-(l) ==> Br2(g) + 2e-

Method 2: A good demonstration in the school laboratory - brown vapour and silvery lump provide good evidence of what's happened
molten calcium chloride


solid calcium Ca2+(l) + 2e- ==> Ca(s) chlorine gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2(g) + 2e-

Method 2 in principle: The basis of the industrial method for the manufacture of calcium metal
************** ************ **************************************************** ********** ******************************************************* *************************************

3b. Summary of Industrial Processes using Electrolysis

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A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc b4. Simple Cells or batteries

  • In electrolysis, electrical energy is taken in (endothermic) to enforce the oxidation and reduction to produce the products.

  • The chemistry of simple voltaic cells or batteries is in principle the opposite of electrolysis.

  • A redox reaction occurs to produce products and energy is given out because it is an exothermic reaction, BUT the energy is released as electrical energy NOT heat energy so the system shouldn't heat up.

  • A simple cell can be made by dipping two different pieces of metal (of different reactivity) into a solution of ions e.g. a salt or dilute acid. If you use the same metal for both strips, they 'cancel' each other out, so no potential difference (voltage) so no current of electrical energy.

  • All you need is a solution of charged positive and negative particles called ions e.g. sodium Na+, chloride Cl-, hydrogen H+, sulphate SO42- etc.

  • The greater the difference in reactivity, the bigger the voltage produced. However this is not a satisfactory 'battery' for producing even a small continuous current.

  • BUT a simple demonstration cell can be made by dipping strips of magnesium and copper into an a salt solution and connecting them via a voltmeter (e.g. as in diagram) and a voltage is readily recorded.

    • The electrode reactions are:

      • at the (+) electrode 2H+(aq) + 2e- ==> H2(g) (hydrogen ions reduced)

        • here the copper is inert and the hydrogen ions come from water.

      • at the (-) electrode Mg(s) - 2e- ==> Mg2+(aq) (magnesium atoms oxidised)

        • the magnesium dissolves into solution by a chemical reaction

      • overall the redox reaction is 2H+(aq) + Mg(s) ==> Mg2+(aq) + H2(g) 

      • and the electrons from the oxidation of the magnesium move round through the magnesium strip, along the external wire to the copper electrode.

      • top sub-indexNote the (+) and (-) polarity of the electrodes in a cell, is the opposite of electrolysis because the process is operating in the opposite direction i.e.

        • in electrolysis electrical energy induces chemical changes,

        • but in a cell, chemical changes produce electricity.

  • One of the first practical batteries is called the 'Daniel cell' which is illustrated below.

    • The Daniel Cell

    • This 'voltaic 'or galvanic' electrochemical cell uses a half-cell of copper dipped in copper(II) sulphate,

    • and in electrical contact with a 2nd half-cell of zinc dipped in zinc sulphate solution.

    • The zinc is the more reactive, and is the negative electrode, releasing electrons because

      • on it zinc atoms lose electrons to form zinc ions, Zn(s) ==> Zn2+(aq) + 2e-

    • The less reactive metal copper, is the positive electrode, and gains electrons from the negative electrode through the external wire connection and here ..

      • the copper(II) ions are reduced to copper atoms, Cu2+(aq) + 2e- ==> Cu(s)

    • Overall the reactions is: Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s)

      • or ionically: Zn(s) + Cu2+(aq) ==> Zn2+(aq) + Cu(s)

    • The overall reaction is therefore the same as displacement reaction, and it is a redox reaction involving electron transfer and the movement of the electrons through the external wire to the bulb or voltmeter etc. forms the working electric current.

  • A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc bThe cell voltage can be predicted by subtracting the less positive voltage from the more positive voltage:

    • e.g. referring to the list of electrode potentials on the right and choosing two different metals coupled together in the electrolyte solution ...

    • a magnesium and copper cell will produce a voltage of (+0.34) - (-2.35) = 2.69 Volts

    • or an iron and tin cell will only produce a voltage of (-0.15) - (-0.45) = 0.30 Volts.

    • Note (i) the bigger the difference in reactivity, the bigger the cell voltage produced

    • and (ii) the 'half-cell' voltages quoted in the diagram are measured against the H+(aq)/H2(g) system which is given the standard potential of zero volts.

