School Science Lessons
(topic15.html)
2024-11-28
Electrochemistry
Contents
15.6.0 Cathodic protection
15.1.0 Corrosion
15.2.0 Electrolytes, electrolytic
15.3.0 Oxidation and reduction, redox reactions
15.4.0 Oxidizing agents and reducing agents
15.5.0 Rusting
15.1.0 Corrosion
Corrosive substances can destroy living tissue on contact.
Corrosive to metals: A substance or a mixture that by chemical action will materially damage or destroy metals.
15.1.1 Corrosion
15.1.2 Galvanized iron, galvanized steel
15.1.3 Tin plate
Experiments
15.1.4 Clean tarnished silver
15.1.5 Corrosion of alloys, restore bronze coins
15.1.6 Corrosion of aluminium
15.1.7 Galvanic Series
15.1.7a Standard electrode potential
15.1.8 Oxidation can affect air pressure
15.2.0 Electrolytes, electrolytic
15.2.1 Electrolytes
15.2.2 Electrolytic cells
15.2.3 Conductivity of solutions of electrolytes
15.2.4 Electrolytes in the blood and urine
Experiments
15.2.5 Strong electrolytes
15.2.6 Weak electrolytes, ammonia solution.
15.2.11 Tests for alcohol, breath tests
15.3.0 Oxidation and reduction
Oxidation is a loss of electrons
Reduction is a gain of electrons
15.3.1 Oxidants (oxidizing agents), (Safety)
15.3.2 Redox reactions, oxidation number
15.3.3 Reductants (reducing agents), (Safety)
15.3.4 When oxidation occurs:
15.3.5 When reduction occurs:
15.4.0 Oxidizing agents and reducing agents
15.4.1 Manganese oxidising agents
15.4.2 Oxygen as an oxidizing agent
Experiments
15.4.3 Bromine as an oxidizing agent
15.4.4 Chlorine as an oxidizing agent
15.4.5 Concentrated nitric acid as an oxidizing agent
15.4.6 Hydrogen peroxide as an oxidizing agent
15.4.7 Hydrogen sulfide as a reducing agent
15.4.8 Nitrous acid as oxidizing agent or reducing agent
15.4.10 Sulfurous acid as a reducing agent.
15.4.12 Gummi bears (gummi fruit) with potassium chlorate
15.4.13 Potassium chlorate and potassium persulfate as oxidizing agents
15.4.14 Potassium chlorate in pyrotechnic flash powders
15.4.15 Sparklers, potassium chlorate
15.4.16 Potassium dichromate as an oxidizing agent
15.4.17 Tests for oxidizing agents
15.5.0 Rusting
15.5.1 Rusting
Experiments
15.5.2 Conditions necessary for rusting
15.5.3 Electrode potential order of metals
15.5.4 Metals can prevent rusting
15.5.5 Need for oxygen for rusting
15.5.6 Oxygen gas combines with iron during rusting
15.5.7 Rusting of iron wire
15.5.8 Rusting of steel wool.
15.5.9 Rusting paper clips
15.5.10 Stainless steel
15.1.1 Corrosion
Corrosion is a general term for the environmental attack of materials including iron based materials.
So the corrosion product of aluminium is not called rust, but by its chemical name.
The corrosion product of iron is called rust, or iron oxides to give rust its chemical name.
Corrosion refers to the unwanted oxidation of metals.
Corrosion is an electrochemical reaction between a material and its environment that results in the deterioration or destruction of the material.
Metallic corrosion requires at least one metal and an electrolyte, water, soil, concrete or a moist atmosphere.
Metallic corrosion is the release energy stored during refining as electric current.
Galvanic corrosion occurs with at least two electrically-connected metals where one metal corrodes (anode) instead of another (cathode).
A simple cell consists of a zinc anode plate and a copper cathode plate immersed in an electrolyte solution.
When the plates are electrically connected, the zinc plate dissolves (corrodes) to form positively charged zinc ions in the solution and electrons migrate from the anode to the copper cathode across the metallic bridge.
Simultaneously, oxygen dissolved in the solution is also consumed at the cathode.
Ions migrate through the solution and combine with other ions to complete the corrosion reaction.
If a voltmeter were placed in the circuit between the anode and the cathode, the difference in energy levels would be measured as a DC voltage.
Galvanic anodes, as distinct from impressed current anodes, do not require an external power source so are a cheap and simple method of protecting against corrosion.
