School Science Lessons
Topic 15a Electrolysis, electroplating
2014-07-12
Please send comments to: J.Elfick@uq.edu.au
Table of contents
15.5.0 Electrolysis
15.1.0 Electroplating

15.5.0 Electrolysis
Electrolysis
15.5.28 Chemical reaction forms electricity, magnesium / copper cell
18.7.6 Dissolve chlorine in swimming pool water by electrolysis
15.5.30 Electrolysis and mass of sodium atom by electrolysis with an electrolytic rectifier
15.5.29 Electrolysis and mass transfer
15.5.31 Electrolysis and oxidation of ferrous to ferric iron
15.5.24 Electrolysis of acids, acetic acid solution
15.5.25 Electrolysis of acids, hydrochloric acid
15.5.7 Electrolysis of aqueous salt solutions with variable voltage supply or 12 V battery
15.5.19 Electrolysis of copper (II) sulfate solution, electrochemical equivalent of copper
15.5.15 Electrolysis of copper (II) sulfate solution, Faraday's laws
15.5.16 Electrolysis of copper (II) sulfate solution, microscale electrolysis
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
3.68 Electrolysis of lead (II) bromide melt
15.5.23 Electrolysis of potassium iodide solution, electrolytic writing
15.5.27 Electrolysis of potassium iodide solution, prepare iodine solution
3.69.6 Electrolysis of silver nitrate solution, with an OHP
15.5.1 Electrolysis of sodium chloride melt
3.69.1 Electrolysis of sodium chloride solution
15.5.12 Electrolysis of sodium chloride solution
3.69.2 Electrolysis of sodium chloride solution, saturated solution
15.5.13 Electrolysis of sodium chloride solution, Nelson cell
15.5.32 Electrolysis of sodium ions through glass
15.5.21 Electrolysis of sodium sulfate solution
3.69.6 Electrolysis of silver nitrate solution, with an OHP
15.5.20 Electrolysis of tin (II) chloride solution, with an OHP
See pdf Electrolysis of water
15.5.4 Electrolysis of water, decomposition of water, Hofmann voltameter, electrochemical coulometer
3.69.1.1 Electrolysis with carbon electrodes
15.5.33 Electric forge
15.1.0 Electroplating
15.1.0
Electroplating
15.1.8 Anodize aluminium
15.1.8.1 Anodize iron nails
15.1.3 Chromium plating
15.1.2 Copper plating
15.1.13 Cucumber pickle frying
15.1.9.1 Disposal of photography wastes
15.1.7 Electroforming using copper
15.1.10 Electroplating copper, copper plating, copper flashing of iron
15.1.1 Faraday's first law
15.1.11 Lead tree and tin tree
15.1.4 Nickel plating
15.1.14 Silver coulometer
15.1.5 Silver plating of copper or nickel
15.1.9 Silvering and desilvering, plating and deplating silver
15.1.1.1 Tests for Faraday's first law with copper and copper (II) sulfate solution
15.1.6 Zinc plating of copper

15.1.0 Electroplating
To electroplate a metal is to coat it by electrolytic deposition, e.g. chromium plating, silver plating, to protect the metal from corrosion and to improve the appearance. You can pass electric current through an electric cell, called an electroplating bath, to deposit one metal on another. The metal to be electroplated must first be cleaned and then placed in a solution of a compound of the used for the
coating in the electroplating. The metal is then connected to a battery and it becomes the cathode, the negative electrode, in the electroplating process. Positive metallic ions, e.g. copper, gold, silver zinc, form a coating on the metal. Metal may be taken from the anode and deposited on metal articles that act as cathodes. The electrolyte contains the metal to be deposited as ions. A salt of the plating metal is the electrolyte, e.g. chromium salt for chromium plating. Electroplating is not the same as anodizing that forms an oxide on the metal.
Put two pieces of the same metal, e.g. pieces of wire, or two metal spoons or knives in a copper (II) sulfate solution. Connect the pieces of metal to a battery. The metal joined to the negative terminal of the battery will become the cathode and the other piece of metal will become the anode. Close the circuit and note the coating of copper on the surface of the cathode.

