topic15

School Science Lessons
Topic 15 Electrochemistry, electrochemical cells, electrolytic cells
Updated 2009-06-23
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites

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

32.5.0 Power and energy
32.6.0 Circuit analysis, house circuits
32.7.0 Instruments to detect electric current
33.3.0 Cells and Batteries
33.4.0 Dry cells, Leclanche cell, flashlight battery
33.5.0 Lead accumulator cell, car battery
33.6.0 Thermoelectricity
33.7.0 Piezoelectricity

3.59.0 Conductors of electricity, electrical conductivity
3.59.1 Substances that conduct electricity
3.59.2 Electrical conductivity of solids
3.59.3 Electrical conductivity of melted solids, fused solids
3.59.4 Electrical conductivity of liquids
6.38 Electricity conductors (primary)

15.00 Electrolytes
8.6 Prepare electrolyte for a lead accumulator cell
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
15.05 Electrolytes in the blood and urine

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
15.1.10 Electroplating copper, copper flashing of iron
15.1.11 Lead tree and tin tree
15.1.12 Electroplating with silver
15.1.13 Cucumber pickle frying
15.1.14 Silver coulometer

15.2.0 Oxidation and reduction, redox reactions, Oxidation is a loss of electrons and reduction is a gain of electrons
15.2.01 Oxidation occurs when:
15.2.02 Reduction occurs when:
3.76 Reduction of potassium permanganate with sulfur dioxide
12.17.1 Reactions of manganese (II) salts, Mn
12.17.2 Prepare manganates, MnO₄^2-12.17.3 Prepare potassium permanganate, KMnO₄
12.17.4 Reactions of permanganate ion, MnO₄^-
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
12.5.6 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
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

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
12.8.1 Reactions of iron (II) salts and iron (IlI) salts, Prussian blue
12.8.2 Rusting
12.8.3 Oxidation of iron (II) salt
12.8.4 Burn steel wool
12.8.5 Reduction of iron (IlI) salts
12.8.6 Heat iron filings with powdered sulfur
12.8.7 Prepare iron (II) oxide, FeO
12.8.8 Heat iron (II) sulfide, (FeS₂, pyrite, fool's gold)
12.8.9 Prepare iron (IlI) oxide, Fe₂O₃12.8.10 Show that black iron oxide is a mixed base
12.8.11 Iron displace hydrogen from sulfuric acid to form iron (II) sulfate
12.8.12 Iron displaces hydrogen from hydrochloric acid to form pale green iron (II) chloride
12.8.13 Heat hydrated iron chlorides
12.8.14 Prepare iron (II) ammonium sulfate (NH₄)₂SO₄.FeSO₄.6H₂O
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 Iron gains weight during 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 Clean tarnished silver

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
15.7.0 Electrode potential of metals
33.3.1 Simple electric cell
33.3.2 Voltaic cell, Daniell cell, with salt bridge
33.3.3 Coin cells
33.3.4 Lemon cell
33.3.5 Simple chemical rectifier
33.3.6 Put chocolate wrapper cell in the mouth
33.3.7 Noisy potato cell
33.3.8 Hydrogen / oxygen fuel cell
33.3.9 Ionic migration
33.3.10 Ionic friction
33.3.11 EMF dependence on electrode material
33.3.12 Contact potential difference
33.3.13 Crowsfoot or gravity cell

15.7.0 Electrode potential of metals
15.7.1 Potential difference from combining half cells, zinc and iron
15.7.2 Potential difference from combining half cells, Zn and Cu, Zn and Pb
15.7.3 Differences in potential on an iron nail

3.59.1 Substances that conduct electricity
See diagram 3.59.1
The carbon rods in the stopper should be 5 min apart. Use crocodile clips to attach a conducting wire between one battery terminals and the bulb and the other battery terminal and one of the carbon rods. The bulb should light when a current passes so lightly touch both carbon rods with copper wire to make the bulb light and show that the circuit is works. Prepare separate beakers of sugar, sodium carbonate, sodium chloride and laundry starch. Dip in the carbon rods in each beaker and record whether the bulb lights up to show that the solution is conducting electricity between the carbon rods. Add water to each beaker. Dip in the carbon rods in each solution beaker and record whether the bulb lights. Wash the rods thoroughly under the tap after dipping in each solution. Note any signs of chemical reaction in the beaker. None of the original solid substances conduct electricity. Sodium carbonate solution and sodium chloride solution conduct electricity. These solutions are electrolytes. Solutions of sugar, starch and methylated spirit do not conduct electricity. They are non-electrolytes. Repeat the experiment by testing dilute hydrochloric acid and dilute sodium hydroxide. Acids, salts, and alkalis are electrolytes. When dissolved in water to form solutions or melted into liquids by heating, they conduct electricity. Electrolytes are usually decomposed when electric current passes through them, electrolysis. In electrolysis, the carbon rod (electrode) connected to the negative (-) terminal of the battery is the cathode, and the electrode connected to the positive (+) terminal is the anode. Gases from the decomposition of electrolytes may be seen as bubbles on the electrodes.

3.59.2 Electrical conductivity of solids
See diagram 3.59: Electrical conductivity apparatus
Use two carbon electrodes from torch batteries, a non-conducting support for the electrodes, crocodile clips or crunched aluminium foil for connections, light bulbs to show when current flows, and a 6 V dry cell power source. Test the conductivity of solids by making a good contact between the cleaned surface of the solid and the two electrodes. Confirm that metals and carbon conduct electricity. Test the conductivity of non-metallic and crystals, e.g. calcite (crystalline calcium carbonate) candle wax, copper (II) sulfate-5-water, ethanedioic acid-2-water (oxalic acid) glass rod, naphthalene, plastics, octadecanoic acid, sucrose (cane sugar) sodium chloride crystals, sodium nitrate, sugar crystals, sulfur, wax. None of these solid compounds is a good conductor.

