School Science Lessons
Topic 15a Electroplating, electrolysis
Updated 2009-09-16
Please send comments to: J.Elfick@uq.edu.au
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Table of contents
3.59.0 Electrical conductivity
15.1.0 Electroplating
15.5.0 Electrolysis
3.59.0 Electrical
conductivity
3.59.1
Substances that conduct electricity
3.59.2
Electrical
conductivity of solids
3.59.3
Electrical conductivity of melted solids,
fused solids
3.59.4
Electrical conductivity of liquids
15.1.0
Electroplating
15.1.1 Faraday's first law
15.1.1.1 Test Faraday's first law
with copper and copper (II) sulfate solution
15.1.2 Electroplating,
copper plating
15.1.3 Electroplating, chromium
plating
15.1.4 Electroplating, nickel plating
15.1.5 Electroplating, silver plating
15.1.6 Electroplating, zinc plating
of copper
15.1.7 Electroforming with copper
15.1.8 Anodize aluminium
15.1.9 Silvering and desilvering, plating and
deplating silver
15.1.10 Electroplating copper, copper flashing
of iron
15.1.11 Lead tree and tin tree
15.1.12 Electroplating with silver
15.1.13 Cucumber pickle frying
15.1.14 Silver coulometer
15.5.0
Electrolysis
15.5.1 Electrolysis of fused sodium chloride
15.5.2 Electrolysis of melted lead (II) bromide
15.5.7 Electrolysis of aqueous salt
solutions
with variable voltage supply or 12 volt battery
15.5.10 Electrolysis of sodium
chloride to
prepare sodium hydroxide solution and chlorine
15.5.11 Electrolysis of sodium
chloride solution
to prepare sodium hypochlorite
15.5.13 Electrolysis of
concentrated sodium
chloride solution, Nelson cell
15.5.15 Electrolysis of copper (II)
sulfate
solution, Faraday's laws
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
15.5.19 Electrolysis of copper (II)
sulfate
solution, electrochemical equivalent of copper
15.5.20 Electrolysis of tin (II)
chloride
solution
15.5.21 Electrolysis of tin (II)
chloride, with
overhead projector or microscope
15.5.22 Electrolysis of silver
nitrate, with
overhead projector or microscope
15.5.23 Electrolysis of potassium
iodide
solution, electrolytic writing
15.5.25 Electrolysis of acids,
hydrochloric acid
15.5.27 Electrolysis of potassium
iodide
solution, prepare iodine solution
15.5.29 Electrolysis and mass
transfer
15.5.30 Electrolysis and mass of
sodium atom by
electrolysis with an electrolytic rectifier
15.5.31 Electrolysis and oxidation
of ferrous to
ferric iron
15.5.32 Electrolysis of sodium ions
through glass
15.5.33 Electric forge
15.5.01
Electrolysis
15.5.3 Electrolysis of water, conduction of water
15.5.4 Electrolysis of water (decomposition of
water) Hoffman electrolysis apparatus
15.5.5 Electrolysis of water, measure volume of
hydrogen gas generated
15.5.6 Electrolysis of water, Decomposition of
acidified water by electricity
15.5.8 Electrolysis of dilute sodium chloride
solution
15.5.9 Electrolysis of dilute sodium chloride
solution with a low voltage DC current
15.5.12 Electrolysis of sodium chloride solution
15.5.14 Electrolysis of copper (II) sulfate
solution
15.5.16 Electrolysis of copper (II) sulfate
solution, microscale electrolysis
15.5.24 Electrolysis of acids, acetic acid
solution
15.5.26 Electrolysis with carbon electrodes
15.5.28 Chemical reaction forms electricity
18.7.6 Dissolve chlorine in pool
water by
electrolysis,
3.59.1
Substances that conduct electricity
Insert two carbon rods each through a 13 mm one-hole stopper. Bind the
two stoppers together so that the wider end of one stopper (up) is next
to the narrower end of the the other stopper (up). The carbon rods in
the
stopper should be 5 min apart. Use crocodile clips to attach a
conducting wire between one battery
terminals and the bulb and the other battery terminal and one of the
carbon rods. The bulb should light when a current passes so lightly
touch both carbon rods with copper wire to make
the bulb light and show that the circuit is works. Prepare separate
beakers of sugar, sodium carbonate, sodium chloride and laundry starch.
