Topic 15 Electrochemistry, electrochemical cells, electrolytic cells
Updated 2008-08-19 R
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites

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Table of contents
15.00 Electrolytes
3.59.0 Electrical conductivity
15.1.0 Electroplating
15.2.0 Oxidation and reduction, redox reactions
15.3.0 Rusting, corrosion
15.4.0 Electrical conductivity of a substance
15.5.0 Electrolysis
15.6.0 Electrochemical cells
15.7.0 Electrode potential of metals

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15.00 Electrolytes
15.01 Conductivity of solutions of different electrolytes
15.02 Strong electrolytes
15.03 Identify lead ions in an unknown solution
15.04 Weak electrolytes

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3.59.0 Electrical conductivity
3.59 Electrical conductivity of solids
3.59.1 Electrical conductivity of melted solids, fused solids
3.60 Electrical conductivity of liquids
6.38 Electricity conductors (Primary)

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15.1.0 Electroplating
15.1.1 Faraday's first law
15.1.1.1 Test Faraday's first law with copper and copper (II) sulfate solution
15.1.1.2 Test Faraday's first law with other metals
15.1.2 Electroplating, copper plating
15.1.3 Electroplating, chromium plating
15.1.4 Electroplate, nickel plating
15.1.5 Electroplating, silver plating
15.1.6 Electroplating, zinc plating of copper
15.1.7 Electroforming with copper
15.1.8 Anodize aluminium
15.1.9 Silvering and desilvering, plating and deplating silver

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15.2.0 Oxidation and reduction, redox reactions
15.2.1 Oxygen as an oxidizing agent
15.2.2 Chlorine as an oxidizing agent
15.2.2.1 Bromine as an oxidizing agent
15.2.3 Potassium dichromate as an oxidizing agent
12.05.6 Prepare potassium dichromate, K₂Cr₂O₇
15.2.4.1 Potassium manganate (VII) oxidizes iron (II) to iron (III)
15.2.4.2 Potassium manganate (VII) oxidizes glycerol to carbon dioxide and water
15.2.4.3 Potassium manganate (VII) solution liberates chlorine from hydrochloric acid
12.08a.1 Reactions of manganese (II) salts, Mn
3.76 Reduction of potassium permanganate with sulfur dioxide
12.08a.2 Prepare manganates, MnO₄^2-
12.08a.3 Prepare potassium permanganate, KMnO₄
12.08a.4 Reactions of permanganate ion, MnO₄^-
15.2.5 Concentrated nitric acid as an oxidizing agent
15.2.5.1 Nitrous acid as an oxidizing agent or a reducing agent
15.2.6 Sulfuric acid as an oxidizing agent
15.2.7 Hydrogen peroxide as an oxidizing agent
15.2.8 Tests for oxidizing agents by change in colour of iron (II) to iron (III)
15.2.9 Tests for oxidizing agents by change of colour of iron with copper (II) sulfate
15.2.10 Tests for oxidizing agents by change of colour of zinc with copper (II) sulfate
15.2.11 Breath test for alcohol using potassium dichromate
15.2.12 Breath test for alcohol using a breath analyser ("breathalyser")
15.2.13 Potassium chlorate and potassium persulfate as oxidizing agents
15.2.14 Hydrogen sulfide as a reducing agent
15.2.15 Sulfurous acid as a reducing agent, ionization reaction

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15.3.0 Rusting, corrosion
3.52 Conditions necessary for rusting
3.52.1 The mass of iron and its temperature increases during rusting
3.52.2 Oxygen combines with iron during rusting
3.52.3 Metals can prevent rusting
15.3.1 Rusting of iron wire
15.3.2 Corrosion of magnesium
15.3.3 Rusting of steel wool
15.3.4 Need for oxygen for rusting
15.3.5 Need for oxygen for corrosion of magnesium
15.3.6 Gain in weight of iron when rusting
15.3.8 Oxidation can affect air pressure
15.3.9 Rate of rusting under separates conditions
15.3.10 Rate of rusting of iron wire
15.3.11 Rate of rusting with steel wool
15.3.12 Conditions necessary for rusting
15.3.13 Electrochemical prevention of rusting, cathodic protection
15.3.13.01 Cathodic protection
15.3.13.1 Rate of corrosion affected by formation of electric cells
15.3.14 Restore bronze coins, corrosion of alloys
15.3.15 Corrosion of aluminium
15.3.16 Cleaning tarnished silver
12.06a.1 Reactions of iron (II) and iron (IlI) salts, Fe²⁺, Fe³⁺
12.06a.2 Oxidation of iron (II) salt, Fe²⁺
12.06a.3 Reduction of iron (IlI) salt, Fe³⁺
12.06a.4 Prepare iron (II) oxide, FeO
12.06a.5 Prepare iron (IlI) oxide, Fe₂O₃
12.06a.6 Show that black iron oxide is a mixed base, Fe₃O₄
12.06a.7 Action of heat on hydrated iron chlorides, Fe₂O₃
12.06a.8 Prepare iron (II) ammonium sulfate (NH₄)₂SO₄.FeSO₄.6H₂O

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15.4.0 Electrical conductivity of a substance
2.59 Electrical conductivity of solids
2.59.1 Electrical conductivity of melted solids
2.60 Electrical conductivity of liquids

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15.5.0 Electrolysis
3.68 Electrolysis of melted lead bromide
3.69 Electrolysis of water
3.69.1 Electrolysis of salt solutions
3.69.1.1 Examples of electrolysis with carbon electrodes
3.69.2 Electrolysis of saturated sodium chloride solution
3.69.3 Electrolysis of copper (II) sulfate solution with copper and platinum electrodes
3.69.4 Electrolysis of copper (II) sulfate solution with copper electrodes
3.69.5 Electrolysis of solutions of ionic salts with an overhead projector or microscope, tin (II) chloride, silver nitrate
15.5.2 Electrolysis of fused sodium chloride
15.5.3 Electrolysis of tin (II) chloride solution and silver nitrate solution
15.5.6.1 Electrolysis of water, measure volume of hydrogen generated
15.5.7 Electrolysis of dilute sodium chloride solution
15.5.9 Electrolysis of copper (II) sulfate solution, copper carried across is proportional to current x time, Faraday's laws
15.5.9.1 Microscale electrolysis of an aqueous solution of copper (II) sulfate
15.5.9.2 Measure electrochemical equivalent of copper
15.5.10 Electrolysis of concentrated sodium chloride solution, Nelson cell
15.5.11 Electrolysis of potassium iodide solution, electrolytic writing
15.5.12 Electrolysis of acetic acid solution
33.3.8 Hydrogen / oxygen fuel cell
33.2.9 Ionic migration
33.2.10 Ionic friction

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15.6.0 Electrochemical cells
3.84 Electrical energy from a simple cell, displacement of copper by zinc
3.84.1 Electrochemical cell, voltaic cell, galvanic cell
3.84.2 Test a simple cell with different metals
3.84.3 Test a simple electric cell with copper and zinc in dilute sulfuric acid
3.84.4 Simple galvanic cell, zinc in hydrochloric acid
3.84.5 A voltaic cell with a salt bridge
3.85 Daniell cell
3.86 Electrode potentials of metals
3.87 Lead accumulator cell
3.88 Dry cells, Leclanche cell
3.89 Movement of copper and chromate ions
3.90 Movement of ions between microscope slides, Cu²⁺ ions, CO²⁺ ions
15.6.13 Magnesium / copper battery
15.6.14 Nickel / cadmium battery, NiCad battery

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15.7.0 Electrode potential of metals
15.7.1 Measure potential difference by combining half cells, zinc and iron
15.7.2 Measure potential difference by combining half cells, zinc and copper, zinc and lead
15.7.3 Differences in potential on iron nail

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15.01 Conductivity of solutions of different electrolytes
An electrolyte can conduct an electric current in the fused state, or in solution, and it is decomposed while conducting the current. Electrolytes dissolve in water to produce solutions that conduct electric current. As the concentration of the electrolyte in solution increases, the conductivity of the solution increases. A strong electrolyte breaks up almost entirely when it dissolves to produce an aqueous solution. Water is a very weak electrolyte and a poor conductor of electricity so some electrolyte must be dissolved in it to increase its conductivity.

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15.02 Strong electrolytes
1. Add drops of sodium hydroxide solution to each of three separate solutions of copper (II) sulfate, copper (II) chloride, copper (II) nitrate. Observe the blue precipitate in each case. These solutions contain only the copper (II) ion in common, so assume that this ion was responsible for the formation of the precipitate.

