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INTRODUCTION TO ELECTROLYSIS - experiments, apparatus, electrode equations explained



(Suitable for AQA, Edexcel and OCR GCSE chemistry students)

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

1. Introduction to electrolysis – apparatus, electrolytes and non–electrolytes

Sub-index for this page

(a) A summary of common electrical conductors

(b) Apparatus for investigations - the electrolysis cell

(c) What happens in electrolysis? and what products can be formed

(d) A simple experiment to show the movement of coloured ions

(e) Examples of diagrams to explain electrolysis and product formation on the electrodes

(f) Summary of the criteria for electrolysis to take place and the splitting up of a compound

This page introduces you to the principles of electrolysis and explains all the technical terms you need to know about e.g. ions, electrolyte, non–electrolyte, cell, electrodes, electrode equations, negative cathode electrode, positive anode electrode. All the technical terms are defined and explained. The apparatus you can use are described and seven examples of electrolysis processes are described in detail (links 2. to 7. in the index at the end of each page). These revision notes on an introduction to electrolysis 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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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 or producing gases like hydrogen, oxygen and chlorine from salt solutions.

(a) A summary of common electrical conductors

  • What makes up a circuit in cells – batteries or electrolysis? What carries the current?
  • Conductors are materials (solid or liquid/solution) that 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), non–metal carbon (graphite) and some recently developed 'smart materials' (which get no further mention in this section of notes).
    • This conduction involves the movement of free or delocalised electrons (e charged particles) and does not involve any chemical change.
      • That doesn't mean to say these 'electrical currents' can't promote chemical change, which happens as we will see later in the process we call electrolysis.
  • Any compound, molten or dissolved in solution, in which the liquid contains free moving ions is called an electrolyte and can conduct an electrical current. (see non–electrolyte)
    • Ions are electrically 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 is a solution containing mobile ions that conduct electrical energy as a stream of moving charged particles.
    • An electrolyte may consist of ...
      • (i) a molten ionic compound i.e. on melting the ions are free to move to carry the current, or
      • (ii) any compound that dissolves in a liquid to give a solution of ions that are free to move.
        • The compound is usually ionic and the liquid is usually water, so in most of the examples described, the electrolyte is an aqueous solution of ions with a few molten salt examples.
        • Note: Aqueous means to do with water, so an aqueous solution is a solution made by dissolving something in the solvent water.
    • When an electric current is passed through such an electrolyte chemical changes can occur on the electrical contacts (called electrodes) and chemical changes happen to break down the compound in a process called electrolysis.
    • Water is very poor conductor because it is a covalent compound and only minute amounts of it ionises to form hydrogen and hydroxide ions, so water it is not an effective electrolyte.
      • The majority of liquid water consists of covalent H2O molecules, but there are trace quantities of H+ and OH ions from the self–ionisation of water,

      • H2O(l) H+(aq) + OH(aq)

        • (about 1 in 200 million does this!, the reaction is reversible, so the longer half-arrow to the left tells you that most water remains as water molecules!).

        • However, once you dissolve an ionic salt like compound or a stronger acid, water (as an aqueous solution) becomes a good electrical conductor i.e. a good electrolyte.

  • What doesn't conduct and why? why can't solid ionic compounds not conduct electricity?
    • Apart from metals, most solids do not conduct electricity because there are no free electrons or ions to carry an electric current.
    • However, although molten ionic compounds and solutions in of ionic compounds in water can carry a current (and undergo electrolysis - chemical change) solid ionic compounds cannot conduct electricity and undergo electrolysis. This is because the ions are tightly held in the crystal lattice and cannot move around and migrate to any electrical contact placed on the solid.