Galvanic electrodes are cast from high purity primary magnesium, aluminium and zinc ingot, with additional "activating" elements strictly controlled according to international standards.
A metal in an electrolyte (water, sea water, soils, concrete) generates an electrical current depending on the type of electrolyte.
Aluminium and zinc anode alloys are used in sea water, very conductive (low resistance) electrolyte.
Magnesium alloy anodes are used in fresh water that is much less conductive (high resistivity).
Any metal will usually corrode ten times faster in sea water than fresh water, because the resistivity of salt water is 0.25 ohm metres, while fresh water is typically 50 ohm metres.
The aim of cathodic protection is to shift the natural voltage of a metal in a negative direction to a point at which corrosion is significantly reduced.
Steel has a natural potential of -5OO mV in sea water (with respect to a standard silver / silver chloride reference electrode, e.g. "Rust Seeker".
If the potential can be shifted to about -800 mV, corrosion of the steel is significantly reduced.
15.1.2 Galvanized iron, galvanized steel
Zinc is used as a protective coating of iron for galvanized iron sheets and galvanized screws and bolts.
If the coating is scratched, in the zinc iron / rain water cell the zinc corrodes to protect the iron.
Also, blocks of zinc are attached to iron ships, bridges and wharfs, called sacrificial anodes.
In this sacrificial corrosion, the zinc corrodes away to protect the less active iron.
Nowadays, "Zincalume" steel outperforms galvanized steel.
This zinc / aluminium alloy-coated steel is composed of 55% aluminium, 43.5% zinc and 1.5% silicon, which provides a superior coating for steel.
The aluminium in the coating provides barrier protection to the steel and the zinc component provides sacrificial cut edge protection.
This double protection mechanism is responsible for Zincalume steel's superior corrosion performance.
15.1.3 Tin plate
Iron is coated with tin to make tin plate for tin cans, beverage cans, and jam tins.
However, if the tin is scratched, the iron corrodes more rapidly than if it were not covered by tin.
15.1.4 Clean tarnished silver
Silver is not readily oxidized, but the surface of silver can react with sulfur in the air to form black silver sulfide.
For example, silver spoons used for eating eggs that contain sulfur proteins change from shiny and bright to black and matt.
Polish off the silver sulfide or dissolve it using a household "silver dip" that contains ammonia or thiourea, but in each case some silver is lost.
Save the silver by using the following oxidation reduction reaction that reverses the corrosion process.
However, some jewellery designers deliberately create a black patina on sunken surfaces as a background contrast to bright silver surfaces.
The jewellers soak the jewellery in potassium sulfide, liver of sulfur, and later buff polish the silver surfaces.
Never try to clean silver with household bleach, because a hard coat of oxide forms that is very difficult to remove using the methods below.
1. To clean the silver, put a sheet of aluminium in the bottom of a beaker.
Put the silver to be cleaned on the aluminium and add baking soda solution (sodium hydrogen carbonate).
Warm the solution.
The sulfur transfers to the aluminium to form aluminium sulfide and the silver becomes shiny again.
2. Clean tarnished silver with aluminium foil.
Add 10 g of sodium bicarbonate (NaHCO3, baking soda) to hot water in a plastic container.
Wrap the tarnished silver in aluminium foil and immerse it in the solution for hours until the silver sheen is restored.
The sodium bicarbonate dissolves any aluminium oxide on the aluminium surface.
3Ag2S (s) + 2Al (s) --> 2Al3+ (aq) + 3S2- (aq) + 6Ag (s)
Rub the tarnished silver with "Brasso" or toothpaste, not the gel-type toothpaste, or buff polish the surfaces.
Soak the tarnished silver in dilute ammonia solution, cloudy ammonia.
Soak the tarnished silver in borax and soap solution in hot water.
Make pans containing hot baking soda solution to clean tarnished silverware, wrap steel wool to prevent rusting, reflect heat up through ironing board cover.
3. Line a metal cake pan with aluminium foil and fill with enough water to cover the silverware.
Add 30 mL baking soda per litre of water.
Heat the water above 65oC, but not boiling.
Place the tarnished silverware in the pan so it touches the aluminium foil.
The hydrogen produced by boiling baking soda combines with the sulfur in the tarnish to removing stains.
4. The bright, shiny surface of silver gradually darkens and less shiny, because silver reacts with substances containing sulfur in the air to form black silver sulfide.