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2. Add drops of barium chloride solution to separate solutions of copper (II) sulfate, sulfuric acid and sodium sulfate. In each case you can attribute the result the presence of the sulfate ion.

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3. Add drops of ferric chloride solution to separate solutions of sodium hydroxide, potassium hydroxide and calcium hydroxide. These experiments with solutions of strong electrolytes suggest that the properties of such solutions are the sum of the properties of the ions present. The properties of a solution of copper (II) sulfate are made up of the properties of the copper (II) ion and the sulfate ion. The copper (II) ion, Cu²⁺, causes the blue green colour of the solution, and is responsible for the formation of many precipitates when other substances are added to the solution. The sulfate ion contributes no colour but forms precipitates with many other ions, such as Ba²⁺, when these are added to copper (II) sulfate solution.

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15.03 Identify lead ions in an unknown solution
Prepare separate solutions of lead nitrate, iron (III) chloride and barium chloride. Test a small portion of each solution in turn with dilute hydrochloric acid, dilute sulfuric acid and sodium hydroxide solution. Tabulate your results. Note that lead nitrate solution always produces a precipitate. Also, iron (III) chloride solution gives a precipitate only when sodium hydroxide solution is added. Barium chloride solution gives a precipitate with both sulfuric acid and sodium hydroxide solutions.

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15.04 Weak electrolytes
Smell very carefully a bottle containing some dilute ammonia solution. The smell of ammonia suggests the presence of ammonia molecules that must have come from the solution. Add a few drops of iron (III) chloride to a little ammonia solution. From the results of a previous experiment with iron (III) chloride, the brown precipitate obtained confirms the presence of the hydroxide ion in ammonia solution. Thus ammonia solution has properties due not only to the ions that are present but also because of ammonia molecules. From similar experiments, you can find that the properties of solutions of weak electrolytes are made up from the properties of the unionized molecules and the properties of the ions produced from them. In such solutions, the ions and molecules are in equilibrium with each other.

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15.1.0 Electroplating
You can pass electric current through an electric cell, called an electroplating bath, to deposit one metal on another. Metal is taken from the anode and deposited on metal articles that act as cathodes. The electrolyte contains the metal to be deposited as ions. The object to be plated is the cathode on an electrolytic cell. A salt of the plating metal is the electrolyte, e.g. chromium salt for chromium plating.

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15.1.1 Faraday's first law
Faraday's first law states that the amount of chemical change during electrolysis is proportional to the charge passed, i.e. the quantity of electricity passed. A coulomb is the quantity of electricity that passes when one amp of current passes for one second. One Faraday (F) = 96,500 coulombs.

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15.1.1.1 Test Faraday's first law with copper and copper (II) sulfate solution
Rinse a strip of thin copper cathode in deionized water. Dry and weigh to the nearest 0.01 g. Use a copper strip as anode. Add 200 g copper (II) sulfate crystals and 80 g concentrated sulfuric acid to 1 g of water. Immerse the electrodes. Pass current at 0.25 amps. Every 15 minutes rinse the cathode with deionized water dry with air, then weigh it. Compare the weights of deposited copper to check whether they agree with Faraday's first law.

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15.1.1.2 Test Faraday's first law with other metals
See diagram 15.01.1
Repeat 15.1.1 1 with other Metals.

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15.1.2 Electroplating, copper plating
See diagram 15.1.2: Copper plating
1. Dip a clean iron nail into copper (II) sulfate solution. The quickly becomes coated with a layer of copper.
2. Clean a key in dilute hydrochloric acid, wash with water and polish with steel wool. The plating bath contains 70 g copper (II) sulfate crystals dissolved in 500 mL of water, 25 mL of methylated spirits, 15 mL concentrated sulfuric acid and grains of clear gelatine. The plating current is 0.5 amps. If the plating current is too strong, a spongy layer forms on the electrode and you can easily rub off the layer of copper.
3. Use a 6 volt battery. Attach a copper electrode to the positive terminal and attach a platinum electrode to the negative terminal. Immerse both electrodes in copper (II) sulfate solution in a beaker. Observe any changes. Copper is plated onto the platinum electrode and is corroded from the copper anode. After a few minutes, reverse the connections and observe any changes. The copper that had plated onto the platinum electrode was corroded from it and plated onto the copper cathode.
4. Coat a non-conductor, e.g. wood, rubber or plastic, with graphite powder. The copper wire in the non-conductor must contact the graphite. Then use it as a cathode to electroplate a non-conductor.

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15.1.3 Electroplating, chromium plating
Steel motor car bumper bars may be chromium plated
Electroplating, copper plating. Add concentrated sulfuric acid to potassium dichromate to form red chromate ion, Cr₂O₇^2-. Decant carefully to remove any solid. Use thin lead as anode and use the object to be plated as a cathode, e.g. a metal spoon. Cover the plating bath with paper to prevent fuming. The plating current is 1 to 20 amps at 50^oC.

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15.1.4 Electroplating, nickel plating
Use thin nickel as an anode. The plating bath is water containing 120 g nickel sulfate crystals, NiSO₄.7H₂O + 15 g ammonium chloride, NH₄Cl + 15 g boric acid, H₃BO₃. The plating current is 0.15 amps for 30 minutes.

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15.1.5 Electroplating, silver plating of copper or nickel
Cutlery may be made of nickel but plated with silver. The letters "EPNS" stamped into some "silver" spoons means "Electroplated Nickel Silver". Use a silver anode and a copper or nickel cathode cleaned in concentrated nitric acid until the surface is matted. The plating current is 1 amp per cm² for 10 minutes. The deposit is dull, but becomes shiny after polishing. The electrolyte, plating bath, contains 1 g silver nitrate in 20 mL water, or 1 litre of water containing 40 g silver nitrate, 560 g potassium iodide and 2.8 mL concentrated sulfuric acid.

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15.1.6 Electroplating, zinc plating of copper
Use a mixture of zinc chloride, boric acid and ammonium chloride. The reaction forms a complex zinc salt. Prepare a plating bath from the following: zinc chloride 25-35 g / L, ammonium chloride 200-280 g / L, boric acid 30-50 g / L, sulfocarbamide 1-2 g / L, polyethylene glycol 2-3 g / L, detergent 0.2-0.5 mL / L. Dissolve the ammonium chloride in warm water. Add the zinc chloride. Stir until dissolved. Dissolve the other salts in a small amount of warm water, and pour them one by one into the chloride solution. After stirring, add more water to reach the desired volume. Adjust the pH value of the solution to between 5.4 and 6.2 by using citric acid or concentrated aqueous ammonia solution. Pour the fresh electrolyte into a beaker. Place a glass rod across the mouth of the beaker. Clean the surface of the object to be plated, i.e. the copper strip, by removing grease, polishing with fine sandpaper, and washing with water. Use a weak acid solution to remove any oxide rust. Hang from the glass rod a strip of polished copper and a strip of polished zinc. Leave a safe distance between the two strips. Connect the copper strip electrode to the negative terminal of a 1.5 V battery. Connect the zinc strip electrode to the positive terminal of the battery. After a few minutes of electric current flowing, a silvery layer of zinc deposits on the copper strip.

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15.1.7 Electroforming with copper
Make a wax impression by pressing a key into soft wax in a crucible. Remove the object. Insert copper wire in the wax. Dust the impression of the key with bronze powder that should also contact the wire. Copper plate for an hour. Remove from mould and dry. Fill with molten solder.