15.1.1 Faraday's first law
(Michael Faraday 1791-1867)
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.
One mol of electrons (one Faraday) corresponds to 96 500 coulombs.
Q = i × t
1 coulomb = 1 amp × 1 second
15.1.1.1 Tests for Faraday's first law with copper and copper (II) sulfate solution
See diagram 3.3.4: Electrolysis
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. Repeat the experiment with other metals.

15.1.2 Copper plating
See diagram 3.3.5: Key in copper sulfate solution
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.
5. In a beaker of copper sulfate solution place a clean strip of brass to act as the cathode (negative electrode) and a carbon rod to act as the anode (positive electrode). The electrodes should not touch. Let current pass until the coating of copper has formed on the brass. Try copper plating other objects, e.g. an old spoon

15.1.3 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, Cr2O72-. 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 50oC.

15.1.4 Nickel plating
1. 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.
2. Thoroughly clean a piece of brass or copper. Make a solution of nickel in the beaker. Connect the piece of brass or copper to the negative terminal of the battery and a carbon rod to the positive. Place both the metal and carbon electrodes in the solution, and examine the metal from time to time. When it is well-covered with nickel, remove it and wash it under the tap. Dry the metal and polish it with metal polish.

15.1.5 Silver plating of copper or nickel
1. 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.
2. The procedure is the same as with copper except that the electrolyte is a solution of about 1 g silver nitrate in 20 mL water. The deposit will be dull. Shiny electroplated deposits are usually obtained by vigorous mechanical polishing of the dull film produced in the first instance. Plate with silver by connecting the object to the negative terminal and using silver nitrate solution.

15.1.6 Zinc plating of copper
Use a mixture of zinc chloride, boric acid and ammonium chloride. The reaction forms a complex zinc salt. Prepare plating baths 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.

15.1.7 Electroforming using 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. .

15.1.8 Anodize aluminium
To anodize is to give a metal a positive coating, usually aluminium, by an electrolytic process where the metal is the anode. Anodizing is the electrochemical conversion of the aluminium surface to a hard aluminium oxide surface that can be coloured with  organic dyes or inorganic metal compounds. It is called anodizing because it thickens the oxide layer on the anode. 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.
Experiments
See  diagram 12.1.11
: Anodize aluminium
1. Line a beaker with aluminium foil to act as a cathode, Add 50 mL of 1 M (10%) sulfuric acid. Clean the surface of a strip of aluminium sheet by putting it 1 M sodium hydroxide for a few seconds, then in 2 M nitric acid for a few seconds. Rinse the strip  in running tap water, dry it with cotton wool, then rinse it again. Fold the aluminium strip over a pencil to act as an anode. Place the strip in the sulfuric acid with the pencil acting as a support over the rim of the beaker. Connect the cathode and anode to a 12 volt DC power supply. Remove the aluminium strip after ten minutes and observe the change in the surface of the metal. Colour the aluminium strip in a boiling organic dye, e.g. cochineal, eosin, alizarin. Seal the dyes into the surface of the strip by putting it into boiling water.
See diagram 15.1.8.1: Anodize aluminium
2. 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 dilute sulfuric acid solution 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 70oC 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.

3. Degrease a 12 cm × 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]
4. Line a 1 litre beaker with double aluminium foil. Fill beaker with 2 mol per litre sulfuric acid at 25oC. 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 cm2.
If the anode area = 3 cm × 3 cm,
then area = 3 × 3 × 2 (2 sides) = 18 cm2,
so current needed = 18 × 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 gas 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.

15.1.8.1 Anodize iron nails
Connect two iron nails dipping into a weak solution of salt water to the terminals of a 4.5 volt battery. Bubbles appear on both nails. The nail connected to the positive terminal of the battery (anode) develops a coating of rust, i.e. iron oxide.