The tarnish on silver can be removed by polishing, but this method removes some of the silver.
The method of dipping tarnished silver into a special liquid then rubbing also removes some silver.
The best method is the following electrochemical reaction:
Line the bottom of a pan with aluminium foil and put the tarnished silver object on the aluminium.
Heat water to boiling and add 125 mL of baking soda in 2 L of water.
Leave the mixture to froth then put it into the pan to cover the tarnished silver.
The tarnished silver and aluminium must remain in contact during the reaction.
If the tarnish does not disappear immediately, reheat the baking soda and water mixture and repeat the treatment.
Aluminium has greater affinity for sulfur than silver, so sulfur atoms are transferred from silver to aluminium, to free the silver metal and form aluminium sulfide.
3Ag2S + 2Al --> 6Ag + Al2S3
silver sulfide + aluminium --> silver + aluminium sulfide
The aluminium sulfide formed may stick to the aluminium foil or form pale yellow flakes in the bottom of the pan.
15.1.5 Corrosion of alloys, restore bronze coins
Brown "copper coins" are usually alloys of zinc and tin in copper.
"Silver" coins are alloys of nickel in copper.
Some "gold" coins are alloys of aluminium and nickel in copper.
Corrosion is common in alloys if the metals are not evenly mixed.
Old coins and statues made of copper alloys and other copper materials exposed to moist air are often covered with blue-green verdigris that is basic copper (II) carbonate, CuCO3.Cu(OH)2.H2O.
New "copper" coins are shiny, but they soon lose their shine and become a dark copper colour, because of a layer of black copper (II) oxide.
Old copper coins may be very black between the raised areas for the numbers.
Experiments
1. Put drops of vinegar on a copper coin.
Leave the coin until the liquid is evaporated.
Green blue crystals are left on the coin surface.
Scrape off the crystals and wash the coin.
The coin now looks shiny, because black copper (II) oxide is removed.
Use dilute hydrochloric acid to make "new" shiny coins.
2. Dissolve 5 g table salt in 50 mL of vinegar in a plastic beaker.
Suspend half a dull "copper" coin in the solution.
The dipped half become shiny, because the coating of copper oxide has been removed.
15.1.6 Corrosion of aluminium
Put a piece of aluminium foil in water.
Put a copper coin on the foil and leave it for some days.
A simple aluminium /copper cell forms and a small electric current can be detected with an ammeter.
The aluminium foil has holes where the coin lies on it.
The water appears cloudy, because of the fine particles of aluminium released during corrosion.
15.1.7 Galvanic Series
The rate of corrosion is affected by formation of electric cells.
All metals have an energy level that can be measured as its natural voltage within a particular electrolyte, e.g. sea water.
The galvanic series table in sea water from anodic actively corrode to cathodic noble passive (more resistant to corrosion) is as follows:
Mg anode alloys, Al anode alloys, Zn anode alloys, Al alloys, Cast iron, C steel (mild steel), Cu alloys (brass / bronze), Cupronickels, Cu, Ni, stainless steel (active), Ag, Ti, stainless steel (passive), Pt, Au, C (graphite).
Metals at the top of the galvanic series have a large voltage difference when connected to metals at the bottom so the rate of corrosion and loss of magnesium, aluminium and zinc on comparative surface areas would be rapid.
Experiment
Thoroughly clean short narrow strips of the metals copper, magnesium, tin, zinc, and five pieces of pure iron wire.
Twist a piece of iron wire tightly around each of the other metals.
Into five clean beakers place about equal volumes of tap water.
Place the single piece of iron wire in the water in one beaker and place one of the twisted pairs of metal strips in each of the other beakers.
Record your observations after one hour, one day, one week.
1. Cu + Fe, 2. Mg + Fe, 3. Sn + Fe, 4. Zn + Fe, 5. Fe.
15.1.7a Standard electrode potential
A simple method of calculating standard electrode potential,
Standard electrode potential.
The photograph shows a petri dish with a filter paper cut as seen.
In the middle of the filter paper was poured just enough salt bridge solution to soak but not run to the ends of the filter paper.
At the ends of each ‘arm’ of the filter paper is a piece of metal and it is soaked just enough to meet the salt bridge solution in its corresponding sulfate solution,
(i.e. copper metal and copper sulfate).
When ready, hold the lead (banana plug) to each piece of metal and take the readings.