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15.1.8.0 Anodize aluminium
See diagram 15.1.8.1
Use anodizing to remove oxides from the surface of objects then coat them with a hard and thick oxide layer which can absorb dyes or hardening agents. Anodized dyed aluminium is used for door and window frames, saucepan lids, and wherever bright reflective surfaces are required. Anodized steel may be used in small objects for marine construction. Anodized titanium is used for shiny jewellery.
1. Put two pieces of aluminium foil in hot sodium hydroxide solution to remove any aluminium oxide layer. Rinse in water, then nitric acid, then water. Use the two pieces of aluminium as electrodes in a solution of dilute sulfuric acid connected to a 6 volt battery. Let electric current flow for 15 minutes. The piece of aluminium attached to the positive terminal, the anode, now has a layer of aluminium oxide. Put this piece of aluminium in a water and a dye, e.g. alcohol solution of congo red (blue in acid and red in alkali). Heat the solution to 70^oC and leave for 15 minutes. The aluminium oxide layer absorbs the dye.
Seal in the dye by putting the anodized metal in boiling water for 15 minutes.
2. Degrease a 12 cm X 3 cm thin aluminium strip by wiping with propanone (acetone). Dip the lower half of the aluminium strip in 1.4 mol per litre sodium hydroxide until effervescence occurs indicating removal of the aluminium oxide layer. Dip the cleaned half in nitric acid to neutralize the sodium hydroxide. Dry the strip without touching it and weigh.
Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) --> 2NaAl(OH)₄(aq)
Al₂O₃(s) + 2OH^-(aq) + 3H₂O(l) --> 2Al(OH)₄^-(aq) [Cleaning the oxide]
2Al(s) + 2NaOH(aq) + 6H₂O(l) --> 2NaAl(OH)₄^-(aq) + 3H₂(g)
2Al(s) + 2OH^-(aq) + 6H₂O(l) --> 2Al(OH)₄^- + 3H₂(g) [Reaction after oxide removed]
3. Line a 1 litre beaker with double aluminium foil. Fill beaker with 2 mol per litre sulfuric acid at 25^oC. Clamp the aluminium strip in the centre of the beaker so that the cleaned half is in the sulfuric acid electrolyte. Use crocodile clips to complete the circuit so that the aluminium strip is positive (the anode) and the and the aluminium foil is negative (the cathode).
Adjust the power pack and rheostat to give an electric current density of 10 to 20 mA per cm^2. If the anode area = 3 cm X 3 cm, then area = 3 X 3 X 2 (2 sides) = 18 cm², so current needed = 18 X 10 to 20 = 180 to 360 mA (0.18 to 0.36 A). Close circuit for about 30 minutes but keep adjusting the rheostat to keep the current constant.
At the anode: 2Al(s) + 3H₂O(l) --> Al₂O₃(s) + 6H⁺(aq) + 6e^-
At the cathode: 6H⁺(aq) + 6e^- --> 3H₂(g)
Combined equation: 2Al(s) + 3H₂O(l) --> Al₂O₃(s) + 3H⁺(g)
Be careful! Hydrogen is given off so no naked flames should be present in the laboratory.
Remove aluminium strip, rinse in water and put in dye solution, e.g. a cold fabric dye. Seal the dye by putting the aluminium strip in boiling water. The oxide coating develops a positive charge that attracts dyes containing coloured anions. The porous oxide layer traps the coloured anions that become sealed in by a layer of Al₂O₃.H₂O formed by heating in boiling water.
Al₂O₃(s) + H₂O(l) --> Al₂O₃H⁺(s) [Positive charged oxide coating] + OH^-(aq)
Measure the gain in mass of the aluminium strip by rinsing in propanone then weighing.

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15.1.9 Silvering and desilvering, plating and deplating silver
Plating and deplating silver on metals or glass are not suitable experiments for schools but perhaps students should know about the reactions. Plating metal surfaces or "resilvering" old mirrors should be done only by chemical companies that specialize in this work because dangerous arsenic compounds must be used. Recovery of silver from photographic and X-ray fixers has commercial significance and perhaps environmental significance because it stops metallic silver entering the water supply. Silver can be electroplated from of fixer solutions using stainless steel cathode to yield a silver flake metal sludge of silver-thiosulfide complex.
Fe + Ag-thiosulfate complex --> Fe²⁺ + Ag(s)
If the current density is too high, sulfide forms.
(S₂)₃)^2- + 2e^- --> S^2- + SO₃^2-
Silver on mirrors or scrap photographic film can be reclaimed with nitric acid to form silver nitrate, or by iron (III) chloride in hydrochloric acid or iron (II) chloride solution to form silver chloride. More active metals, e.g. copper, zinc aluminium and iron, can replace less reactive silver in a galvanic response. However, for a large scale processes, iron is preferred because its salts least pollute the environment. Some people use clean pads of steel wool. Do not encourage students to experiment with the family silver!

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15.2.0 Oxidation and reduction, redox reactions
Oxidation is a loss of electrons and reduction is a gain of electrons.
1. Oxidation occurs when:
1. The substance combines with oxygen, i.e. the addition of oxygen to an element or compound, e.g. burning the substance in air,
C(s) + O₂(g) --> CO₂(g)
2Mg + O₂ –> 2MgO
2CO + O₂ –> 2CO₂
2. The substance loses hydrogen, e.g. In the following reaction, the concentrated acid loses hydrogen and changes to chlorine.
4HCl(aq) + MnO₂(s) --> MnCl₂(aq) + 2H₂O(l) + Cl₂(g)
The removal of hydrogen from a compound. In the following equations H₂S is oxidized:
2H₂S + O₂ –> 2S + 2H₂O
H₂S + Cl₂ –> S + 2HCl
3. Oxidation is an increase of valence. In the following equation divalent iron is oxidized to trivalent iron.
2FeCl₂ + Cl₂ --> 2FeCl₃
2Fe²⁺ + Cl₂ --> 2Fe³⁺ + 2Cl^-
4. Oxidation is the loss of electron(s), e.g. when a ferrous ion changes to a ferric ion
e²⁺ - e^- --> Fe³⁺
2. Reduction occurs when:
1. A substance loses oxygen, e.g. In the following reaction Copper (II) oxide loses oxygen and changes to copper.
CuO(s) + H₂(g) --> Cu(s) + H₂O(g)
2. A substance gains hydrogen, e.g. In this reaction, nitrogen gains hydrogen to become ammonia.
N₂(g) + 3H₂(g) --> NH₃(g)
nitrogen + hydrogen --> ammonia
3. An oxidizing agent helps the oxidation of another chemical. An oxidizing agent is a substance which causes oxidation. An oxidizing agent is easily reduced, i.e. it gains electrons easily. The oxidizing agent gains the electrons and the substance being oxidized loses electrons. During oxidation, the oxidizing agent is reduced. When ferric chloride solution is added to stannous chloride solution, ferric chloride is reduced and stannous chloride is oxidized.
Sn²⁺ + 2Fe³⁺ --> Sn⁴⁺ + 2Fe²⁺
Sn²⁺ - 2e^- --> Sn⁴⁺ (oxidation)
2Fe³⁺ + 2e^- --> 2Fe³⁺ (reduction)
4. A reducing agent helps the reduction of another chemical. A reducing agent is easily oxidized, i.e. it loses electrons easily. Examples of reducing agents include the following:
Zn metal that is easily oxidized to zinc ion, Zn²⁺
Hydrogen sulfide that reacts with chlorine to form sulfur
Carbon reduces lead (II) oxide to lead.
Carbon monoxide reduces Fe (III) oxide to iron in a blast furnace.
5. Oxidation and reduction reactions (redox reactions) must occur together. In a redox reaction, the same number of electrons is gained in the reduction as is lost in the oxidation. In the following reaction, O₂ is an oxidizing agent and the H₂ is a reducing agent:
2H₂(g) + O₂(g) --> 2H₂O(l)

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15.2.1 Oxygen as an oxidizing agent
See 7.1.1: Chemical changes, burn magnesium
Oxygen molecules (O₂) gain electrons to form oxide ions (O^2-).

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15.2.2 Chlorine as an oxidizing agent
See 13.4.7: Reactions of chlorine with sodium
Chlorine molecules (Cl₂) gain electrons to form chloride ions (Cl^-).
2FeCl₂ + Cl₂ --> 2FeCl₃
2Fe²⁺(aq) + Cl₂ --> 2Fe³⁺(aq) + 2Cl^-(aq)

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15.2.2.1 Bromine as an oxidizing agent
Add drops of bromine water to 2 cm of ferrous sulfate in a test-tube. The green ferrous salt turns yellow, forming a ferric salt.
2Fe²⁺ + Br₂ --> 2Fe³⁺ + 2Br^-
Fe²⁺ - e^- -> Fe³⁺ (ferrous ion oxidized)
Br₂ + 2e^- + 2Br^- (bromine reduced)
To prove the presence of a ferric salt, add sodium hydroxide solution to form a brown precipitate of ferric hydroxide.