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- --> S2- + SO32-
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!

15.1.9.1 Disposal of photography wastes
Treat photographic 'fixer' wastes in the three separate steps below or sell the wastes to photographers, or chemical recyclers.
1. To precipitate silver as insoluble silver chloride, add 20 g of sodium chloride per litre of waste solution
1.1 Leave the solution for days until the salt dissolves. Silver will precipitate as a fine white solid of silver chloride.
1.2 Decant the solution and collect the solid silver chloride for recycling
2. To react with thiosulfates, add 5 M hydrochloric acid to the decant in a plastic container in a fume cupboard or outside with the wind behind.
2.1 Leave for one hour or until any reaction is completed.
3. To prepare the waste solution for discharge into the sewer, add phenolphthalein or litmus indicator. Then add 2 M sodium hydroxide until the solution is just basic, pH 7-8
3.1 Add a further 20 mL sodium hydroxide until the indicator is pH 8-10.
3.2 Discharge this basic mixture of common salts directly to the sewer. This solution is free from silver and ions that can release noxious gases in the sewer.

15.1.10 Electroplating copper, copper flashing of iron
See diagram 15.1.10: Electroplating
1. Obtain two carbon rods from old torch cells. Put them, not touching, in a 10% solution of copper (II) sulfate, 10 g copper (II) sulfate crystals, 90 mL water. Connect to two torch cells in series. Observe the surface of the rods after ten minutes. Note any changes. Replace the rods and reverse the leads to the cell. Note what happens.
2. Take an article, say of brass, iron or silver, which you wish to electroplate with copper. Iron is not very suitable because when immersed in copper (II) sulfate solution it partly dissolves and a loose adherent of coating of copper is formed. Connect the article to the battery connection from which the hydrogen gas was produced in the previous experiment. The electrolyte is a solution of copper (II) sulfate (about 10%) in water. The other electrode can be copper wire. When a current is passed through the circuit, a film of
copper gradually appears on the article being plated. Simultaneously copper will be dissolved from the copper wire electrode that after a time becomes noticeably eaten away. Copper is deposited at one electrode and passes into solution at the other.
3. Pass electric current through copper (II) sulfate solution You will need a 250 mL beaker, piece of cardboard, two carbon rods from old dry cells, and dilute copper (II) sulfate solution. The carbon electrode connected to the positive wire is the anode and the electrode connected to the negative wire is the cathode. The copper (II) sulfate solution is called the electrolyte. The cathode becomes coated with copper. The coating becomes thicker the longer the current is flowing. The copper (II) sulfate solution becomes less blue after about one hour because the copper is removed from the solution and placed on the cathode. The blue colour of the solution was caused by the copper in it.

4. Use copper and carbon electrodes in a copper (II) sulfate bath to plate copper onto a carbon electrode.

5. Plate polished iron in a copper (II) sulfate solution. Plate with copper by connecting the object to the negative terminal and using copper (II) sulfate solution.

15.1.11 Lead tree and tin tree
1. Make a tin tree pass current between lead electrodes in a saturated solution of lead acetate to cause fern-like clusters to form on the cathode.
2. Make a tin tree pass current between electrodes of copper and tin in an acid solution of stannic chloride so that with copper as the cathode, tin crystallizes as long needles.

15.1.12 Silver plating
The procedure is the same as with copper except that the electrolyte is a solution of about 1 g silver nitrate in 20 mL water. The deposit will be dull. Shiny electroplated deposits are usually obtained by vigorous mechanical polishing of the dull film produced in the first instance. Plate with silver by connecting the object to the negative terminal and using silver nitrate solution.

15.1.13 Cucumber pickle frying
Apply high voltage across a cucumber pickle and it lights at one end.

15.1.14 Silver coulometer
Plate silver in a silver nitrate bath onto a platinum cup. A silver coulometer shows a 1 g change in anode weight when 1 amp is passed for 1 sec.

