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Extra Electrochemistry Information Notes on Electrolysis, batteries and fuel cells

Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  (revise courses equal to US grades 9-10)

Index for this page 1. Introduction to electrolysis * 2. Summary of practical experimental arrangements for laboratory (lab) electrolysis experiments and electrode reactions * 3. Summary of Industrial Electrolysis Process links * 4. Simple cells or batteries * 5. Fuel cells

Associated LINKS to other pages on this site: full LIST of GCSE-KS4 Chemistry, Earth Science and Radioactivity REVISION NOTES * Metal Extraction * Industrial chemistry * Metal reactivity (redox introduction) * Types of chemical reaction * Electrolysis calculations * EMAIL query?comment

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.

• SUMMARY OF COMMON ELECTRICAL CONDUCTORS and what makes up the circuit?
  • 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)
    • Ions 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.

A simple experiment to show the movement of coloured ions

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, CrO₄^2--, 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!).

• 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 or melt of ions 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.
• ELECTROLYSIS 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 -ve electrode,
  • and negatively charged ions move to the positive electrode (anions attracted to anode), e.g. in the diagram, chloride ions Cl^-, move to the +ve electrode. 
• 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
  • 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
  • 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.

Eq.

no.

(-) 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.

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.

(-) Na⁺_(l) + e^- ==> Na_(l) (sodium metal)

(+) 2Cl^-_(l/aq) - 2e^- ==> Cl_2(g)

(-) 2H⁺_(aq) + 2e^- ==> H_2(g) (hydrogen gas)

or 2H₃O⁺_(aq) + 2e^- ==> H_2(g) + 2H₂O_(l)

or 2H₂O_(l) + 2e^- ==> H_2(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?

(-) Cu²⁺_(aq) + 2e^- ==> Cu_(s) (copper deposit)

(+) Cu_(s) - 2e^- ==> Cu²⁺_(aq) (copper dissolves)

(-) Al³⁺_(l) + 3e^- ==> Al_(l) (aluminium)

(+) 2O^2-_(l) - 4e^- ==> O_2(g) (oxygen gas)

(+) 4OH^-_(aq) - 4e^- ==> 2H₂O_(l) + O_2(g) (oxygen gas)

or (+) 2H₂O_(l) - 4e^- ==> 4H⁺_(l) + O_2(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^- ==> Br_2(g/l) (bromine)

(-) Zn²⁺_(aq) + 2e^- ==> Zn_(s) (zinc deposit)

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

(-) Ca²⁺_(l) + 2e^- ==> Ca_(s) (calcium metal)

Electrolysis calculations section 13. of the Chemical Calculations pages

3. Summary of Industrial Processes using Electrolysis

• Below is a list of links to other web pages on this site with sections on electrolysis.

• The extraction of aluminium from purified molten bauxite ore.

• Anodising aluminium to strengthen the protective oxide layer.

• The extraction of sodium from molten sodium chloride using the 'Down's Cell'.

• The purification of copper using copper electrodes.

• The purification of zinc.

• Making chlorine, hydrogen and sodium hydroxide from sodium chloride salt solution by brine electrolysis and also an electrolysis Q on the Halogens task sheet.

• Electroplating conducting surfaces to coat it with a metal layer.

4. 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 SO₄^2- 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^- ==> H_{2 (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) + H_{2 (g)} 

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

    • Note 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.

  • 

  • 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) + CuSO_4(aq) ==> ZnSO_4(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.

• The cell voltage can be predicted by subtracting the less positive voltage from the more positive voltage:

  • e.g. 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)/H_2(g) system which is given the standard potential of zero volts.

• Cells or batteries are useful and convenient portable sources of energy but they are expensive compared to what you pay for 'mains' electricity.

5. Fuel Cells - another sort of battery

• Hydrogen gas can be used as fuel.

  • It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water.

    • hydrogen + oxygen ==> water or 2H_2(g) + O_2(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.

  • Hydrogen can be used to power fuel cells. (GCSE-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'.

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

* Note the +ve and -ve electrode charges are reversed compared to electrolysis, because the system is operating in the opposite direction.

Other notes on ADVANCED chemistry pages: The alkaline hydrogen-oxygen fuel cell is described in Equilibria Part 7 Redox Chemistry  and organic fuel cells are described in Redox Chemistry Part 3.