Dip in the carbon rods in each beaker and record whether the bulb
lights up to show that the solution is conducting
electricity between the carbon rods. Add water to each beaker. Dip in
the carbon rods in each solution beaker and record whether the bulb
lights. Wash the rods thoroughly under the tap after dipping in each
solution. Note any signs of chemical reaction in the beaker. None of
the original solid substances conduct electricity. Sodium carbonate
solution and sodium chloride solution conduct electricity. These
solutions are electrolytes. Solutions of sugar, starch and methylated
spirit do not conduct electricity. They are non-electrolytes. Repeat
the experiment by testing dilute hydrochloric acid and dilute sodium
hydroxide. Acids, salts, and alkalis are electrolytes. When dissolved
in water to
form solutions or melted into liquids by heating, they conduct
electricity. Electrolytes are usually decomposed when electric
current passes through them, electrolysis. In electrolysis, the carbon
rod (electrode) connected to the negative
(-) terminal of the battery is the cathode, and the electrode connected
to the positive (+) terminal is the anode. Gases from the decomposition
of electrolytes may be seen as bubbles on the electrodes.
3.59.2 Electrical
conductivity of solids
See diagram 3.59: Electrical conductivity
apparatus
Use two carbon electrodes from torch batteries, a non-conducting
support for
the
electrodes, crocodile clips or crunched aluminium foil for connections,
light bulbs to show when current flows, and a 6 V dry cell power
source. Test the conductivity of solids by making a good contact
between
the cleaned surface of the solid and the two electrodes. Confirm that
metals and carbon conduct electricity. Test the conductivity of
non-metallic and crystals, e.g. calcite
(crystalline calcium carbonate) candle wax, copper
(II) sulfate-5-water, ethanedioic acid-2-water (oxalic acid) glass
rod, naphthalene, plastics, octadecanoic acid, sucrose (cane sugar)
sodium chloride crystals, sodium nitrate, sugar crystals, sulfur, wax.
None of these
solid compounds is a good conductor.
3.59.3 Electrical conductivity of melted solids,
fused solids
Be careful! Do not let the
carbon
electrodes ignite and burn.
Grip two carbon electrodes from used dry cell batteries with the
crocodile clips. Test the conductivity of the melt by dipping in the
electrodes. Wait for the electrodes to reach the same temperature. This
ensures that the electrodes are in contact with the liquid and not the
solidified melt. Scrape and clean the electrodes between each test.
1. Melt substances that are solids at room temperature, but heat very
gently, otherwise they may ignite and burn, e.g. candle wax, cellulose
acetate (acetate rayon) lead metal, lead bromide, naphthalene, nylon,
octadecanoic acid (stearic acid) polyethylene, polythene, Perspex,
potassium iodide (m.p. 682oC) sodium chloride, sodium
nitrate, solder, sulfur, tin metal. Melted solids vary in their
conductivity. Only molten metals, alkalis and salts are good
conductors. Sugar and sulfur are non-conductors.
2. Glass can be a conductor. Heat a glass rod until it becomes very
hot and begins to soften. Test the hot, soft part with the conductivity
apparatus. When molten, glass is a good conductor of electricity.
3.59.4 Electrical conductivity of liquids
Pure substances that are gases or liquids at room temperature are not
good conductors, but the liquid metal mercury is a good conductor.
1. Clean and dry the carbon electrodes between each test. To test the
conductivity of liquids, immerse the ends of carbon electrodes 3 mm
apart in acetone, copper (II) sulfate
solution, methylated spirit, liquid paraffin, olive oil,
peanut oil, sodium chloride solution, sugar solution, turpentine
(mineral turps) vinegar.
2. Test the conductivity of solutions, e.g. 2 M concentration
of the following:
2.1 Strong electrolytes, e.g. copper (II) sulfate,
hydrochloric acid, potassium hydroxide, sodium chloride, sodium
hydroxide, sodium nitrate, sulfuric acid.
2.2 Weak electrolytes:
ammonia solution, benzoic acid, ethanoic acid (acetic acid). Always
wash the electrodes thoroughly after testing each solution. Solutions
of acids alkalis and metallic salts are generally good
conductors. Solutions of sugar and alcohol are non-conductors.
Solutions of other types of substances in water and in other liquids
are generally non-conductors.
3. Test demineralized water for
conductivity. The bulb does not light. Very gradually stir small
crystals of sodium chloride into the water. Note any light from the
light bulb as the salt dissolves. Similarly test distilled water, tap
water and mineral water.
4. If a commercial conductivity meter is available,
nonelectrolytes show a very small current but a completely dissociated
strong electrolyte e.g. 0.1 M HCl, shows a current > 100 mA.