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15.2.3 Potassium dichromate as an oxidizing agent
Add potassium dichromate solution and drops of dilute sulfuric acid to iron (II) sulfate solution. The dichromate ion (Cr₂O₇²⁺) is reduced to Cr³⁺ and the solution changes from orange to green. The iron (II) ions (Fe²⁺) are oxidized to iron (III) ions (Fe³⁺).
Cr₂O₇²⁺(aq) + 14H⁺(aq) + 6e^-- --> 2Cr³⁺(aq) + 7H₂O(l)
6Fe²⁺(aq) --> 6Fe³⁺(aq) + e^-
Cr₂O₇²⁺(aq) + 14H⁺(aq) + 6Fe²⁺(aq) --> 2Cr³⁺ + 7H₂O(l) + 6Fe³⁺(aq)

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15.2.4.1 Potassium permanganate (VII) oxidizes iron (II) to iron (III)
Add potassium permanganate solution and drops of dilute sulfuric acid to iron (II) sulfate solution. The manganate (VII) ions (MnO₄^-) are reduced to manganese (II) ions (Mn²⁺). The iron (II) ions (Fe²⁺), are oxidized to iron (III) ions (Fe³⁺).
MnO₄^-(aq) + 8H⁺(aq) + 5e^--> Mn²⁺(aq) + 4H₂O(l)
5Fe²⁺(aq) --> 5Fe³⁺(aq) + e-
MnO₄^-(aq) + 8H⁺(aq) + 5Fe²⁺(aq) --> Mn²⁺(aq) + 4H₂O(l) + 5Fe³⁺(aq)

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15.2.4.2 Potassium permanganate (VII) oxidizes glycerol to carbon dioxide and water
Put 3 g of fine crystal potassium permanganate on a coffee in lid on sand. Make a hole in the centre of the potassium permanganate and pour 1 mL glycerol (propane-1,2,3-triol) into the hole. Boil then cool the glycerol first if it has already absorbed water. Observe a bright pink flame and steam given off. Dissolve the residue in water and note a green solution [Mn (VI)] and brown solid [Mn (IV)].

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15.2.4.3 Potassium permanganate solution liberates chlorine from hydrochloric acid
Do this experiment in a fume cupboard. Add potassium permanganate solution to 2 cm of concentrated hydrochloric acid solution in a test-tube with damp filter paper over the edge of the opening. Chlorine gas is given off. Be careful! The damp filter paper becomes bleached.
2MnO4^-(aq) + 16H(aq)⁺ + 10Cl^- --> 2Mn²⁺(aq) + 8H₂O(l) + 5Cl₂(g)
10Cl^-(aq) - 10e^- --> 5Cl₂ (chloride ion oxidized)
2MnO^4-(aq) + 16H⁺(aq) + 10e^- --> 2Mn²⁺(aq) + 6H₂O(l) (permanganate ion reduced)

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15.2.5 Concentrated nitric acid as an oxidizing agent
See 12.3.12: Reactions of concentrated nitric acid and copper
Concentrated nitric acid as an oxidizing agent precipitates sulfur from hydrogen sulfide as a yellow suspension.
H₂S <–> 2H⁺ + S^2-
2H⁺ + S^2- + 2H⁺ + NO₃^- --> S + 2H₂O + 2NO₂
S^2- - 2e^- --> S (sulfide ion oxidized)
4H⁺ + 2NO₃^- + 2e^- --> 2H₂O + 2NO₂ (nitric acid reduced)

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15.2.5.1 Nitrous acid as an oxidizing agent or a reducing agent
1. Nitrous acid can act as an oxidizing agent. Slowly add a solution of sodium nitrite in water to a solution of potassium iodide acidified with dilute sulfuric acid. Iodine forms showing that the nitrous acid produced by the action of the dilute acid on the sodium nitrite has oxidized the potassium iodide. The nitrous acid has itself been reduced to nitric oxide. The nitric oxide forms brown fumes of nitrogen dioxide when it contacts the oxygen of the air.
2NO₂^- + 2I^- + 4H⁺ –> I₂ + 2NO + 2H₂O
When acting as an oxidizing agent, nitrous acid gains electrons and is reduced to nitric oxide.
2NO2^- + 4H⁺ + 2e^- --> 2H₂O + 2NO
2. Nitrous acid can act as a reducing agent. Acidify a solution of potassium permanganate with dilute sulfuric acid and add a solution of sodium nitrite until the colour of the potassium permanganate just disappears. Note the absence of brown fumes of nitrogen dioxide. The solution contains nitric acid and can be tested by the nitrate test. The potassium permanganate has oxidized the nitrous acid to nitric acid. The potassium permanganate is reduced to manganous salts.
2MnO₄^- + 6H⁺ + 5NO₂^- --> 2Mn²⁺ + 3H₂O + 5NO₃^-
Nitrous acid here acts as a reducing agent; it loses electrons and is oxidized to nitric acid.
NO₂^- + H₂O - 2e^- -> NO₃^- + 2H⁺

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15.2.6 Sulfuric acid as an oxidizing agent
See 12.3.13: Reactions of concentrated sulfuric acid and copper

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15.2.7 Hydrogen peroxide as an oxidizing agent
Hydrogen peroxide turns an iodide solution brown, forming iodine and perhaps precipitating black crystals of iodine.
1. Add drops of hydrogen peroxide solution to 2 cm of potassium iodide solution in a test-tube.
2H⁺ + 2I^- + H₂O -->. 2H₂O + I₂
2I^- - 2e^- --> I₂ (iodide ion oxidized)
2H⁺ + H₂O₂ + 2e^- -> 2H₂O (H₂O₂ is reduced)
2. Add drops of potassium iodide solution to 20 vols (6%) hydrogen peroxide solution. Then add the same number of drops of dilute sulfuric acid. Heat gently. Note any colour change. Add drops of starch solution. A blue black colour suggested oxidation of 2I^- to I₂.
H₂O₂(aq) + 2H⁺(aq) + 2e^-- --> 2H₂O(l)
2I^-(aq) --> I₂(s) + 2e^-
H₂O₂(aq) + 2H⁺(aq) + 2I^-(aq) --> I₂(s) + 2H₂O(l)
Or
I₂(s) + I^-(aq) --> I₃^-(aq)
H₂O₂(aq) + 2H⁺(aq) + 3I^-(aq) --> I₃^-(aq) + 2H₂O(l)

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15.2.8 Tests for oxidizing agents by change in colour of iron (II) to iron (III)
Prepare a fresh solution of iron (II) sulfate by dissolving iron filings in dilute sulfuric acid. When the reaction stops, filter the solution. The filtrate is acidified iron (II) sulfate solution that is green. Add the test solutions and gently heat. If the solution turns brown, Fe²⁺ has changed to Fe³⁺ because of the presence of an oxidizing agent.

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15.2.9 Tests for oxidizing agents by change of colour of iron with copper (II) sulfate
Add Iron to copper (II) sulfate solution. Note the colour change. The copper ion is an oxidizing agent. The blue colour is removed as copper forms.
Cu²⁺(aq) + Fe(s) --> Fe²⁺(aq) + Cu(s)

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15.2.10 Tests for oxidizing agents by change of colour of zinc with copper (II) sulfate
In this reaction, the copper ion Cu²⁺ attracts electrons better than the zinc ion, Zn²⁺. The Zn is oxidized to zinc ions and the copper is reduced to copper metal. Red copper precipitates and the solution lose its blue colour.
Add pieces of zinc to copper (II) sulfate solution. The zinc corrodes and goes into solution. Red copper precipitates and the solution lose its blue colour. Add excess zinc so that all the copper precipitates.
Decant the solution and evaporate to leave zinc sulfate crystals. Add excess zinc so that all the copper precipitates. Decant the solution and evaporate to leave zinc sulfate crystals.
Zn(s) + CuSO₄(aq) --> Cu(s) + ZnSO₄(aq)
Zn(s) + Cu²⁺(aq) --> Cu(s) + Zn2+

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15.2.11 Breath test for alcohol using potassium dichromate
The breath after drinking contains ethanol vapour, which can be oxidized by potassium dichromate (K₂Cr₂O₇), and the orange dichromate will be reduced to green chromium ions (Cr³⁺).
Add 1 mL of 0.05% potassium dichromate solution and one drop of concentrated sulfuric acid to a small test-tube. Pour 10 mL pure ethanol (absolute alcohol) into a small distilling flask. Heat the flask slowly. Pass the ethanol vapour through the potassium dichromate solution. The colour of the solution changes from orange to green.
Cr₂O₇^2-(aq) + 8H⁺(aq) + 3C₂H₅OH(l) --> 2Cr³⁺(aq) + 3CH₃CHO(l) + H₂O(l)
K₂Cr₂O₇ + 4H₂SO₄ + 3C₂H₅OH --> K₂SO₄ + Cr₂(SO₄)₃ + 3CH₃CHO + H₂O
K₂Cr₂O₇ + 4H₂SO₄ + 7CH₃CHO --> K₂SO₄ + Cr₂(SO₄)₃ + 7CH₃COOH

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15.2.12 Breath test for alcohol using a breath analyser ("breathalyser")
Test a breath analyser used by police or hospital staff. In some countries the breath testing apparatus used by police to detect motorists who have consumed too much alcohol is called a "breathalyser". Borrow a breath testing apparatus from the police. Ethanol vapour in the breath reduces orange potassium dichromate (K₂Cr₂O₇) to green chromium ions (Cr³⁺).