15.5.1 Electrolysis of sodium chloride melt
See diagram 15.5.2: Electrolysis of fused sodium chloride
In electrolysis of a melt, only two ions are present in the dissolved salt. Molten sodium chloride, above 800oC, forms sodium at the cathode and chlorine gas at the anode. The melting point of the white crystalline solid sodium chloride is 800oC 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

15.5.4 Electrolysis of water, decomposition of water, voltameter, Hofmann voltameter, electrochemical coulometer
This experiment is also called "electrolysis of an aqueous salt solution, "electrolysis of acidified water" and "Hoffman electrolysis apparatus". (Wilhelm von Hofmann, Germany, 1818–1892)
Hofmann voltameter acrylic with Pt electrodes, "Scientrific", (commercial website)
Coulomb meter, "Scientrific", (commercial website)
M64 Water Electrolysis Demonstrator, "Prof Bunsen", (commercial website)
See diagram 3.2.69: Electrolysis apparatus | See diagram 15.5.4: Hofmann voltameter | See diagram 15.5.6: Electrolysis
See diagram: 15.5.6: Decomposition of acidified water | See diagram 15.5.3: Electrolysis of an aqueous salt solution
See 15.59.1: Substances that conduct electricity
See diagram 3.69: Electrolysis apparatus: A pencil "lead", B rubber stopper, C screen wire, D soldered joint, E copper wire, F battery, G test-tube
1. A voltmeter measures electric potential in volts.
2. A voltameter, coulometer, is an electrolytic cell, consisting of a container, electrolyte and electrodes, used to find the amount of electricity that has flowed during electrolysis by measuring products of the process, e.g. the weight of substances gained or lost by the electrodes. A voltameter, coulometer,  measures electric current by measuring the amount of metal deposited or gas liberated from an electrolyte in given time caused by the passage of the current, with SI unit the coulomb.
3. A Hofmann voltameter, is used to demonstrate the decomposition of water into hydrogen gas and oxygen gas. It has platinum electrodes that are too inert to allow deposition of any chemical on them.
4. Voltameters were used to measure current before the invention of the ammeter.
5. Electrochemical coulometers are used to measure the concentration of different ions in solution by plotting the relation of current and voltage in a micro electrode and for chemical analysis in which the quantities are determined from the amount of electricity needed for the electrolysis..                                 

Experiments
1. Pure water is a poor conductor of electricity so add a small quantity of an electrolyte to improve conductivity, e.g. dilute H2SO4 or KNO3 or Na2SO4. Fill the cylinder and two test-tubes with the acidified water. Put a finger over each test-tube and invert it over an electrode. Add bromothymol blue indicator. Connect the cell to a 12 V d.c. supply and use a current of 1 A. Watch for bubbles of gas at both electrodes, but if no bubbles appear, add more acid. Note that when the first test-tube is full of gas, the second test-tube is only half full of gas. Remove each test-tube when it becomes full of gas. Keep the test-tube inverted and apply a stopper.
Test for hydrogen gas with a lighted splint, match test, that causes a sharp popping sound.
Test for oxygen with a glowing splint, match test, that reignites to form a flame again. Two volumes of hydrogen gas form at the cathode for each volume of oxygen that forms at the anode.
Repeat the experiment by swapping the leads to the battery and replacing the test-tubes, then the sequence of filling the test-tubes is reversed. The acid in the electrolyte remains constant, so the hydrogen gas and oxygen gas have come from the electrolysed water that has decreased in volume.
At the anode: 2H2O (l) --> O2 (g) + 4H+ (aq) + 4e- [loss of electrons from the solution to the circuit]
[Bromothymol blue is yellow in acidic solutions.]
At the cathode: 4H2O (l) + 4e- --> 2H2 (aq) + 4OH- (aq) [gain of electrons from the circuit to the solution]
[Bromothymol blue is blue in basic solutions.]
Overall reaction: 2H2O (l) --> 2H2 (g) + O2 (g) [Bromothymol blue is green in neutral solutions.]
2. Stir into the water sodium sulfate and test it again. Compare the results. Water alone does not conduct electricity, but the addition of sodium sulfate, or any electrolyte, makes it conduct. Pure water does not conduct an electric current, but if an electrolyte is added to it the water conducts and, with many electrolytes, is decomposed into its elements, hydrogen and oxygen. Hydrogen gas forms at the cathode and oxygen forms at the anode. But sometimes the electrolyte itself is decomposed.