4.1 Dilute 5 mL of 0.1 M solution of:
4.1.1 HCl,
4.1.2 NaOH to 50 mL.
Test
each reactant solution then mix the two solutions and test half
the volume of the product solution. The conductivity of the
product solution is less than the conductivity of each of the reactant
solutions.
4.2 Test 5 mL of 0.1 M solutions of
4.2.1 acetic acid, HC2H3O2,
4.2.2 aqueous ammonia solution.
Test each
reactant solution then mix the two solutions and test half the volume
of the product solution.
4.3 Test 5 mL of 0.1 M solutions of:
4.3.1 H2SO4,
4.3.2 Ba(OH)2.
Add 3 drops of 0.1% thymol blue indicator
solution to the sulfuric acid solution then add drops of the 0.1
M Ba(OH)2 solution while stirring until the indicator
changes from pink to yellow to blue. Test the conductivity of the
product solution.
15.1.0 Electroplating
To electroplate something is to coat it by electrolytic deposition, e.g
chromium plating, silver plating. You can pass electric current through
an electric cell, called an
electroplating bath, to deposit one metal on another. Metal is taken
from the anode and deposited on metal articles that act as cathodes.
The electrolyte contains the metal to be deposited as ions. The object
to be plated is the cathode on an electrolytic cell. A salt of the
plating
metal is the electrolyte, e.g. chromium salt for chromium
plating.
15.1.1 Faraday's first law
Faraday's first law states that the amount of chemical change during
electrolysis is proportional to the charge passed, i.e. the quantity of
electricity passed. A coulomb is the quantity of electricity that
passes when one amp of current passes for one second. One Faraday (F) =
96,500 coulombs.
15.1.1.1 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 Electroplating, copper plating
See diagram 3.3.7: Copper plating
1. Dip a clean iron nail into copper (II) sulfate solution. The
quickly
becomes coated with a layer of copper.
2. Clean a key in dilute hydrochloric acid, wash with water and
polish with steel wool. The plating bath contains 70 g copper (II)
sulfate crystals dissolved in 500 mL of water, 25 mL of methylated
spirits, 15 mL concentrated sulfuric acid and grains of clear gelatine.
The plating current is 0.5 amps. If the plating current is too strong,
a spongy layer forms on the electrode and you can easily rub off the
layer of copper.
3. Use a 6 volt battery. Attach a copper electrode to the positive
terminal and attach a platinum electrode to the negative terminal.
Immerse both electrodes in copper (II) sulfate solution in a beaker.
Observe
any changes. Copper is plated onto the platinum electrode and is
corroded from the copper anode. After a few minutes, reverse the
connections and observe any changes. The copper that had plated onto
the platinum electrode was corroded from it and plated onto the copper
cathode.
4. Coat a non-conductor, e.g. wood, rubber or plastic, with
graphite powder. The copper wire in the non-conductor must contact the
graphite. Then use it as a cathode to electroplate a non-conductor.
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 Electroplating, chromium plating
Steel motor car bumper bars may be chromium plated
Electroplating, copper plating. Add concentrated sulfuric acid to
potassium dichromate to form red chromate ion, 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 Electroplating, 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 Electroplating, silver plating of copper or
nickel
Cutlery may be made of nickel but plated with silver. The letters
"EPNS" stamped into some "silver" spoons means "Electroplated Nickel
Silver". Use a silver anode and a copper or nickel cathode cleaned in
concentrated nitric acid until the surface is matted. The plating
current is 1 amp per cm2 for 10 minutes. The deposit is
dull, but becomes shiny after polishing. The electrolyte, plating bath,
contains 1 g silver nitrate in 20 mL water, or 1 litre of water
containing 40 g silver nitrate, 560 g potassium iodide and 2.8 mL
concentrated sulfuric acid.
15.1.6 Electroplating, zinc plating of copper
Use a mixture of zinc chloride, boric acid and ammonium chloride.
The reaction forms a complex zinc salt.
Prepare 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.
2. Degrease a 12 cm X 3 cm thin aluminium
strip by wiping with
propanone (acetone). Dip the lower half of the aluminium strip in 1.4
mol per litre sodium hydroxide until effervescence occurs indicating
removal of the aluminium oxide layer. Dip the cleaned half in nitric
acid to neutralize the sodium hydroxide. Dry the strip without touching
it and weigh.
