School Science Lessons
Topic 15a Electrolysis, electroplating
2012-05-10 SPwp
Please send comments to: J.Elfick@uq.edu.au
Table of contents
15.5.0 Electrolysis
15.1.0 Electroplating
15.5.0 Electrolysis
15.5.28 Chemical reaction forms electricity
18.7.6 Dissolve chlorine in 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 volt battery
15.5.1 Electrolysis of fused sodium chloride
15.5.2 Electrolysis of melted lead (II) bromide
15.5.14 Electrolysis of copper (II) sulfate solution
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
15.5.17 Electrolysis of copper (II)
sulfate solution with copper and platinum electrodes
15.5.18 Electrolysis of copper (II)
sulfate solution with copper electrodes
15.5.23 Electrolysis of potassium iodide solution,
electrolytic writing
15.5.27 Electrolysis of potassium iodide solution,
prepare iodine solution
15.5.22 Electrolysis of silver nitrate, with overhead
projector or microscope
15.5.12 Electrolysis of sodium chloride solution
15.5.13 Electrolysis of sodium chloride
solution, Nelson cell
15.5.32 Electrolysis of sodium ions through glass
15.5.20 Electrolysis of tin (II) chloride solution
15.5.3 Electrolysis of water, conduction of water
15.5.4 Electrolysis of water, decomposition of
water, Hofmann voltameter
15.5.6 Electrolysis of water, Decomposition of
acidified water by electricity
15.5.5 Electrolysis of water, measure volume of
hydrogen gas generated
See
pdf Electrolysis, Pencil and Paper Gland Electrolyser
15.5.26 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 with 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
15.1.6 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
Test 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 with copper
Make a wax impression by pressing a key into soft wax in a crucible.
Remove the object. Insert copper wire in the wax. Dust the impression of
the key with bronze powder that should also contact the wire. Copper plate
for an hour. Remove from mould and dry. Fill with molten solder. .
15.1.8 Anodize aluminium
See diagram 15.1.8.1: Anodize aluminium
Use anodizing to remove oxides from the surface of objects then coat
them with a hard and thick oxide layer which can absorb dyes or hardening
agents. Anodized dyed aluminium is used for door and window frames, saucepan
lids, and wherever bright reflective surfaces are required. Anodized steel
may be used in small objects for marine construction. Anodized titanium
is used for shiny jewellery.
1. Put two pieces of aluminium foil in hot sodium hydroxide solution
to remove any aluminium oxide layer. Rinse in water, then nitric acid,
then water. Use the two pieces of aluminium as electrodes in 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.
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.
2. 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]
3. 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.
15.1.9 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.
15.5.0 Electrolysis, voltameter, coulometer
Chemical reactions in
a liquid, an electrolyte, caused by passing of electric current is called
electrolysis. The chemical reactions are usually the decomposition of a substance
by the application of electric current. Electrolytes are acids, bases or
salts dissolved in water. Electric current enters or leaves the electrolyte
through conductors called electrodes. The electrode joined to the positive
terminal of the battery is called the anode. Conventional electric current
passes from the positive terminal of the battery to the anode. The electrode
joined to the negative terminal of the battery is called the cathode. Conventional
electric current passes from the cathode to the negative terminal of the
battery. An electrolytic cell consisting of a container, electrolyte and
electrodes is called a voltameter (not "voltmeter"!). Such a cell used to
measure electric charge is called a coulometer. Electrolysis uses a source
of electricity to break apart an ionic compound. In an electrolytic cell,
an external electricity source, e.g. a battery, forces electrons around
the circuit away from the negative terminal of the battery and towards the
positive terminal of the battery. Electric current as ions is carried in
the electrolyte within the electrolytic cell. Positive ions, called cations,
are attracted to the negative cathode and negative ions anions are attracted
to the positive anode. Oxidation as loss of electrons from the ions in solution
to the electrode occurs at the anode. Reduction, as gain of electrons from
the electrode to the ions in solution, occurs at the cathode. If the electrolyte
is a salt consisting of a metal and non-metal, the metal precipitates at
the cathode and the non-metal precipitates at the anode. Electrolysis is
used to reduce metal ores to the metal.
15.5.1 Electrolysis of fused sodium chloride
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
15.5.2 Electrolysis
of melted lead (II) bromide
This experiment is not advisable in schools.
See diagram 3.2.68: Electrolysis of a melt
apparatus | See 1.13a: Simple fume hood
Electricity is carried through a solution by ions. At the electrodes,
these or other ions are neutralized and discharged as neutral atoms or molecules.