3. Use a voltameter to prepare hydrogen electrolytically and collect the gas in an inverted small borosilicate test-tube.
Pass d.c. current through slightly acidic water evolves hydrogen gas and oxygen at the electrodes. Use a gas coulometer to measure the volume of gas from electrolysis. Use phenolphthalein as an indicator in electrolysis demonstrations. Use purple cabbage as an indicator to show electrolysis of sodium sulfate. Use the standard commercial Hofmann apparatus for electrolysis of water. Place Tygon tubing over the wire coming out the bottom to protect it from the acid. Use a projection electrolytic cell to show the
evolution of gas. Make soap bubbles with the gases from electrolysis of water and blow them to droplets.
4. Prepare hydrogen / oxygen mixture by electrolysis of dilute acid / detergent mixture in a beaker. Collect the gases produced from both electrodes a funnel set beneath the surface of the liquid. The bubbles that float on the surface explode away from the funnel when ignited with a taper.

5. Water by itself does not conduct an electric current but does so if an electrolyte is added to it. Fill your small 2 delivery tubes beaker with water and stir in a finger width of sodium hydrogen sulfate. When the powder has dissolved, pour the clear liquid into the U-tube so that the level is at least 3 mm below the sidearms. The U-tube must be well supported, exactly vertically. One way to do this is to place it in a cup, with wads of paper to keep it in position. Push the stoppers containing the carbon rods firmly but into the U-tube. fit the delivery tubes to the sidearms of the U-tube with the special rubber connections.
Make the delivery tubes from your glass tubing. The end of each delivery tube projects into an inverted test-tube of water in a deep dish. Support the test-tubes with cork-lined apparatus clamps. When all is ready, the chemical alone in the beaker. It is important to wash the rods thoroughly between each test, so have a container of water available so that the rods can be dipped into it before each substance is tested. Record your observations as to any signs of chemical reaction in the beaker.
As soon as the connection is made bubbles of gas will be seen coming from the carbon rods. The gases enter the inverted test-tubes, thereby displacing water into the dish. The gases are hydrogen gas and oxygen, from the decomposition of the water, and there is twice as much of one as of the other. Note whether the hydrogen gas comes from the cathode (negative connection to the battery) or the anode.
Remove each test-tube when full, place your finger or thumb over its mouth, and tests for hydrogen gas and oxygen. Describe what you see. The experiment forms twice as much hydrogen gas as oxygen by volume.
Fit the delivery tubes to the sidearms of the U-tube with the special rubber connections. You have to make the delivery tubes from your glass tubing. The end of each delivery tube projects into an inverted test-tube of water in a deep dish. Support the test-tubes with cork-lined apparatus clamps. When all is ready, the carbon rods projecting from the stoppers are connected by conducting wire and crocodile clips to the battery. Two batteries connected in series form quicker results. Hydrogen gas forms at the cathode.