Al2O3 (s) + 2NaOH (aq) + 3H2O (l)
-->
2NaAl(OH)4 (aq)
Al2O3 (s) + 2OH- (aq) + 3H2O
(l)
--> 2Al(OH)4- (aq) [Cleaning the oxide]
2Al (s) + 2NaOH (aq) + 6H2O (l) --> 2NaAl(OH)4-
(aq)
+ 3H2 (g)
2Al (s) + 2OH- (aq) + 6H2O (l) --> 2Al(OH)4-
+ 3H2 (g) [Reaction after oxide removed]
3. Line a 1 litre beaker with double
aluminium foil. Fill beaker with 2
mol per litre sulfuric acid at 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 X 3 cm,
then area = 3 X 3 X 2 (2 sides) = 18 cm2, so current needed
= 18 X 10 to 20 = 180 to 360 mA (0.18 to 0.36 A). Close circuit for
about 30 minutes but keep adjusting the rheostat to
keep the current constant.
At the anode: 2Al (s) + 3H2O (l) --> Al2O3
(s)
+ 6H+ (aq) + 6e-
At the cathode: 6H+ (aq) + 6e- --> 3H2
(g)
Combined equation: 2Al (s) + 3H2O (l) --> Al2O3
(s)
+ 3H+ (g)
Be careful! hydrogen 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.
15.1.10 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 Electroplating
with
silver
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
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
See diagram 3.68: 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 waterr
See diagram 15.5.3: Electrolysis of an
aqueous salt solution
See 3.59.1: 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) Hoffman
electrolysis apparatus
See diagram 3.69: Electrolysis apparatus
This experiment is also called "electrolysis of acidified
water", "Hoffman
electrolysis apparatus", and "decomposition of water".
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 that causes a sharp popping
sound.
Tests for oxygen with a glowing splint that re-ignites 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
Pass d.c. 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 Hoffman 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.
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. 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 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 is evolved 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
15.5.8 Electrolysis of dilute sodium chloride
solution
1. 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.
2. 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 X 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.
15.5.9
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 localised
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.
15.5.10 Electrolysis of
sodium chloride 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.
15.5.11
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.
15.5.12
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-
Tests
hydrogen gas and oxygen gas appear as bubbles because they have very
low
solubilities in water.
Test 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.
Test oxygen gas with a glowing splinter of wood that ignites.
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. Chlorine gas is more soluble in
water, so it first dissolves then forms bubbles of gas when the
solution is saturated with chlorine. Chlorine turns wet blue litmus red
then bleaches
it white. Chlorine causes a potassium
iodide solution to become brown and a potassium iodide / starch
solution to become dark blue.
The sodium chloride solution electrolyte becomes sodium hydroxide
solution with the loss of chlorine and turns universal indicator
purple.
2. 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- --> 2OH- + H2
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)
+ 2OH-
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
1. The electrolyte is a saturated solution of copper (II) sulfate + 5%
sulfuric acid. The copper voltameter has two clean copper electrodes
attached to the sides of a glass jar by clips fitted with terminals so
that the cathode can be removed and replaced in the same place. Connect
three similar circuits carrying currents 1. 2. and 3. adjusted by
rheostats to carry 1. 1 amp 2. 0.5 amp and 3. 0.5 amp.
If electrodes have an immersed area of 8 x 5 cm, currents of 1 A
corresponds to a current density of about 0.025 A per cm2 and
current of 0.5 A corresponds to current density of about 0.012
A per cm2. Wash and dry the cathodes then clean with emery
paper. Weigh the cathodes and place into the three circuits. Close the
three switches simultaneously. After ten minutes, open the switch in
the circuit 1. carrying 1 A and open the switch in circuit 2. carrying
0.5 A. After twenty minutes, open the switch in circuit 3. carrying 0.5
A. Remove the cathodes then wash, dry and weigh them
again.
The weight of copper carried across is proportional to current x time.
The first law of electrolysis, discovered by Michael Faraday, states:
The mass of substance liberated during electrolysis is proportional to
the charge passed. If mass/charge = the electrochemical equivalent
constant of the substance, Faraday's second law states: The amount of
chemical produced in different substances by a quantity of electricity
is proportional to the electrochemical equivalent constant of the
substance.
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
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.
15.5.21 Electrolysis of tin (II) chloride, with
overhead projector or microscope
1. Pour 2 M tin (II) chloride (stannous chloride) solution into a
flat transparent dish on an overhead projector. Focus on two parallel
electrodes
made of tin or lead, or use solder sticks. Pass 5 V of electric current
and note flakes of tin appearing on the cathode. Reverse the current to
see the tin flakes dissolve then appear on the anode.
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- --> 2H20 + 02 + 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..