(b) Apparatus for investigations - the electrolysis cell

  • When an appropriate d.c. current is passed through an electrolyte, chemical changes occur where the external circuit connections (electrodes) are dipped into the electrolyte.
    • These chemical changes ONLY occur on the surface of the electrodes where they are in contact with the electrolyte solution and (usually) elements are released as the compound is broken down by the process called electrolysis.
    • What does a simple, but complete, electrical circuit for electrolysis consist of?
      • A simple cell system for electrolysis
      • The electrolyte container can be made from a short piece of wide glass/plastic tubing with a rubber bung base with two holes drilled in it to take to carbon rods as electrodes.
      • Small test tubes filled with electrolyte are inverted over the carbon electrodes.
      • This even simpler set-up is recommended by the RSC and consists of two wire electrodes bent in a S shape so the gases can be collected in little test tubes filled with the electrolyte and inverted over the electrodes in the beaker of electrolyte.
      • The diagrams above illustrate simple electrolysis experiments you will see or (hopefully) do in a school laboratory or college laboratory.
      • The electrolyte solution (in this case sulfuric acid, can be sodium chloride etc.) is contained within the electrolysis cell (e.g. section of wide plastic piping).
      • Two electrical connectors called electrodes (e.g. graphite/carbon rods) protrude upwards into the electrolyte solution pushed through two holes drilled in a larger rubber bung. This is the same function as the two wires in the other simpler electrolysis set-up illustrated above.
      • The circuit is completed when connected to an external electrical power supply of d.c. current, and usually a voltage of 2-3 V is quite sufficient to give a good rate of electrolysis.
      • So, in sequence from the negative terminal, through the external copper wire electrons flow clockwise from the positive electrode to the negative electrode (cathode), then ions (NOT electrons) carry the current through the electrolyte across to the positive electrode (anode), and then electrons again carry the current through another external wire to the positive terminal of the battery or other power supply.
      • When you switch on the d.c. power on, or connect the battery, the electrolysis process should start.  Often, but not always, you will see bubbles of gas appearing on the electrode surface, because that's where the chemical changes we call electrolysis take place!
      • Flowing in one direction only, the electrons carry the current in the external copper wires BUT not in the electrolyte solution. However in the electrolyte solution there are two ion currents flowing in opposite directions, and it is important that this is understood because no chemical change can take place if the ions are not attracted to their oppositely charged electrode.


(c) What happens in electrolysis? and what products can be formed?

  • All salts and many acids are good electrolytes, good electrical conductors of a d.c. current because they provide high concentrations of positive and negative ions.
  • Positive ions (cations) e.g. hydrogen H+, copper Cu2+, sodium Na+ are attracted to the negative electrode (cathode).
    • Positive cations migrate to the negative cathode.
  • Negative ions (anions) e.g. chloride Cl , sulfate SO42–, bromide Br, are attracted to the positive electrode (anode).
    • Negative anions migrate to the positive anode.
  • It is possible to demonstrate this flow using a coloured ion experiment (see diagram and text below).
  • Remember no electrons flow in the solution, but they do flow in the external metal wires or carbon (graphite)/metal electrodes of the external circuit.
  • So, what is the chemistry of electrolysis?
  • When an ion meets its oppositely charged electrode, one of two things can happen. Either the ion hangs around the electrode and does nothing OR the ion undergoes chemical change, sometimes referred to as 'the ion is discharged'.
    • The chemical changes that occur on the surface of an electrode are either a REDUCTION (on the negative cathode electrode) or an OXIDATION (on the positive anode electrode).
    • Each of the oxidation or reduction changes is written as a half-equation, so you see the electrons lost or gained
    • At the negative cathode electrode, positive ions (cations) are attracted and these positive ions may gain electrons and are reduced to some chemical product e.g. typical half-reactions ...
    • Either hydrogen gas or a metal deposit is formed on the negative cathode electrode.
    • 2H+(aq) + 2e ==> H2(g)    (colourless hydrogen gas from acid solutions)
    • Cu2+(aq) + 2e ==> Cu(s)    (copper deposit from copper sulfate solution)
    • Pb2+(l) + 2e ==> Pb(l)   (lead formed from a hot molten salt)
    • Al3+(l) + 3e ==> Al(l)   (aluminium formed from molten oxide)
    • Note that ...
    • (i) reduction = electron gain, at the negative cathode electrode,
    • (ii) hydrogen and metals are formed at the negative cathode electrode,
    • (iii) not all the positive ions will be discharged i.e. reduced, so in a mixture of H+ and Na+ ions in aqueous solution, the hydrogen ions are preferentially reduced to hydrogen, leaving the sodium ions unchanged in solution,
    • (c) doc bAND generally speaking, the less reactive a metal, the more easily its ion is reduced to the metal on the electrode surface e.g. in a mixture of positive ions the preference order is
      • Cu2+ (==> Cu) > H+ (==> H2) > Na+ (==> Na)
    • A general rule with reference to the reactivity series of metals:
      • If the metal in the salt solution is more reactive than hydrogen, then hydrogen the hydrogen ion is most likely to be discharged at the negative cathode giving hydrogen.
        • If a reactive metal like sodium was discharged, it would immediately react with water giving hydrogen anyway!
      • If the metal in the salt solution is less reactive than hydrogen it is the metal ion that is likely to be discharged forming a deposit of the metal on the electrode surface.
        • In practice you can get a deposit of lead from a lead nitrate solution.
    • At the positive anode electrode, negative ions (anions) are attracted and these negative ions may lose electrons and are oxidised to some chemical product e.g. typical half-reactions ...
      • A non-metal like oxygen or chlorine is released at the positive anode electrode.
      • 2Cl(aq) – 2e ==> Cl2(g)   (chloride ion ==> pale green chlorine gas from NaCl)
      • 2Br(l) – 2e ==> Br2(g)   (bromide ion ==> brown bromine vapour from molten bromide)
      • 2O2–(l) – 4e ==> O2(g)  (oxide ion ==> colourless oxygen gas from a molten oxide)
      • 4OH(aq) – 4e ==> 2H2O(l) + O2(g) (hydroxide ==> colourless oxygen gas, from alkaline hydroxide solution or traces of OH- from water)
        • or     4OH(aq) ==> 2H2O(l) + O2(g) + 4e
      • Note that ...
      • (i) oxidation = electron loss, at the positive anode electrode,
      • (ii) apart from hydrogen, other non–metals are formed at the positive anode electrode,
      • (iii) AND not all the negative ions will be discharged i.e. oxidised, so in a mixture of OH and Cl ions in aqueous solution, the chloride ions are preferentially oxidised to chlorine, leaving most of the hydroxide ions unchanged in solution.
      • Oxygen is usually produced unless a halide ion is present,
        • so chloride would give chlorine, bromide gives bromine and iodide gives iodine at the positive anode.
    • All the above electrode equations showing the formation of the electrolysis products are fully described in the context of a laboratory experiment or industrial process.
    • Three 'connected' sub–notes on electrolysis investigations:
      • As well as investigating the products of electrolysis, you can also vary experimental conditions e.g. changes in voltage p.d. or electrolyte concentration can be studied. Possible investigations can show ....
      • (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 greater the voltage, the faster the rate of electrolysis, but don't over do it!
  • 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 molecule liquids or solutions of covalent compounds. Even if a covalent compound dissolves in water, if no ions are formed, there will be no electrical conduction.
  • 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.