15.5.7 Electrolysis of aqueous salt solutions with variable voltage supply or 12 volt battery
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, Eo = + 0.82 V
4H2O + 4e- --> 2H2 + 4OH-, Eo = -0.41 V
So the minimum voltage for electrolysis of pure water = + 0.82 -(-0.41) = 1.23 V

15.5.13 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) + OH-

15.5.14 Electrolysis of copper (II) sulfate solution
Pass a current through copper sulfate solution. Pass the current for three or four minutes and examine the electrodes. Oxygen from the water is formed at the anode, but no hydrogen gas is evolved. Instead, the metal, copper, is deposited as a film on the cathode. The cathode has a deposit of copper on it, which can be wiped off.
15.5.15 Electrolysis of copper (II) sulfate solution, Faraday's laws
See diagram 15.5.15: Electrolysis of copper (II) sulfate solution
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 adjusted by rheostats to carry A. 1 amp, B. 0.5 amp and C. 0.5 amp. If electrodes have an immersed area of 8 × 5 cm, current of 1 amp corresponds to a current density of about 0.025 amp per cm2 and current of 0.5 amp corresponds to current density of about 0.012 amp 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 carrying A. 1 amp and open the switch in the circuit carrying B. 0.5 amp.
After twenty minutes, open the switch in circuit carrying C. 0.5 amp. Remove the cathodes then wash, dry and weigh them again.
The weight of copper carried across is proportional to current × 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.
15.5.16 Electrolysis of copper (II) sulfate solution, microscale electrolysis
See diagram 15.5.16: Electrolysis of copper (II) sulfate solution, microscale electrolysis
Microscale electrolysis allows very fine observation of changes during electrolysis.
1. 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.
2. 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 gas. 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.

15.5.19 Electrolysis of copper (II) sulfate solution, electrochemical equivalent of copper
See diagram 32.2.65: Electrolysis of copper (II) sulfate solution
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.

15.5.20 Electrolysis of tin (II) chloride solution, with an OHP
1. Pour 2 M tin (II) chloride (stannous chloride) solution into a Petri dish on an overhead projector. Focus on two parallel electrodes made of tin or lead. Pass 5 V of electric current and observe flakes of tin appearing on the cathode. Reverse the current to see the tin flakes dissolve, then appear on the anode.
2. Dissolve 113 g of tin (II) chloride dihydrate in 200 mL of concentrated acid. Add small pieces of metallic tin to the solution then add deionized water to make up the volume to 1 litre. This solution is an irritant so wear eye protection. For the electrolysis, use a fume cupboard to pass 2 to 5 volts through 40 mL of the solution using a carbon anode (+ terminal), e.g. graphite rod from a torch battery, and an iron cathode (-ve terminal), e.g. steel nail. Tin crystals form on the cathode and chlorine forms at the anode, so run the electrolysis just long enough to see the tin crystals starting to form.
At the cathode: Sn2+ + 2e- --> Sn

15.5.21 Electrolysis of  sodium sulfate solution
Use a car battery, large dry cells or a 2-12 volt transformer rectifier as a source of current. Use two copper wires as electrodes for the electrolysis of dilute sodium sulfate solution. Bubbles of gas (hydrogen gas) rise from one electrode. The other electrode is attacked. If instead of copper, the electrode is a short length of platinum wire or platinum foil, bubbles of oxygen will be produced. Water is decomposed into hydrogen gas and oxygen. Hydrogen gas and oxygen are obtained similarly by electrolysis of dilute
solutions of many common substances..

15.5.23 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)

15.5.24 Electrolysis of acids, 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- --> 2H2O + O2 + 4e-
Repeat the experiment with copper wire electrodes dipping into the acid solution.

15.5.25 Electrolysis of acids, hydrochloric acid
Pass a current through a beaker of hydrochloric acid. Note the gases formed, hydrogen gas and chlorine. The hydrochloric acid decomposes into its elements, hydrogen and chlorine, and the water is not affected. Hydrogen gas forms at the cathode and chlorine forms at the anode.