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15.2.13 Potassium chlorate and potassium persulfate as oxidizing agents
Arrange in test-tube pairs 2 cm of 1. acidified potassium iodide solution 2. acidified ferrous sulfate solution 3. hydrogen sulfide solution 4. concentrated hydrochloric acid. Adding 0.55 cc of potassium chlorate to one set and add 0.55 cc potassium persulfate to the other set. Note the reaction and warm to completion if necessary. Note in which case the reaction occurs more readily. Both potassium chlorate and potassium persulfate are powerful oxidizing agents. The persulfate ion oxidizes by accepting electrons to become sulfate ions, e.g., using potassium iodide.
S₂O₈^2- + 2I^- --> 2SO₄^2- + I₂
S₂O₈^2- + 2e^- -->2SO₄^2- (persulfate ion reduced)

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15.2.14 Hydrogen sulfide as a reducing agent
The use of Kipp's apparatus as a source of hydrogen sulfide is NOT recommended in this document.
1. In a fume cupboard, pass hydrogen sulfide gas into a dilute acidified potassium permanganate solution. The colour of the potassium permanganate disappears but a milky precipitate of sulfur remains.
2MnO4^- + 6H⁺ + 5H₂S –> 2Mn²⁺ + 8H₂O + 5S(s)
2. Pass hydrogen sulfide for ten minutes through a dilute solution of ferric chloride acidified with a few drops of hydrochloric acid. The colour will change from yellow to green. Boil the solution in a dish for two minutes to expel hydrogen sulfide, filter through a double filter paper to remove sulfur, and add caustic soda solution in excess to the filtrate. A dirty green precipitate of ferrous hydroxide will be obtained showing that the ferric ion has been reduced to ferrous ions
2Fe³⁺ + H₂S --> 2Fe²⁺ + 2H⁺ + S(s)

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15.2.15 Sulfurous acid as a reducing agent, ionization reaction
Ionization reaction
H₂SO₃ + H₂O <--> H₃O⁺ + HSO₃^-
HSO₃^- + H₂O <--> H₃O⁺ + SO₃^2-
1. In a fume cupboard, pass sulfur dioxide or sulfurous acid into a dilute acidified potassium permanganate solution. The colour of the potassium permanganate disappears but no precipitate of sulfur is formed. The sulfurous acid has been oxidized to sulfuric acid
2MnO4^- + 6H⁺ + 5SO₃^2- --> 2Mn²⁺ + 3H₂O + 5SO₄^2-
2. Pass sulfur dioxide continuously through a dilute solution of ferric chloride. The liquid turns red because of the formation of a complex sulfite. Transfer the solution to a dish and boil for a few minutes on a tripod and gauze. The resulting solution will be pale green or colourless. Add caustic soda solution in excess to a sample where a dirty green precipitate of ferrous hydroxide shows that reduction is complete.
2Fe³⁺ + SO₃²⁺ + H₂O -->. 2Fe²⁺ + SO₄^2- + 2H⁺
3. Dissolve potassium iodate in water in a boiling tube and pass of sulfur dioxide through it. The iodate is reduced to iodine that is deposited as black crystals.
IO₃^- + 3SO₃^2- --> I^- + 3SO₄^3-
5I^- + IO₃^- + 6H⁺ –> 3I₂ (s) + 3H₂O
If the stream of sulfur dioxide continues for a few minutes, the solution goes clear because of the formation of hydrogen iodide.
I₂ + SO_3^2- + H2O –> 2I^- + SO₄^2- + 2H⁺

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15.3.0 Rusting, corrosion
See also 3.52.1: The mass of iron and its temperature increases during rusting
Rusting is an electrochemical process that needs water and oxygen.
At the anode:
Fe(s) --> Fe²⁺(aq) + 2e^-
At the cathode:
O₂(aq) + 2H₂O(l) + 4e^--> 4OH^-(aq) or
1/2O₂ + H₂O + 2e^- --> 2OH^-
The Fe(OH)₂ solution oxidizes to rust (Fe₂O₃.xH₂O, hydrated iron oxide) Corrosion refers to the unwanted oxidation of metals. Both air and water are necessary for the corrosion of iron. Corrosion is caused by the unwanted oxidation of metals. Both air and water are necessary for corrosion of iron. When in moist air, iron is very liable to form rust, most of which is Fe₂O₃.xH₂O. Rust forms on the surface because of the action of water and oxygen on it. You can show that oxygen occupies about one fifth of the atmosphere by volume based on the decrease in the air volume during rusting.

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15.3.1 Rusting of iron wire
See diagram 15.3.1
Polish 0.4 g (about 130 cm long) of thin iron wire (or thin wire gauze) and curl it into a small ball. Push the ball into the bottom of a 10 mL graduated cylinder. Add water to immerse the iron wire and cover the mouth of the cylinder with a slice of glass. Holding the glass slice, invert the cylinder and adjust the water height to a certain mark (say, "9.0 mL") by carefully moving the glass slice. Stand the inverted graduated cylinder over a dish containing water and remove the glass slice. After two days, much reddish brown rust forms on the surface of the iron wire and the water level rises to show a one fifth decrease (about 1.8 mL if the original water level is adjusted to "9.0") in the air volume inside the cylinder.

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15.3.2 Corrosion of magnesium
Repeat the experiment with magnesium ribbon replacing iron wire. The water height inside the graduated cylinder will go down to give an increase in the air volume. This result comes from the hydrogen gas formed in the reaction of magnesium with water.

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15.3.3 Rusting of steel wool
See diagram 15.3.3
1. Use two measuring cylinders. Push steel wool into the bottom of one measuring cylinder. Leave the other as a control. Pour 50 mL water into each measuring cylinder. Hold a piece of cardboard over the mouth of each measuring cylinder and invert it over a shallow dish containing water. Remove the cardboard. Adjust the height of the water in each inverted measuring cylinder by blowing in air with a bent pipette so that the height of water in the two measuring cylinders is the same. Leave the experiment for several days.
2. Repeat the experiment with salty water. The rusting occurs more quickly not because the sodium chloride takes part in the reaction but because it makes the water more conduction. Similarly the presence of sulfur dioxide in the air in cities and industrial sites increases the rate of rusting.
Fe + 1/2O₂ + H₂ (from water) --> Fe(OH)₂ [iron (II) hydroxide]
4Fe(OH)₂ + O₂ --> 2Fe₂O₃.3H₂O + H₂O [iron (III) oxide]
The Fe(OH)₂ in solution is oxidized to Fe₂O₃.

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15.3.4 Need for oxygen for rusting
1. Compare the heights of water in the two measuring cylinders in the previous experiment. The water level is higher in the cylinder containing the rusted steel wool. The height of water rises until the original volume of air in the cylinder decrease by one fifth. The proportion represents how much oxygen is in air. The lost oxygen is combined with the iron of the steel wool to form rust.
2. Moisten inside a test-tube with water. Put iron filings in the bottom of the test-tube and insert a piece of cotton wool to keep them in place. Invert the test-tube in a beaker that is one third full of water. The water levels inside and outside the test-tube should be the same. Mark the original water level on the outside surface of the test-tube. After two days, the iron rusts and the water level rise inside the tube until it is steady. About one fifth of the original air in the test-tube is used up. This suggests that when iron filings rust, oxygen is used.

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15.3.5 Need for oxygen for corrosion of magnesium
Repeat the experiment with magnesium ribbon replacing steel wool.

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15.3.6 Iron gains weight during rusting
When iron rusts, it changes from Fe to Fe₂O₃.xH₂O.
Weigh some dry iron filings. Leave in moist air for two days. Note any increase in weight as rust forms.

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15.3.8 Oxidation can affect air pressure
Wash a small piece of steel wool in methylated spirit to remove any grease. When it is dry, put it in a test-tube with a 1-hole stopper fitted with a 40 cm length of glass tubing. Clamp the test-tube with the end of the glass tubing under water. Note the level of the water in the tubing at the start of the experiment and after one hour and two hours. Water rises up the tubing as oxygen is used to form rust.

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15.3.9 Rate of rusting under separate conditions
Use three test-tubes inverted over water. Push steel wool moistened with ethanoic acid (acetic acid), water, oil. The reaction forms rust first in 2.1 then 2.2 then 2.3.