3.59.3 Electrical conductivity of melted solids, fused solids
Be careful! Do not let the carbon electrodes ignite and burn.
Grip two carbon electrodes from used dry cell batteries with the crocodile clips. Test the conductivity of the melt by dipping in the electrodes. Wait for the electrodes to reach the same temperature. This ensures that the electrodes are in contact with the liquid and not the solidified melt. Scrape and clean the electrodes between each test.
1. Melt substances that are solids at room temperature, but heat very gently, otherwise they may ignite and burn, e.g. candle wax, cellulose acetate (acetate rayon) lead metal, lead bromide, naphthalene, nylon, octadecanoic acid (stearic acid) polyethylene, polythene, Perspex, potassium iodide (m.p. 682^oC) sodium chloride, sodium nitrate, solder, sulfur, tin metal. Melted solids vary in their conductivity. Only molten metals, alkalis and salts are good conductors. Sugar and sulfur are non-conductors.
2. Glass can be a conductor. Heat a glass rod until it becomes very hot and begins to soften. Test the hot, soft part with the conductivity apparatus. When molten, glass is a good conductor of electricity.

3.59.4 Electrical conductivity of liquids
Pure substances that are gases or liquids at room temperature are not good conductors, but the liquid metal mercury is a good conductor.
1. Clean and dry the carbon electrodes between each test. To test the conductivity of liquids, immerse the ends of carbon electrodes 3 mm apart in acetone, copper (II) sulfate solution, methylated spirit, liquid paraffin, olive oil, peanut oil, sodium chloride solution, sugar solution, turpentine (mineral turps) vinegar.
2. Test the conductivity of solutions, e.g. 2 M concentration of the following:
2.1 Strong electrolytes, e.g. copper (II) sulfate, hydrochloric acid, potassium hydroxide, sodium chloride, sodium hydroxide, sodium nitrate, sulfuric acid.
2.2 Weak electrolytes: ammonia solution, benzoic acid, ethanoic acid (acetic acid). Always wash the electrodes thoroughly after testing each solution. Solutions of acids alkalis and metallic salts are generally good conductors. Solutions of sugar and alcohol are non-conductors. Solutions of other types of substances in water and in other liquids are generally non-conductors.
3. Test demineralized water for conductivity. The bulb does not light. Very gradually stir small crystals of sodium chloride into the water. Note any light from the light bulb as the salt dissolves. Similarly test distilled water, tap water and mineral water.
4. If a commercial conductivity meter is available, nonelectrolytes show a very small current but a completely dissociated strong electrolyte e.g. 0.1 M HCl, shows a current > 100 mA.
4.1 Dilute 5 mL of 0.1 M solution of:
4.1.1 HCl,
4.1.2 NaOH to 50 mL.
Test each reactant solution then mix the two solutions and test half the volume of the product solution. The conductivity of the product solution is less than the conductivity of each of the reactant solutions.
4.2 Test 5 mL of 0.1 M solutions of
4.2.1 acetic acid, HC₂H₃O₂,
4.2.2 aqueous ammonia solution.
Test each reactant solution then mix the two solutions and test half the volume of the product solution.
4.3 Test 5 mL of 0.1 M solutions of:
4.3.1 H₂SO₄,
4.3.2 Ba(OH)₂.
Add 3 drops of 0.1% thymol blue indicator solution to the sulfuric acid solution then add drops of the 0.1 M Ba(OH)₂ solution while stirring until the indicator changes from pink to yellow to blue. Test the conductivity of the product solution.

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.

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.

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.

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 copper (II) sulfate solution 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.

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.

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.

15.05 Electrolytes in the blood and urine
In medical use electrolyte refers to the ions. So serum electrolyte refers to sodium, potassium or chloride ions that function in cardiac rhythm, skeletal muscle contraction and nerve transmission. The level of bicarbonate ion is important for the acid-base balance in the blood. The urine electrolytes, sodium and potassium, indicate electrolyte balance and how hormones affect the function of the kidney.

15.2.01 Oxidation occurs when:
1.1 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
1.2 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
1.3 Oxidation is an increase of valence. In the following equation divalent iron is oxidized to trivalent iron.
2FeCl₂ + Cl₂ --> 2FeCl₃
2Fe²⁺ + Cl₂ --> 2Fe³⁺ + 2Cl^-
1.4 Oxidation is the loss of electron (s), e.g. when a ferrous ion changes to a ferric ion.
e²⁺ - e^- --> Fe³⁺

15.2.02 Reduction occurs when:
2.1 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)
2.2 A substance gains hydrogen, e.g. In this reaction, nitrogen gains hydrogen to become ammonia.
N₂ (g) + 3H₂ (g) --> NH₃ (g)
nitrogen + hydrogen --> ammonia
2.3 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)
2.4 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.
2.5 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 H₂is a reducing agent:
2H₂ (g) + O₂ (g) --> 2H₂O (l)

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-).

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)

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.

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)

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)

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)].

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)

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)

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 sodium nitrite solution to potassium iodide solution 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 potassium permanganate solution with dilute sulfuric acid and add sodium nitrite solution 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⁺

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

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 + I₂2I^- - 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)

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.

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)

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⁺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

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⁺). The legal limit in some countries is 80 mg of ethanol per 100 mL of blood.

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)

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)

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⁺

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.

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.

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.

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)₂in solution is oxidized to Fe₂O₃.

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.

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

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.

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 one-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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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 one of the metal salt solutions.

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^o2.36 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.)

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.

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 E₀will attract electrons more easily. A positive value for E₀means 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

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)

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 E₀values of copper (copper in copper (II) sulfate solution) and lead (lead in lead (II) nitrate solution).

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)