The results are comparable to full scale half cells.
Note the strip of the metal folded it around the edge of the petri dish to allow the alligator clips to be clipped onto the metal, instead of just touching it.
All the pieces of metal are cut to a similar size.
The folded pieces are 0.5 cm wide and long enough to make for easy folding.
The volume of electrolytes/salt bridge is just enough to wet the filter paper as much as is needed.
Wet each arm with a 0.1M solution and then added the salt bridge into the middle before the ends came together.
The small pieces of each metal touch the solution soaked into the filter paper.
The apparatus acts like a normal half cell.
If one electrode is placed on the magnesium and another on the the copper, gives readings for the magnesium/copper half cell.
Then lift the electrode off the copper metal and place it on the zinc metal and get the readings for the magnesium-zinc cells.
Again, lift off the zinc and place it onto the lead for magnesium/lead cell.
Do the same process for all the other combinations.
So long as the salt bridge keeps the other solutions separate, gives individual half cells.
See photo Standard Electrode Potential
15.1.8 Oxidation can affect air pressure
Wash a small piece of steel wool in methylated spirit to remove any grease.
When it is dry, put it in a test-tube with a one-hole stopper fitted with a 40 cm length of glass tubing.
Clamp the test-tube with the end of the glass tubing under water.
Note the level of the water in the tubing at the start of the experiment and after one hour and two hours.
Water rises up the tubing as oxygen is used to form rust.
15.2.1 Electrolytes
Electrolytes may be acids, bases or salts dissolved in water.
Electrolytes are liquids containing ions that may be decomposed by electrolysis, e.g. the solution in a car battery.
An electrolyte is a solution contains ions that are electrically conductive through the movement of those ions, but not electrically conductive by moving electrons.
Electrolytes form ions when dissolves in water.
Living cells, blood and other tissues contain electrolytes.
O.R.S. (oral rehydration salts), taken to treat diarrhoea and traveller's gastro-enteritis in children and adults, contains balanced mix of glucose and electrolytes.
15.2.2 Electrolytic cells
Electrolytic cells use electrons from an external source.
The anode is the positive terminal of an electrolytic cell through which electrons leave and conventional current enters.
The cathode is the negative electrode of an electrolytic cell through which electrons enter and conventional current leaves.
If reverse voltage is applied to a Daniell cell, electrons enter the zinc cathode and zinc is deposited on it from solution.
Also, electrons leave the copper anode and copper goes into solution.
Zn2+ (aq) + 2e- --> Zn (s)
Cu (s) --> Cu2+ (aq) + 2e-
Cu (s) + Zn2+ (aq) --> Zn (s) + Cu2+ (aq)
The electrolytic cell is used to break up chemical compounds by electrolysis, e.g. the decomposition of water, using a Hofmann voltameter.
(August Wilheim von Hofmann, 1818-1892, Germany)
15.2.3 Conductivity of solutions of different electrolytes
An electrolyte can conduct an electric current in the fused state, or in solution, and it is decomposed while conducting the current.
Electrolytes dissolve in water to produce solutions that conduct electric current.
As the concentration of the electrolyte in solution increases, the conductivity of the solution increases.
A strong electrolyte breaks up almost entirely when it dissolves to produce an aqueous solution.
Water is a very weak electrolyte and a poor conductor of electricity so some electrolyte must be dissolved in it to increase its conductivity.
15.2.4 Electrolytes in the blood and urine
In medical use electrolyte refers to the ions.
So serum electrolyte refers to sodium, potassium or chloride ions that function in cardiac rhythm, skeletal muscle contraction and nerve transmission.
The level of bicarbonate ion is important for the acid-base balance in the blood.
The urine electrolytes, sodium and potassium, indicate electrolyte balance and how hormones affect the function of the kidney.
15.2.5 Strong electrolytes
Strong electrolytes completely ionize (dissociate), in solution.
1. Add drops of sodium hydroxide solution to each of three separate solutions of copper (II) sulfate, copper (II) chloride, copper (II) nitrate.
Observe the blue precipitate in each case.
These solutions contain only the copper (II) ion in common, so assume that this ion was responsible for the formation of the precipitate.
2. Add drops of barium chloride solution to separate solutions of copper (II) sulfate, sulfuric acid and sodium sulfate
In each case you can attribute the result the presence of the sulfate ion.