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15.3.10 Rate of rusting of iron wire
Fill a 30 mL wide necked bottle with a big ball of polished thin iron wire (about 0.6 g). Add water to soak the iron wire and then pour the water out. Stopper the mouth of the bottle with a rubber stopper fitted with a 40 cm straight glass tube. Invert the bottle and clamp it on an iron stand with the end of the glass tube under the water in a beaker. Mark the original water level on the outside of the glass tube. Note the water height every hour. The water level rises slowly in the first five hours and then goes up at a faster speed of about 0.5-0.6 cm an hour. After one day, rising of the water level slows again.

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15.3.11 Rate of rusting of steel wool
Fit a small wide mouth bottle with a rubber stopper and a glass tube about 3 m long. Fit the bottle with a rubber stopper and a glass tube about 3 m long. Use a bundle of steel wool that is big enough to fill the bottle. Remove any oil from the steel wool by washing it in petrol then leaving it to dry. Put the steel wool in the bottle and insert the stopper fitted with a glass tube. Invert the bottle and support it with the end of the tube under water. Record the water level in the tube each hour.

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15.3.12 Conditions necessary for rusting
See diagram 15.3.6
Rusting needs air and water and increases if the water contains salts. You can prevent rusting by painting outside surfaces or by oiling machinery surfaces or by absorbing moisture with silica gel to protect delicate machinery, e.g. cameras or microscope parts.
Use four test-tubes fitted with corks each containing two identical clean nails. Use rainwater. Half the nail is in contact with water and half the nail is in contact with air. This is the control test-tube. Put anhydrous calcium chloride or silica gel in the test-tube. Plug the test-tube with cotton wool. The nail is in contact with air, but is not in contact with moisture. Pour water into the test-tube and boil for some minutes to expel all the dissolved air. Pour oil on the surface of the water to form an airtight layer. The nail is contact with water, but is not in contact with air. Use salt water. Half the nail is in contact with the salt water and half the nail is in contact with air. The nail is in contact with air and salt water. You can see more rusting in test-tube 1.4 than test-tube 1.1. You see no rusting in test-tubes 1.2. and 1.3.

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15.3.13 Electrochemical prevention of rusting, cathodic protection
A "tin can" is made by covering sheets of iron with tin plate to exclude oxygen. If the "tin can" is scratched and it is wet, the iron corrodes very rapidly because an electrochemical cell is set up.
Wrap a piece of aluminium foil around the lower part of a nail. Put the nail and metal in a test-tube. Add tap water to cover the lower part of the nail. Use these metals: control (no metal), magnesium ribbon, zinc foil, copper wire, tin foil. Put the test-tubes in a test-tube rack put stoppers on the test-tubes and leave them undisturbed for several days. If a very small amount of sodium chloride is added to each test-tube, rusting can occur within an hour. Rusting first starts in the test-tubes containing copper or tin, then it starts in the control. Iron is more active than copper or tin, so the iron forms the positive ion Fe²⁺ to react with negative ions in solution to form precipitates of rust on the nail. No rusting occurs in the test-tubes containing magnesium ribbon or zinc, but the more active magnesium or zinc form ions that react with negative ions to form white precipitates.

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15.3.13.01 Cathodic protection
Cathodic protection can protect iron ships and bridges from corrosion. A more electronegative metal, e.g. zinc, is attached as a "sacrificial anode" that goes into solution instead of the iron. Also, you can apply direct current to make the iron into a cathode.
Wooden sailing ships were protected from fouling organisms by the release of copper ions from copper sheathing of the ship's bottom. However, copper sheathing on an iron bottom ship produced an electrochemical cell in the sea water that corroded the iron. This could be prevented by attaching blocks of zinc the bottom to give cathodic protection to the copper.

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15.3.13.1 Rate of corrosion affected by formation of electric cells
Thoroughly clean short narrow strips of the metals magnesium, zinc, copper and tin and also clean five pieces of pure iron wire. Twist a piece of iron wire tightly around each of the other metals. Into five clean beakers place about equal volumes of tap water. Place the single piece of iron wire in the water in one beaker and place one of the twisted pairs of metal strips in each of the other beakers. Record your observations after one hour, one day, one week. Zinc is used as a protective coating of iron for galvanized iron sheets and galvanized screws and bolts. If the coating is scratched, in the zinc iron rain water cell the zinc corrodes to protect the iron. Also, blocks of zinc are attached to iron ships, bridges and wharfs. In this sacrificial corrosion the zinc corrodes away to protect the less active iron. Iron is coated with tin to make tin plate for tin cans and jam tins. However, if the tin is scratched, the iron corrodes more rapidly than if it were not covered by tin.

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15.3.14 Restore bronze coins, corrosion of alloys
Brown "copper coins" are usually alloys of zinc and tin in copper. "Silver" coins are alloys of nickel in copper. Some "gold" coins are alloys of aluminium and nickel in copper. Corrosion is common in alloys if the metals are not evenly mixed. Old coins and statues made of copper alloys and other copper materials exposed to moist air are often covered with blue green verdigris that is basic copper (II) carbonate CuCO₃.Cu(OH)₂.H₂O. New "copper" coins are shiny, but they soon lose their shine and become a dark copper colour because of a layer of black copper (II) oxide. Old copper coins may be very black between the raised areas for the numbers.
Put drops of vinegar on a copper coin. Leave the coin until the liquid is evaporated. Green blue crystals are left on the coin surface. Scrape off the crystals and wash the coin. The coin now looks shiny because black copper (II) oxide is removed. Use dilute hydrochloric acid to make "new" shiny coins.

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15.3.15 Corrosion of aluminium
Put a piece of aluminium foil in water. Put a copper coin on the foil and leave it for some days. A simple aluminium /copper cell forms and a small electric current can be detected with an ammeter. The aluminium foil has holes where the coin lies on it. The water appears cloudy because of the fine particles of aluminium released during corrosion.

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15.3.16 Clean tarnished silver
The surface of silver can react with chemicals in the air to form black silver sulfide, e.g. silver spoons used for eating eggs that contain sulfur proteins. You can polish off the silver sulfide or dissolve it using a commercial silver dip that contains ammonia or thiourea but in each case you lose some silver. You can save the silver by using the following oxidation reduction reaction that reverses the corrosion process. However, some jewellery designers deliberately create a black patina on sunken surfaces as a background contrast to bright silver surfaces, so they soak the jewellery in potassium sulfide, liver of sulphur, and later buff polish the silver surfaces. Never try to clean silver with household bleach because a hard coat of oxide forms that is very difficult to remove using the methods below.
1. To clean the silver, put a sheet of aluminium in the bottom of a beaker. Put the silver to be cleaned on the aluminium and add baking soda solution (sodium hydrogen carbonate). Warm the solution. The sulfur transfers to the aluminium to form aluminium sulfide and the silver becomes shiny again. Clean tarnished silver with aluminium foil.
2. Add 10 g of sodium bicarbonate (NaHCO₃, baking soda) to hot water in a plastic container. Wrap the tarnished silver in aluminium foil and immerse it in the solution for hours until the silver sheen is restored. The sodium bicarbonate dissolves any aluminium oxide on the aluminium surface.
3Ag₂S(s) + 2Al(s) –> 2Al³⁺(aq) + 3S₂^-(aq) + 6Ag(s)
3. Rub the tarnished silver with "Brasso" or toothpaste, not the gel-type toothpaste, or buff polish the surfaces.
4. Soak the tarnished silver in dilute ammonia solution, cloudy ammonia.
5. Soak the tarnished silver in borax and soap solution in hot water.

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15.4.0 Electrical conductivity of a substance
Conductance or conductivity or is the ratio of the current flowing though a conductor to the potential difference between its ends, i.e. the electric field causing the current to flow. Conductance or conductivity is the reciprocal of resistance or resistivity. The SI unit for conductance is the "siemens", S. The SI unit for its reciprocal is ohms (omega). Pure substances that are gases or liquids at room temperature are not good conductors, e.g. water, alcohol, and olive oil. The liquid metal, mercury, is an exception. Fused solids vary in their conductivity. Molten metals, alkalis and salts are good conductors. Other fused solids are not good conductors. The salts sodium chloride and sodium nitrate, as fused liquids, are good conductors but fused sugar and sulfur are non-conductors.
Use a 6 volt battery and two crocodile clips to grasp the cleaned surface of the solids. Use a light bulb to show when current is flowing. Record the solids, melted solids, liquids, and aqueous solutions that do or do not conduct electricity.