Do this experiment in a fume cupboard,
fume hood. Be careful! A hot molten solid is formed and choking poisonous
fumes are given off.
Lead (II) bromide has a comparatively low melting point salt at 373oC.
Melt the lead (II) bromide in a crucible and pass 12 V through it with carbon
electrodes. The only ions in this melt are bromide ions and lead ions. Lead
has both a lower melting point, 328oC, and a greater density
than lead (II) bromide, so it appears as a melt at the bottom of the crucible.
Note the small globule of lead at the negative electrode, the cathode, after
10 minutes of electrolysis. Decant the molten lead bromide carefully into
another crucible. Note the bromine gas, b.p. 59oC, at the positive
electrode, the anode. The electric current has split crystalline lead bromide
into bromine gas and lead metal.
Reduction at the cathode: Pb2+ (l) + 2e- -->
Pb (l)
Oxidation at the anode: 2Br- (l) - 2e- --> Br2
(g)
Summary equation: PbBr2 (l) --> Pb (s) + Br2
(g)
15.5.3 Electrolysis
of water, conduction of water
See diagram 15.5.3: Electrolysis of an aqueous
salt solution
See 15.59.1: Substances that conduct
electricity
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.
15.5.4 Electrolysis
of water, decomposition of water, Hofmann voltameter
See diagram 3.2.69: Electrolysis apparatus
| See diagram 15.5.4: Hofmann voltameter
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. Tests 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. 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.]
15.5.5 Electrolysis
of water, measure volume of hydrogen gas generated
1. 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 coulombmeter 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.
2. 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.
15.5.6 Electrolysis
of water, decomposition of acidified water by electricity
See diagram: 15.5.6: Decomposition of acidified
water
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 side-arms. 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 side-arms 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 side-arms 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
15.5.12 Electrolysis
of sodium chloride solution
See diagram: 3.69.1: Electrolysis of sodium
chloride solution
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. Hydrogen gas and oxygen gas appear as bubbles because they have very
low solubility in water.
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.
3. Test for oxygen gas with a glowing splinter of wood that ignites.
4. Tests for chlorine
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.
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.
4.3 Chlorine turns wet blue litmus red then bleaches it white.
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 saturated sodium
chloride solution
Use a U-tube with carbon electrodes that dip into the solution. Add some
6 M hydrochloric acid. Fill the U-tube to 2 cm from the openings. Insert
a two holes rubber stopper in each opening of the U-tube for the electrode
and for a delivery tube for gas collection. Use 6 to 12 V current. Higher
voltage speeds the reaction. Test the liquid at each electrode with litmus
paper. At the anode, the reaction produces chlorine instead of oxygen:
2Cl-
(aq) --> Cl2 (g) + 2e-
At the cathode, the reaction forms hydrogen gas from water molecules.
Sodium ions collect in the solution but are not discharged so the cathode
is surrounded by sodium hydroxide solution.
H2O + 2e- --> OH- + H2
6. 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.
7. 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.
8. 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.
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.
Attach fine copper wire to platinum wire and pass the end of the wire
under the lid of a Petri dish to form the anode. In bright light, clean the
tip of a piece of the fine copper wire, inspect it with a magnifier then
pass the end under the lid of the Petri dish to form the cathode. Tape the
electrode wires to the bottom of the Petri dish with tips separated by 5
millimetres and tape the electrode wires to the side of the Petri dish where
they pass over the sides. Put two drops of concentrated copper (II) sulfate
solution (10 g to 100 mL water) in the Petri dish so that the tips of both
electrodes are touching the solution. Place an extra two drops of copper
(II) sulfate solution aside to compare colour change. Put specks of solid
copper (II) oxide in the solution between the tips of the electrodes. Spread
the specks to form a continuous band in the solution between the electrodes.
Put the lid on the petri dish and connect the electrodes to a 3V source of
direct current. Connect the platinum anode to the positive terminal. Connect
the copper cathode
to the negative terminal.
Use a magnifying glass to observe
changes around the electrodes and the specks of copper oxide. When the circuit
is closed, deposits of copper appear on the cathode but the rate of deposition
later changes. The grains of copper oxide start to disappear into the solution.
When the growth of copper deposited on the cathode reaches a grain of copper
oxide, the coppers is deposited very rapidly around the grain. Note the
bubbles around the anode and later around the cathode. The bubbles may stream
from one electrode towards the other electrode. Hold a lighted taper above
bubbles appearing at the electrodes. Note the popping noise indicating hydrogen
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.