3. Add drops of ferric chloride solution to separate solutions of sodium hydroxide, potassium hydroxide and calcium hydroxide.
These experiments with solutions of strong electrolytes suggest that the properties of such solutions are the sum of the properties of the ions present.
The properties of copper (II) sulfate solution are made up of the properties of the copper (II) ion and the sulfate ion.
The copper (II) ion, Cu2+, causes the blue-green colour of the solution, and forms precipitates when other substances are added.
The sulfate ion contributes no colour, but forms precipitates with many other ions, such as Ba2+, when these are added to copper (II) sulfate solution.
15.2.6 Weak electrolytes, ammonia solution
Weak electroytes do not completely ionize, dissociate, in solution.
Smell very carefully a bottle containing some dilute ammonia solution.
The smell of ammonia suggests the presence of ammonia molecules that must have come from the solution.
Add a few drops of iron (III) chloride to a little ammonia solution.
From the results of a previous experiment with iron (III) chloride, the brown precipitate obtained confirms the presence of the hydroxide ion in ammonia solution.
Thus ammonia solution has properties due not only to the ions that are present, but also, because of ammonia molecules.
The properties of solutions of weak electrolytes are made up from the properties of the unionized molecules and the properties of the ions produced from them.
In such solutions, the ions and molecules are in equilibrium with each other.
15.2.11 Tests for alcohol, breath tests
1. Breath test for alcohol using potassium dichromate
The breath after drinking alcoholic beverages contains ethanol vapour, which can be oxidized by orange potassium dichromate (K2Cr2O7), that is reduced to green chromium (III) ions (Cr3+).
Add 1 mL of 0.05% potassium dichromate solution and one drop of concentrated sulfuric acid to a small test-tube.
Pour 10 mL pure ethanol (absolute alcohol) into a small distilling flask and heat the flask slowly.
Pass the ethanol vapour through the potassium dichromate solution.
The colour of the solution changes from orange to green.
Cr2O72- (aq) + 8H+ (aq) + 3C2H5OH (l) --> 2Cr3+ (aq) + 3CH3CHO (l) + H2O (l)
K2Cr2O7 + 4H2SO4 + 3C2H5OH --> K2SO4 + Cr2(SO4)3 + 3CH3CHO + H2O
K2Cr2O7 + 4H2SO4 + 7CH3CHO --> K2SO4 + Cr2(SO4)3 + 7CH3COOH.
1. Breath test for alcohol using breath analyser ("breathalyser")
Test a breath analyser used by police or hospital staff.
In some countries, the breath testing apparatus used by police to detect motorists who have consumed too much alcohol is called a "breathalyser".
Borrow a breath testing apparatus from the police.
Ethanol vapour in the breath reduces orange potassium dichromate (K2Cr2O7) to green chromium ions (Cr3+).
The legal limit in some countries is 80 mg of ethanol per 100 mL of blood.
15.3.1 Oxidants (oxidizing agents)
Oxidation is loss of electrons.
An oxidant causes the oxidation of another substance as it is reduced.
A "good" oxidant is easily reduced.
An oxidizing agent helps the oxidation of another chemical.
An oxidizing agent is a substance that causes oxidation.
An oxidizing agent is easily reduced, i.e. it gains electrons easily.
The oxidizing agent gains the electrons and the substance being oxidized loses electrons.
During oxidation, the oxidizing agent is reduced.
When ferric chloride solution is added to stannous chloride solution, ferric chloride is reduced and stannous chloride is oxidized.
Sn2+ + 2Fe3+ --> Sn4+ + 2Fe2+
Sn2+ - 2e- --> Sn4+ (oxidation)
2Fe3+ + 2e- --> 2Fe3+ (reduction)
Oxidants may react violently, with reductants and with common organic materials, e.g. paper, cotton.
Mix oxidants with reductants under carefully controlled conditions.
Store oxidants separately from reductants.
Do not heat or grind in a mortar heat-sensitive and friction-sensitive oxidants, e.g. nitrates and peroxides.
Perchloric acid, potassium and potassium chlorate can explode when mixed with organic material.
Permanganates, nitrates, iodates, and periodates can form explosive mixtures with combustible materials, e.g. alcohols, aluminium dust, zinc dust and sulfur.
Nitric acid forms explosive mixtures with many organic compounds.
Do not mix nitric acid with ethanol, alcohols of low molecular weight, and oxygenated compounds, e.g. ketones and aldehydes.