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15.5.0 Electrolysis
Chemical reactions in a liquid, an electrolyte, caused by passing of electric current is called electrolysis. Electrolytes are acids, bases or salts dissolved in water. Electric current enters or leaves the electrolyte through conductors called electrodes. The electrode joined to the positive terminal of the battery is called the anode. Conventional electric current passes from the positive terminal of the battery to the anode. The electrode joined to the negative terminal of the battery is called the cathode. Conventional electric current passes from the cathode to the negative terminal of the battery. An electrolytic cell consisting of a container, electrolyte and electrodes is called a voltameter (not "voltmeter"!). Such a cell used to measure electric charge is called a coulometer. Electrolysis uses a source of electricity to break apart an ionic compound. In an electrolytic cell, an external electricity source, e.g. a battery, forces electrons around the circuit away from the negative terminal of the battery and towards the positive terminal of the battery.
Electric current as ions is carried in the electrolyte within the electrolytic cell. Positive ions, called cations, are attracted to the negative cathode and negative ions anions are attracted to the positive anode.
Oxidation as loss of electrons from the ions in solution to the electrode occurs at the anode. Reduction, as gain of electrons from the electrode to the ions in solution, occurs at the cathode. If the electrolyte is a salt consisting of a metal and non-metal, the metal precipitates at the cathode and the non-metal precipitates at the anode. Electrolysis is used to reduce metal ores to the metal.

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15.5.2 Electrolysis of fused sodium chloride
See diagram 15.5.2: Electrolysis of fused sodium chloride
The melting point of the white crystalline solid sodium chloride is 800^oC but the melting point can be lowered by mixing calcium chloride with the sodium chloride. The molten salt can be decomposed by electrolysis to form molten sodium at the negative cathode and chlorine at the positive anode. Reduction occurs at the cathode and oxidation occurs at the anode.
2Na⁺ + 2e^- --> 2Na
2Cl^- --> Cl₂ + 2e^-
2Na⁺ + 2Cl^- --> 2Na + Cl₂

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15.5.3 Electrolysis of tin (II) chloride
Pour some 2 M tin (II) chloride solution into a Petri dish on an overhead projector.
Focus on two parallel electrodes made of tin or lead.
Pass about 5 V of electric current and observe flakes of tin appearing on the cathode.
Reverse the current to see the tin flakes dissolve

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15.5.6.1 Electrolysis of water, measure volume of hydrogen generated
See diagram 15.5.6.1
Electrolysis of aqueous salt solutions using a variable voltage supply or a 12 volt battery and rheostat
Place two 250 mL burettes over the electrodes. Open the taps of the burettes and fill with acidified water until the burettes are completely filled. Close the switch and adjust to a value of 1 amp. Allow the current to flow for twenty minutes. After opening the switch, slide the burettes in holding clips until the levels of the water inside and outside the tube are the same. Observe the volumes of the gases evolved. With acidified water and platinum electrodes the graph of current against voltage shows current almost zero until voltage exceeds 1.7 volts so Ohm's law does not apply.
During electrolysis of water, or electrolysis of an aqueous solution of a salt, e.g. KNO₃ or Na₂SO₄, the following reactions occur:
O₂ + 4H⁺ + 4e^- <-- 2H₂O, E^o = + 0.82 V
4H₂O + 4e^- --> 2H₂ + 4OH^-, E^o = -0.41 V
So the minimum voltage for electrolysis of pure water = + 0.82 -(-0.41) = 1.23 V

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15.5.7 Electrolysis of dilute sodium chloride solution
Fill the electrolysis apparatus with a dilute solution of sodium chloride. Test the electrolyte with drops of litmus solution. Connect the electrodes to a 6 volt battery. Collect the gas in each arm and test each gas. The reaction forms hydrogen gas at the cathode and oxygen gas at the anode. The gases evolved are the elements in water, not those in sodium chloride. The reduction of sodium ions to sodium does not occur in the presence of water. [Note that the solution next to the anode is acidic! What does NOT happen at the anode is: 2Cl^- --> Cl₂ + 2e^-, because in this experiment it is easier to oxidize water molecules than chloride ions.]
At the cathode: Na⁺(aq) + e^- <--> Na(s) -2.71 V
At the cathode: 2H₂O + 2e^- <--> H₂(g) + 2OH^-(aq) -0.41 V
Sodium ions are not reduced to sodium in the presence of water.
At the anode: 2H₂O(l) + 2e^- --> H₂ (g) + 2OH^-(aq)
At the anode: 6H₂O(l) --> 4e^- + O₂(g) + 4H₃O⁺(aq) (Only if the salt solution is very dilute.)

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15.5.9 Electrolysis of copper (II) sulfate solution, copper carried across is proportional to current x time, Faraday's laws
See diagram 15.5.9
The electrolyte is a saturated solution of copper (II) sulfate + 5% sulfuric acid. The copper voltameter has two clean copper electrodes attached to the sides of a glass jar by clips fitted with terminals so that the cathode can be removed and replaced in the same place. Connect three similar circuits carrying currents 1. 2. and 3. adjusted by rheostats to carry 1. 1 amp 2. 0.5 amp and 3. 0.5 amp.
If electrodes have an immersed area of 8 x 5 cm, currents of 1 A corresponds to a current density of about 0.025 A per cm² and current of 0.5 A corresponds to current density of about 0.012 A per cm². Wash and dry the cathodes then clean with emery paper. Weigh the cathodes and place into the three circuits. Close the three switches simultaneously. After ten minutes, open the switch in the circuit 1. carrying 1 A and open the switch in circuit 2. carrying 0.5 A. After twenty minutes, open the switch in circuit 3. carrying 0.5 A. Remove the cathodes then wash, dry and weigh them again.
The weight of copper carried across is proportional to current x time. The first law of electrolysis, discovered by Michael Faraday, states: The mass of substance liberated during electrolysis is proportional to the charge passed. If mass/charge = the electrochemical equivalent constant of the substance, Faraday's second law states: The amount of chemical produced in different substances by a quantity of electricity is proportional to the electrochemical equivalent constant of the substance.

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15.5.9.1 Microscale electrolysis of an aqueous solution of copper (II) sulfate
See diagram 15.5.9.1
Microscale electrolysis allows very fine observation of changes during electrolysis.
Attach fine copper wire to platinum wire and pass the end of the wire under the lid of a Petri dish to form the anode. In bright light, clean the tip of a piece of the fine copper wire, inspect it with a magnifier then pass the end under the lid of the Petri dish to form the cathode. Tape the electrode wires to the bottom of the Petri dish with tips separated by 5 millimetres and tape the electrode wires to the side of the Petri dish where they pass over the sides. Put two drops of concentrated copper (II) sulfate solution (10 g to 100 mL water) in the Petri dish so that the tips of both electrodes are touching the solution. Place an extra two drops of copper (II) sulfate solution aside to compare colour change. Put specks of solid copper (II) oxide in the solution between the tips of the electrodes. Spread the specks to form a continuous band in the solution between the electrodes. Put the lid on the petri dish and connect the electrodes to a 3V source of direct current. Connect the platinum anode to the positive terminal. Connect the copper cathode to the negative terminal. Use a magnifying glass to observe changes around the electrodes and the specks of copper oxide. When the circuit is closed, deposits of copper appear on the cathode but the rate of deposition later changes. The grains of copper oxide start to disappear into the solution. When the growth of copper deposited on the cathode reaches a grain of copper oxide, the coppers is deposited very rapidly around the grain. Note the bubbles around the anode and later around the cathode. The bubbles may stream from one electrode towards the other electrode. Hold a lighted taper above bubbles appearing at the electrodes. Note the popping noise indicating hydrogen. The blue colour of the solution fades more quickly at the cathode. Increase the voltage briefly to 6 volts then back to 3 volts and observe any changes. Change the space between the electrodes and observe any changes. Place the electrodes parallel instead of tip to tip any observe any changes.

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15.5.9.2 Measure electrochemical equivalent of copper
See diagram 32.2.65
Faraday's first law of electrolysis states that the mass of an element deposited or liberated in electrolysis is proportional to the current and to the time for which the current flows.
Remove the cathode from the voltmeter, thoroughly clean it on both sides, first with emery cloth, and then with deionized water. Calculate the total surface area of the cathode that will be immersed in the electrolyte. Allow 0.02 amps per cm² of the surface area. Replace the cathode, avoiding touching the surface with the fingers, and connect the circuit. Close switch S and adjust the rheostat to the calculated current three minutes to prepare the surface of the cathode. Remove the cathode, wash it in deionized water, then in methylated spirits, dry thoroughly in a current of warm air and weigh it. Replace the cathode, close switch S and record the time. Allow the current to flow for 30 minutes. Use the rheostat to maintain the current constant, I amps. After 30 minutes, open the switch, record the time and remove the cathode. Wash and dry the cathode and weigh again. The mass of copper deposited will be very small.