15.5.27 Electrolysis of potassium iodide solution, prepare iodine solution
Repeat the above experiment but three quarters fill the two small beakers with potassium iodide solution. Observe the brown iodine forming at the anode (positive electrode). When the solution has turned a pale brown, stop the current. Potassium hydroxide forms in the other beaker. Store and label the two solutions.
15.5.28 Chemical reaction forms electricity
See diagram 15.2.28: Magnesium / copper cell
Attach two wires to a light bulb. Attach a piece of magnesium ribbon to one wire . Attach a strip of copper foil to the other wire. Hold the bulb in your hand and dip the magnesium and copper into a beaker of dilute sulfuric acid or a solution of sodium hydrogen sulfate. Describe what you see. The bulb lights, and bubbles of hydrogen gas form on the copper strip. The chemical reaction between the magnesium and the acid causes a current of electricity to flow along the wire and light the bulb. The current also makes the hydrogen gas bubbles come out of the acid at the copper strip, although the copper itself does not react.

15.5.29 Electrolysis and mass transfer
Measure the current while transferring mass by plating copper to obtain a semi-quantitative determination of the Faraday experiment

15.5.30 Electrolysis and mass of sodium atom by electrolysis with an electrolytic rectifier
Electrodes of aluminium and lead in a saturated solution of sodium bicarbonate form a rectifier.

15.5.31 Electrolysis and oxidation of ferrous to ferric iron
Put ferrous iron in hot water with nitric acid and heat.

15.5.32 Electrolysis of sodium ions through glass
Sodium is plated on the inside of a lamp inserted into molten sodium nitrate!

15.5.33 Electric forge
Melt an iron rod cathode in a strong sodium sulfite solution.
15.5.12 Electrolysis of sodium chloride solution
See diagram: 3.69.1: Electrolysis of sodium chloride solution
In electrolysis of a melt, only two ions are present in the dissolved salt. However, in electrolysis of a solution the two ions from the slight dissociation of every 6 X 109 water molecules are also present. A solution of sodium chloride in water contains Na+(aq) H+(aq) Cl-(aq) OH-(aq) and H2O(l) < = > H+(aq) + OH-(aq)
Make an electrolysis apparatus from a wide mouth plastic bottle, two-holes stopper, and two carbon electrodes from the centres of 1.5 V dry cell batteries. Test the gases that form at the anode and the cathode. Repeat the experiment using 7% Na2CO3 with Fe electrodes.
Experiments
1. Electrolysis of a saturated sodium chloride solution, with a low voltage DC source of current
Reduction at the cathode
Both Na+ and H+ ions are attracted. The hydrogen ions are reduced by electron gain to form hydrogen molecules.
2H+ (aq) + 2e- --> H2 (g)
The Na+ ions are not reduced.
Oxidation at the anode
Both OH- and Cl- ions are attracted. The chloride ions are oxidized by electron loss to give chlorine molecules.
2Cl- (aq) --> Cl2 (g) + 2e-
The sodium chloride solution electrolyte becomes sodium hydroxide solution with the loss of chlorine and turns universal indicator purple.
Tests
1.1 Hydrogen gas and oxygen gas appear as bubbles because they have very low solubility in water.
1.2 Test for hydrogen gas with a lighted splint that produces a sudden squeaky pop sound. Also, pass
hydrogen gas into a solution of detergent, then hold a lighted splint near the bubble to ignite the gas with a soft sound.
1.3 Test for oxygen gas with a glowing splinter of wood that ignites.
1.4. Tests for chlorine
1.4.1 Hydrogen gas and oxygen gas are colourless and odourless, but chlorine is a poisonous green yellow gas with has a strong unpleasant odour so it can be detected by careful smelling.
1.4.2 Chlorine gas is more soluble in water, so it first dissolves then forms bubbles of gas when the solution is saturated with chlorine.
1.4.3 Chlorine turns wet blue litmus red then bleaches it white.
1.4.4 Chlorine causes a potassium iodide solution to become brown and a potassium iodide / starch solution to become dark blue.