Many oxidants are toxic heavy metal compounds, e.g. lead oxide and potassium dichromate.
15.3.2 Redox reactions, oxidation number
Oxidation and reduction reactions (redox reactions) must occur together.
In a redox reaction, the same number of electrons is gained in the reduction as is lost in the oxidation.
In the following reaction, O2 is an oxidizing agent and the H2 is a reducing agent:
2H2 (g) + O2 (g) --> 2H2O (l)
Oxidation number, oxidation state, is the "apparent charge" on an atom, molecule or ion.
The oxidation number of an element is zero, of hydrogen is +1, of H2 is zero, of hydride (e.g. NaH) is -1, of oxygen is -2, of O2 is zero, and in H2O2 is -1.
Oxidation number increases when something is oxidized and decreases when something is reduced.
The sum of the oxidation numbers of atoms in a molecule or ion is equal to the total charge on the molecule or ion.
In the following equation:
2NH3 + 3CuO --> 3Cu + 3H2O + N2
2NH3 --> N2, so oxidation number increases by 6.
3CuO --> 3Cu, so oxidation number decreases by 6.
15.3.3 Reductants (reducing agents)
Reduction is gain of electrons.
A reductant causes the reduction of another substance as it is oxidized.
A "good" reductant is easily oxidized.
A reducing agent helps the reduction of another chemical.
A reducing agent is easily oxidized, i.e. it loses electrons easily.
Examples of reducing agents include the following: Zn metal that is easily oxidized to zinc ion, Zn2+.
Hydrogen sulfide that reacts with chlorine to form sulfur.
Carbon reduces lead (II) oxide to lead.
Carbon monoxide reduces Fe (III) oxide to iron in a blast furnace.
Reductants react, often violently, with oxidants and should only be used under carefully controlled conditions.
They should always be- stored dry and separated from oxidants.
Alkali metals (sodium and potassium) and their amalgams react violently with water, often with the evolution and ignition of hydrogen.
Metal hydrides are also very reactive with moisture and can ignite or explode on contact with water.
Contact with skin or eyes is therefore very dangerous and must be avoided.
Metal powders are extremely reactive with oxidizing agents, are toxic if inhaled and may cause skin sensitization.
15.3.4 When oxidation occurs:
1.1 The substance combines with oxygen.
Oxygen is added to an element or compound, e.g. burning a substance in air.
C (s) + O2 (g) --> CO2 (g), carbon is oxidized,
2Mg + O2 --> 2MgO, magnesium oxidized,
2CO + O2 --> 2CO2, carbon monoxide is oxidized.
1.2 The substance loses hydrogen.
The removal of hydrogen from a compound.
4HCl (aq) + MnO2 (s) --> MnCl2 (aq) + 2H2O (l) + Cl2 (g), hydrochloric acid is oxidized,
2H2S + O2 --> 2S + 2H2O, hydrogen sulfide is oxidized,
H2S + Cl2 --> S + 2HCl, hydrogen sulfide is oxidized.
1.3 Oxidation is an increase of valence.
In the following equation divalent iron is oxidized to trivalent iron.
2FeCl2 + Cl2 --> 2FeCl3
2Fe2+ + Cl2 --> 2Fe3+ + 2Cl-.
1.4 Oxidation is the loss of electron (s), e.g. when a ferrous ion changes to a ferric ion.
e2+ - e- --> Fe3+.
15.3.5 When reduction occurs:
2.1 A substance loses oxygen, e.g. In the following reaction Copper (II) oxide loses oxygen and changes to copper.
CuO (s) + H2 (g) --> Cu (s) + H2O (g).
2.2 A substance gains hydrogen, e.g. In this reaction, nitrogen gains hydrogen to become ammonia.
N2 (g) + 3H2 (g) --> NH3 (g)
nitrogen + hydrogen --> ammonia.
15.4.1 Manganese oxidising agents
The standard electrode potential of MnO4- has E value of 1.51, so a strong oxidising agent, then Mn with E value of 1.18 (reversing the Mn2+ equation), then Mn2+ with E value of -1.18 V.
So for weakest to strongest oxidising agents, the order should be Mn2+ , Mn, MnO4-.
15.4.2 Oxygen as an oxidizing agent
Oxygen molecules (O2) gain electrons to form oxide ions (O2-).