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15.5.10 Electrolysis of concentrated sodium chloride solution, Nelson cell
See diagram 15.5.10: Industrial Nelson cell
Fix an iron wire gauze cylinder around a porous pot and place both in a beaker. Fill the porous pot and beaker with a saturated solution of sodium chloride. Add drops of phenolphthalein solution to the solution outside the pot. Connect the iron wire gauze to the negative terminal of a 6 volt battery to make it the cathode. Put a carbon rod into the porous pot to make it the anode. Hydroxyl ions form at the cathode. Phenolphthalein turns red. The solution at the anode bleaches wet litmus paper because chlorine is formed.
At the anode: 2Cl^- (aq) - 2e^- --> Cl₂(g)
At the cathode: H₂O + 2e^- --> H₂(g) + 2OH^-

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15.5.11 Electrolysis of potassium iodide solution, electrolytic writing
Soak filter paper in potassium iodide solution then put it on a glass sheet to drain. Connect wet filter paper to negative terminal of 12 V battery with an alligator clip. Connect a carbon electrode to the positive terminal of 12 V battery. Switch on power supply and write on the wet paper. Reverse polarity to erase the writing. The writing forms when the carbon positive electrode touches the wet paper to form dark brown iodine
At the anode: 2I^-(aq) --> I₂(aq) + 2e^-
At the cathode: 2H₂O(l) + 2e^- --> 2OH^-(aq) + H₂(g)

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15.5.12 Electrolysis of acetic acid solution
Cut two 10 mm diameter holes in the bottom of a plastic food container. Insert a clean carbon rod from a 1.5 V dry cell battery through each hole. Seal around the rod with silicon sealer to keep the container is watertight. Attach wires to each carbon rod with crocodile clips. Fill the container with water to cover the carbon rods. Add 10 mL of vinegar. Fill test-tubes with this solution and mount each test-tube over a carbon rod. Connect the carbon electrodes to 6 volt battery. Bubbles form on the electrodes then rise into the test-tubes. Do not collect more than a few mL of the gases because hydrogen gas is very flammable and is explosive when mixed with oxygen.
At the electrode attached to the negative battery terminal:
2H+ + 2e- --> H2(g)
At the electrode attached to the positive battery terminal:
4OH- --> 2H20 + 02 + 4e-
Repeat the experiment with copper wire electrodes dipping into the acid solution

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15.6.0 Electrochemical cells
Electrochemical cells (Voltaic cells) form electricity from chemical reactions. The cell is made up of two half cells. Each half cell consists of an electrode in contact with an electrolyte. It is usually a metal in contact with a solution of one of its salts.

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15.6.13 Magnesium / copper battery
Connect the external circuit before adding the sodium sulfate solution. Clean copper in dilute nitric acid and clean magnesium ribbon in 1 M hydrochloric acid.
Half cells: 1. Magnesium ribbon in contact with 0.5 M sodium sulfate solution, in a jar. 2. Copper strip in contact with 0.5 M copper (II) sulfate solution, in a dialysis tubing bag, then in the same jar
1. Mg(s) --> Mg²⁺(aq) + 2e^- (E^o 2.36 V at 25^oC, 1 atmosphere pressure)
2. 1/2 H₂ --> H⁺ + 2e^- (E^o = 0)
3. Cu(s) --> Cu²⁺(aq) + 2e^- (E^o = -0.337 V)
4. Mg(s) + Cu²⁺(aq) --> Mg²⁺(aq) + Cu(s)
EMF = E^o (oxidation) - E^o (reduction), so EMF = 2.36 - (-0.337) = 2.7 V
At the anode: Oxygen is liberated: 4OH^- --> O₂+ 2H₂O + 4e^-
At the cathode: Hydrogen ions are not reduced to H₂ because the E^o of reaction 2., where E^o = 0 V, is greater than the E^o of reaction 3. ( where E^o = -0.337 V.)

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15.6.14 Nickel / cadmium battery, NiCad battery
Rechargeable battery used to power various small devices, e.g. electric toothbrush. During discharge:
At the cathode: nickel (IV) hydroxide + 2 electrons --> nickel (II) hydroxide (reduction)
At the anode: cadmium - 2 electrons --> cadmium (II) hydroxide (oxidation)
Some people think that this kind of battery shows a "memory effect", i.e. after recharging, the battery later runs down only to the capacity at which last recharged. The solution is to let the battery discharge almost completely before recharging. Constant recharging after use for a short time may produce overcharging and change the form of cadmium crystals in the battery resulting in slower release of electric current and apparent lower voltage.

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15.7.0 Electrode potential of metals
See also : Standard electrode potential, electrode potential, reduction potential, E⁰ | See also 3.86: Electrode potentials of metals
Values of electrode potentials of metals are derived from comparisons with the hydrogen cell under standardized conditions of 1 M solution at 25^oC and 1 atmosphere (101.2 kPa) pressure. The standard hydrogen cell is hydrogen gas from a platinum electrode in 1 M solution of H⁺. If E₀ value is +ve, then the preferred direction of electron flow is left to right. The ion or atom with the greater value of E₀ will attract electrons more easily. A positive value for E₀ means that particles in the half cell attract electrons more easily than particles in the hydrogen half cell. If more than one reaction could occur, the reaction that does occur is the reaction that would form the greatest voltage.
Standard reduction potentials (E₀)

K+ + e- --> K	E0 = -2.92 V
Ba2+ + 2e- --> Ba	E0 = -2.90 V
Ca2+ + 2e- --> Ca	E0 = -2.87 V
Na+ + e- --> Na	E0 = -2.71 V
Mg2+ + 2e- --> Mg	E0 = -2.34 V
Al3+ + 3e- --> Al	E0 = -1.67 V
Mn2+ + 2e- --> Mn	E0 = -1.05 V
Zn2+ + 2e- --> Zn	E0 = -0.76 V
Cr3+ + 3e- --> Cr	E0 = -0.71 V
Fe2+ + 2e- --> Fe	E0 = -0.44 V
Ni2+ + 2e- --> Ni	E0 = -0.25 V
Sn2+ + 2e- --> Sn	E0 = -0.14 V
Pb2+ + 2e- --> Pb	E0 = -0.13 V
2H+ + 2e- --> 2H	E0 = 0.00 V
Cu2+ + 2e- --> Cu	E0 = +0.35 V
Ag+ + e- --> Ag	E0 = +0.80 V
Hg2+ + 2e- --> Hg	E0 = +0.85 V

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15.7.1 Potential difference from combining half cells, zinc and iron
To measure the potential difference of a zinc half cell connected to an iron half cell. Use a strip of zinc metal in a zinc chloride solution and an iron nail in iron (II) sulfate solution. Connect the two half cells with a strip of filter paper soaked in potassium chloride solution to act as a salt bridge. Complete the circuit by connecting leads from each metal to a voltmeter. Read the voltmeter. Electrons flow with potential difference of 0.32 V.
Zn(s) --> Zn²⁺(aq) + 2e^- (E₀= + 0.76 V)
Fe²⁺ + 2e^- --> Fe(aq) (E₀ = -0.44 V)
Zn(s) + Fe²⁺ --> Zn²⁺ + Fe(s) (E₀= + 0.32 V)

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15.7.2 Potential difference from combining half cells, Zn and Cu, Zn and Pb
If Zn E₀ = -0.76 V set up cells to measure the E₀ values of copper (copper in copper (II) sulfate solution) and lead (lead in lead (II) nitrate solution).

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15.7.3 Differences in potential on an iron nail
Soak 1 gm agar in 100 mL water for two hours then boil until dissolved. Add phenolphthalein indicator and add acid or alkali until pH = 8. Add drops of freshly prepared potassium ferricyanide solution and pour into a Petri dish. Add a very clean nail and place the petri dish on an overhead projector. After some hours, a pink colour forms around the shaft of the nail because of hydroxide ions and blue green colour forms around the head of the nail because of Fe²⁺ ions. The stressed head shows positive potential and the unstressed shaft shows negative potential.
At the anode: Fe(s) --> 6 Fe²⁺(aq) + 2e^-
At the cathode: O₂(aq) + 2H₂O(l) + 4e^-- --> 4OH^-(aq)

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