(d) 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 filter paper and the current switched on. (* Using a mixture of copper(II) sulfate and potassium chromate(VI).

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!).

In the electrolyte, as with any solution, ALL the particles are moving around at random, but they are not distributed at random as the electrolysis proceeds.

During electrolysis positive ions on average move towards the negative cathode electrode,

AND the balancing negative ions on average move towards the positive anode electrode.

Both streams of moving ions constitute the electric current flowing in the electrolyte.


(e) Examples of diagrams to explain electrolysis and the formation of products on the electrodes

Two examples are illustrated below, but the full electrolysis description and explanation is given on individual web pages - see Electrochemistry INDEX.

Diagram showing the net direction of ion movement in water acidified with sulfuric acid.


Diagram showing the net direction of ion movement in molten sodium chloride


(f) Summary of the criteria for electrolysis to take place and the splitting of a compound
  • 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!
    • The particle theory of solids still applies even if you try to pass electricity through it (apart from graphite and metals).
    • 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 molten ionic 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 and is used extensively in the chemical industry for extracting elements like sodium and chlorine from their naturally occurring compounds.
    • Since it requires an 'input' of energy, it is an endothermic process and costly to pay for the electrical energy.
  • 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). This is an example of how electrolysis is used in the chemical industry.
  • 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 maybe have to derive an overall equation for the electrolysis.
      • e.g. for the electrolysis of molten sodium chloride described above, the overall chemical change due to electrolysis can be written as ...
        • 2NaCl(l) ==> 2Na(l) + Cl2(g)
      • At this point it is appropriate and very important to mention and use state symbols in all electrode equations and overall chemical change equations.
        • Reminder: (g) = gas, (l) = liquid, (s) = solid, (aq) = aqueous solution in water
        • e.g. for sodium chloride,   NaCl(l)   for the molten salt,
        • NaCl(aq)   or   Na+ + Cl   for two possible expressions of the aqueous solution
  • When dealing with the electrolysis of aqueous solutions of salts and acids in water, things can be more complicated and sometimes several competing electrode reactions can occur at the same time, and sometimes products differ depending on the nature of the cathode and anode electrodes.
    • There are links from the electrochemistry index page which describe in great detail particular examples of the process of electrolysis and all can be done as college or school student experiments or teacher demonstrations, your pupils can have great fun with these experiments but take great care with the production of chlorine!

Electrolysis Quiz (GCSE 9-1 HT Level (harder)

Electrolysis Quiz (GCSE 9-1 FT Level (easier)


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