15.4.3 Bromine as an oxidizing agent
Add drops of bromine water to 2 cm of ferrous sulfate in a test-tube.
The green ferrous salt turns yellow, forming a ferric salt.
2Fe2+ + Br2 --> 2Fe3+ + 2Br-
Fe2+ - e- --> Fe3+,(ferrous ion oxidized).
Br2 + 2e- + 2Br-,(bromine reduced).
To show the presence of a ferric salt, add sodium hydroxide solution to form a brown precipitate of ferric hydroxide.
15.4.4 Chlorine as an oxidizing agent
Chlorine molecules (Cl2) gain electrons to form chloride ions (Cl-).
2FeCl2 + Cl2 --> 2FeCl3
2Fe2+ (aq) + Cl2 --> 2Fe3+ (aq) + 2Cl- (aq).
15.4.5 Concentrated nitric acid as an oxidizing agent
Pass hydrogen sulfide through concentrated nitric acid.
Observe the precipitation of sulfur as a yellow suspension as the concentrated nitric acid acts as an oxidizing agent.
H2S <--> 2H+ + S2-
2H+ + S2- + 2H+ + NO3- --> S + 2H2O + 2NO2
S2- - 2e- --> S (sulfide ion oxidized)
4H+ + 2NO3- + 2e- --> 2H2O + 2NO2 (nitric acid reduced).
15.4.6 Hydrogen peroxide as an oxidizing agent
Hydrogen peroxide turns an iodide solution brown, forming iodine and perhaps precipitating black crystals of iodine.
1. Add drops of hydrogen peroxide solution to 2 cm of potassium iodide solution in a test-tube.
2H+ + 2I- + H2O --> 2H2O + I2
2I- - 2e- --> I2 (iodide ion oxidized)
2H+ + H2O2 + 2e- --> 2H2O (H2O2 is reduced).
2. Add drops of potassium iodide solution to 20 vols (6%) hydrogen peroxide solution.
Then add the same number of drops of dilute sulfuric acid.
Heat gently and note any colour change.
Add drops of starch solution.
A blue black colour suggests oxidation of 2I- to I2.
H2O2 (aq) + 2H+ (aq) + 2e-- --> 2H2O (l)
2I- (aq) --> I2 (s) + 2e-
H2O2 (aq) + 2H+ (aq) + 2I- (aq) --> I2 (s) + 2H2O (l)
Or
I2 (s) + I- (aq) --> I3- (aq)
H2O2 (aq) + 2H+ (aq) + 3I- (aq) --> I3- (aq) + 2H2O (l).
15.4.7 Hydrogen sulfide as a reducing agent
The use of Kipp's apparatus as a source of hydrogen sulfide is NOT recommended in this document.
1. In a fume cupboard, pass hydrogen sulfide gas into a dilute acidified potassium permanganate solution.
The colour of the potassium permanganate disappears, but a milky precipitate of sulfur remains.
2MnO4- + 6H+ + 5H2S --> 2Mn2+ + 8H2O + 5S (s).
2. Pass hydrogen sulfide for ten minutes through a dilute solution of ferric chloride acidified with a few drops of hydrochloric acid.
The colour will change from yellow to green.
Boil the solution in a dish for two minutes to expel hydrogen sulfide, filter through a double filter paper to remove sulfur, and add sodium hydroxide soda solution in excess to the filtrate.
A dirty green precipitate of ferrous hydroxide will be obtained showing that the ferric ion has been reduced to ferrous ions.
2Fe3+ + H2S --> 2Fe2+ + 2H+ + S (s).
15.4.8 Nitrous acid as oxidizing agent or reducing agent
1. Nitrous acid can act as an oxidizing agent.
Slowly add sodium nitrite solution to potassium iodide solution acidified with dilute sulfuric acid.
Iodine forms showing that the nitrous acid produced by the action of the dilute acid on the sodium nitrite has oxidized the potassium iodide.
The nitrous acid has itself been reduced to nitric oxide.
The nitric oxide forms brown fumes of nitrogen dioxide when it contacts the oxygen of the air.
2NO2- + 2I- + 4H+ --> I2 + 2NO + 2H2O
When acting as an oxidizing agent, nitrous acid gains electrons and is reduced to nitric oxide.
2NO2- + 4H+ + 2e- --> 2H2O + 2NO.