2. Electrolysis of dilute sodium chloride solution
Dissolve a finger width of sodium chloride and add a finger width of litmus solution made red by a few drops of acid. Place the carbon rods in the solution, and connect them to the battery. The carbon rods should be kept well apart. Describe what happens in the liquid around each carbon rod. Around the cathode the liquid turns blue and bubbles of gas evolve. At the anode the liquid goes colourless and there are few bubbles. Notice the “swimming pool smell” of chlorine. Do not inhale this gas. When sodium
chloride solution is electrolysed, chlorine gas is evolved at the anode and much of it dissolves in the water causing the “swimming pool smell”. Chlorine is a bleaching agent, so it turned the litmus colourless near the anode. At the cathode, hydrogen gas is evolved and sodium hydroxide is formed turning the litmus blue.

3. 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.)
In electrolysis of a solution the two ions from the slight dissociation of every 6 × 109 water molecules are also present. A solution of sodium chloride in water, brine, contains Na+ (aq) H+ (aq) Cl- (aq) OH- (aq) and
H2O (l) < = > H+ (aq) + OH- (aq)
Make an electrolysis apparatus from a wide mouth plastic bottle, two holes stopper, and two carbon electrodes from the centres of 1.5 V dry cell batteries or pencil leads. Redox reactions occur at the electrodes. Test the gases that form at the anode and the cathode.

4. Electrolysis of dilute sodium chloride solution with a low voltage DC current
At the cathode, hydrogen ions are reduced more easily than sodium ions, so they will form hydrogen gas. At the anode, hydroxide ions are oxidized more easily than chloride ions in dilute solution, so they will form oxygen gas. The hydrogen gas and oxygen gas formed will be in a 2:1 ratio by volume because the net reaction is electrolysis of water. Local changes in pH There should be no net change in pH of the electrolyte but localized changes may occur. Loss of hydrogen ions at the cathode leaves the electrolyte
near the cathode with a net excess of hydroxide ions, basic solution. Similarly, loss of hydroxide at the anode leaves the electrolyte near the anode with an excess of hydrogen ions, acid solution. Detect these local changes of pH with solid bromothymol blue indicator (green in neutral sodium chloride solution, blue in basic solution, yellow in acidic solution). Repeat the experiment with a platinum or titanium anode and a fine copper wire cathode. The products of the electrolysis are the same.

5. Electrolysis of sodium chloride solution in a Petri dish
See diagram: 3.69.2: Electrolysis of sodium chloride
Dissolve 5 g of sodium chloride in 2 cc of water and add 3 drops of methyl orange indicator. Put a filter paper inside a plastic Petri dish then drop the solution onto the paper with a dropping pipette until the filter paper can hold no more solution. Attach the positive end of a 6 V battery to a lead wire with a crocodile clip to grip one end of the filter paper. Attach the negative end of the battery to a carbon electrode, e.g. a "lead" pencil. Rub the carbon electrode on the wet filter paper to make a mark. Hydrogen ions are attracted to the negative terminal so the mark is red. Repeat the experiment with universal indicator. The mark is red.
Repeat the experiment with other indicators.
6. Electrolysis of sodium chloride solution to prepare sodium hydroxide solution and chlorine
See diagram: 15.5.10: Prepare solutions of sodium hydroxide and chlorine
In this experiment the sodium hydroxide and chlorine are apart because they react with each other. Three quarters fill two small beakers with sodium chloride solution. Connect the beakers with a strip of absorbent paper. The paper soaks up the liquid and forms a connection between the liquids in the two beakers. Place a carbon rod in each cup and connect them to the battery. The carbon rods should not touch the wet absorbent paper. Close the circuit and let current to pass for one hour. Store and label the solutions in the beakers. The absorbent paper serves as a “bridge” between the two vessels, allowing the current to pass while the sodium hydroxide and chlorine form separately in the vessels. The chlorine dissolves in the water, as does the sodium hydroxide.

7. Electrolysis of sodium chloride solution to prepare sodium hypochlorite
If the sodium hydroxide and chlorine, formed in the electrolysis of sodium chloride solution, are allowed to mix, they form sodium hypochlorite, an important germicide and disinfectant. Prepare a saturated solution of sodium chloride in a beaker of water. Pass the current through the solution for two hours. Store and label the solution. Do not let the carbon electrodes touch.