15.5.17 Electrolysis
of copper (II) sulfate solution with copper and platinum electrodes
Refining removes impurities from metals by electrolysis to get pure metals.
The cathode is a thin sheet of copper. The electrolyte is copper (II) sulfate
solution. The anode is the impure copper to be refined. During electrolysis
the pure copper leaves the anode and is deposited on the cathode leaving
the electrode as a mass of impurities. The equations show how impure copper
is purified by the electrolysis of a copper (II) sulfate solution in which
the impure copper is the anode and a sheet of pure copper is the cathode.
The anode corrodes and pure copper is deposited on the sheet of pure copper.
Attach a copper electrode to the positive terminal of a 6 V battery and
a platinum electrode to the negative terminal. Immerse both electrodes
in copper (II) sulfate solution in a beaker. Copper plates on the platinum
electrode and the copper anode corrodes.
Cu2+ + 2e- --> Cu
Reverse the connections. The copper plated on the platinum electrode
corrodes from it and plates on the copper cathode.
Cu2+ + 2e- <-- Cu
15.5.18 Electrolysis of
copper (II) sulfate solution with copper electrodes
See diagram 3.69.4: Electrolysis on an overhead
projector
Place a transparent dish half full of a 10 g in 100 mL copper (II) sulfate
solution on an overhead projector. Cut one end of a 1 cm wide strip of
copper foil into a sharp angle and hang it on one side of the flat transparent
dish with the sharp end under the copper (II) sulfate solution. Peel the
plastic cover off one end of a short length of electric wire then disperse
the thin copper wires. Hang the electric wire on the other side of the dish
with the thin copper wires immersed in the solution. Connect the copper
strip to the
positive terminal of a 12 V d.c. supply and connect the thin
copper wires to the negative terminal. Adjust the distance between the two
electrodes. No bubbles appear at the electrodes.
When electric current passes,
note the changes to the electrodes. The sharp end of the anode gradually
dissolves. At the cathode,
copper deposits on the thin copper wires, like
a branching tree. The cathode is covered with a fresh layer of copper. The
anode looks dull. The weight of copper lost by a pure copper anode equals
the weight of copper gained by the cathode, but the concentration of the
copper (II) sulfate electrolyte remaining the same. When using copper (II)
sulfate solution and copper electrodes, the graph of current against voltage
is a straight line, so the solution acts as an electrical conductor and Ohm's
law applies.
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
1. Pour some 2 M tin (II) chloride solution into a Petri dish on an overhead
projector. Focus on two parallel electrodes made of tin or lead. Pass about
5 V of electric current and observe flakes of tin appearing on the cathode.
Reverse the current to see the tin flakes dissolve.
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
15.5.22 Electrolysis of
silver nitrate, with overhead projector or microscope
Observe the electrolysis of silver nitrate solution under a microscope
or with an overhead projector. Use a 2 mm wide strip silver cathode, a platinum
anode, a 2M solution of silver nitrate and a power source less than 2 V.
Attach the electrodes to a microscope slide with adhesive tape leaving 1.5
mm between the tips. Use a dropper to put 2 drops of 2M silver nitrate solution
between the electrodes. Turn on the power and increase it slowly with a
potentiometer. Crystals of silver form around the cathode depending
on
the voltage and the shape of the cathode. Black Ag (I) Ag (III) O2
and some bubbles, probably oxygen, appear around the anode. Reverse the
circuit to watch reversal of the electrolytic reactions. In this reaction
oxidation of Ag (I) to Ag (III) occurs more easily than oxidation of hydroxide
to oxygen.
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.
15.5.26 Electrolysis with
carbon electrodes
1. Potassium iodide: Iodine forms at the anode (+ ve) and hydrogen gas
forms at the cathode (- ve).
2. Zinc sulfate: Oxygen forms at the anode (+ ve) and zinc forms at the
cathode (- ve).
3. Lead acetate: Oxygen forms at the anode (+ ve) and lead forms at the
cathode (- ve).
4. Copper (II) chloride: Chlorine forms at the anode (+ ve) and copper
forms at the cathode (- ve).
5. Copper (II) sulfate: Oxygen forms at the anode (+ ve) and copper forms
at the cathode (- ve).
6. Sodium chloride (concentrated): Chlorine forms at the anode (+ ve)
and hydrogen gas forms at the cathode (- ve).
7. Sulfuric acid (dilute): Oxygen forms at the anode (+ ve) and hydrogen
gas forms at the cathode (- ve).
8. Sodium hydroxide (dilute): Oxygen forms at the anode (+ ve) and hydrogen
gas forms at the cathode (- ve).
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.