  • Cells or batteries are useful and convenient portable sources of energy for torches, radios, shavers and other gadgets BUT they are expensive compared to what you pay for 'mains' electricity. On the other hand you have no choice for a car battery!

  • (c) doc b Electrolysis and cell theory for Advanced Level Chemistry Students

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5. Fuel Cells - another sort of battery

  • 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) 

    • It is a non-polluting clean fuel since the only combustion product is water. 

    • It would be ideal if it 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. see (c) doc b GCSE/IGCSE-AS notes on fuels

    •  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!

  • Fuel cells are 'battery systems' in which two reactants can be continuously fed in. The consequent redox chemistry produces a working current.

  • 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.

  • 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.

  • DIAGRAM and CHEMISTRY below: A fuel cell 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 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 migrate through the electrolyte and pass through the semi-permeable membrane to the cathode, where they reunite with oxygen and the electrons to produce water.

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

  • Note on reverse action:

    • 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 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 when in the 'darkness of night'.

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

















equation Description of a hydrogen-oxygen fuel cell

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.

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.

keywords help with coursework projects electrode equations: 1. Introduction to electrolysis * 2. Summary of practical experimental arrangements for laboratory (lab) apparatus electrolysis experiments and electrode reactions * Summary of the products of selected electrolysis processes* 3a. Summary of Industrial Electrolysis Process links * 4. Simple cells or batteries * electrode equations anode products cathode products negative electrode products positive electrode products 5. Fuel cells * Na+ + e- ==> Na * 2Cl- - 2e- ==> Cl2 * 2Cl- ==> Cl2 + 2e- * (-) 2H+(aq) + 2e- ==> H2(g) * 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l) * 2H2O(l) + 2e- ==> H2(g) + 2OH-(aq) * (-) Cu2+(aq) + 2e- ==> Cu(s) (copper deposit) (+) Cu(s) - 2e- ==> Cu2+(aq) (copper dissolves) (-) Al3+(l) + 3e- ==> Al(l) (aluminium) (+) 2O2-(l) - 4e- ==> O2(g) (oxygen gas) (+) 4OH-(aq) - 4e- ==> 2H2O(l) + O2(g) (oxygen gas) or (+) 2H2O(l) - 4e- ==> 4H+(l) + O2(g) (oxygen gas) Both equations amount to the same overall change i.e. the formation of hydrogen gas molecules and as far as I know either is acceptable in an exam?  (-) Pb2+(l) + 2e- ==> Pb(l) (lead deposit) (+) 2Br-(l/aq) - 2e- ==> Br2(g/l) (bromine) (-) Zn2+(aq) + 2e- ==> Zn(s) (zinc deposit) (-) Ag+(aq) + e- ==> Ag(s) (silver deposit) (-) Ca2+(l) + 2e- ==> Ca(s) (calcium metal) Na+ + e- ==> Na * 2Cl- - 2e- ==> Cl2 * 2Cl- ==> Cl2 + 2e- * (-) 2H+ + 2e- ==> H2 * 2H3O+ + 2e- ==> H2 + 2H2O * 2H2O + 2e- ==> H2 + 2OH- * (-) Cu2+ + 2e- ==> Cu (copper deposit) (+) Cu - 2e- ==> Cu2+ (copper dissolves) (-) Al3+ + 3e- ==> Al (aluminium) (+) 2O2- - 4e- ==> O2 (oxygen gas) (+) 4OH- - 4e- ==> 2H2O + O2 (oxygen gas) or (+) 2H2O - 4e- ==> 4H+ + O2 (oxygen gas) Both equations amount to the same overall change i.e. the formation of hydrogen gas molecules and as far as I know either is acceptable in an exam?  (-) Pb2+ + 2e- ==> Pb (lead deposit) (+) 2Br-(l/aq) - 2e- ==> Br2(g/l) (bromine) (-) Zn2+ + 2e- ==> Zn (zinc deposit) (-) Ag+ + e- ==> Ag (silver deposit) (-) Ca2+ + 2e- ==> Ca (calcium metal)

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