2. Nitrous acid can act as a reducing agent.
Acidify potassium permanganate solution with dilute sulfuric acid and add sodium nitrite solution until the colour of the potassium permanganate just disappears.
Note the absence of brown fumes of nitrogen dioxide.
The solution contains nitric acid and can be tested by the nitrate test.
The potassium permanganate has oxidized the nitrous acid to nitric acid.
The potassium permanganate is reduced to manganous salts.
2MnO4- + 6H+ + 5NO2- --> 2Mn2+ + 3H2O + 5NO3-
Nitrous acid here acts as a reducing agent; it loses electrons and is oxidized to nitric acid.
NO2- + H2O - 2e- --> NO3- + 2H+.
15.4.10 Sulfurous acid as a reducing agent
Ionization reaction:
H2SO3 + H2O <--> H3O+ + HSO3-
HSO3-+ H2O <--> H3O+ + SO32-
1. In a fume cupboard, pass sulfur dioxide or sulfurous acid into a dilute acidified potassium permanganate solution.
The colour of the potassium permanganate disappears, but no precipitate of sulfur is formed.
The sulfurous acid has been oxidized to sulfuric acid.
2MnO4- + 6H+ + 5SO32- --> 2Mn2+ + 3H2O + 5SO42-.
2. Pass sulfur dioxide continuously through a dilute solution of ferric chloride.
The liquid turns red, because of the formation of a complex sulfite.
Transfer the solution to a dish and boil for a few minutes on a tripod and gauze.
The resulting solution will be pale green or colourless.
Add sodium hydroxide solution in excess to a sample where a dirty green precipitate of ferrous hydroxide shows that reduction is complete.
2Fe3+ + SO32+ + H2O --> 2Fe2+ + SO42- + 2H+.
3. Dissolve potassium iodate in water in a boiling tube and pass sulfur dioxide through it.
The iodate is reduced to iodine that is deposited as black crystals.
IO3- + 3SO32- --> I- + 3SO43-
5I- + IO3- + 6H+ -->3I2 (s) + 3H2O
If the stream of sulfur dioxide continues for a few minutes, the solution goes clear, because of the formation of hydrogen iodide.
I2 + SO32- + H2O --> 2I- + SO42- + 2H+.
15.4.12 Gummi bears, (gummi fruit), with potassium chlorate
This experiment is not allowed in the school laboratory, but it is done in science fairs by experienced staff.
Put 5 g potassium chlorate in large test-tube clamped to a ring stand.
Use long handled tongs to drop a gummi bear into the test-tube.
Sugar in the gummi bear reacts violently with the potassium chlorate.
The gummi bear jumps around in bright purple flames.
Gummi bears may contain high fructose corn or wheat glucose syrup, sugar, gelatine, reconstituted apple juice, food acid 330, flavours, glazing agent 903, colours 102, 110, 122, 133, and wheat starch.
They are designed to give an intensive energy "lift" to people who have been exercising intensively.
15.4.13 Potassium chlorate and potassium persulfate as oxidizing agents
Arrange in test-tube pairs 2 cm of 1. acidified potassium iodide solution, 2. acidified ferrous sulfate solution, 3. hydrogen sulfide solution, 4. concentrated hydrochloric acid.
Add 0.55 cc of potassium chlorate to one set and add 0.55 cc potassium persulfate to the other set.
Note the reaction and warm to completion if necessary.
Note in which case the reaction occurs more readily.
Both potassium chlorate and potassium persulfate are powerful oxidizing agents.
The persulfate ion oxidizes by accepting electrons to become sulfate ions, e.g. using potassium iodide.
S2O82- + 2I- --> 2SO42- + I2 S2O82- + 2e- -->2SO42- (persulfate ion reduced).
15.4.14 Potassium chlorate in pyrotechnic flash powders
The reactions below are not permitted in schools.
potassium chlorate + aluminium powder + optional sulfur powder, burns quickly with loud noise, used in stun grenades.
KClO3 + 2Al --> Al2O3 + KCl
Potassium perchlorate in fireworks.
3KClO4 + 8Al --> 4Al2O3 + 3KCl
Magnesium photographic flash powder.
2KNO3 + 5 Mg --> K2O + N2 + 5MgO
Antimony trisulfide and chlorate in cheap small firecrackers.
3KClO3 + Sb2S3 --> Sb2O3 + 3SO2 + 3KCl.
15.4.15 Sparklers
Sparklers may be illegal in some places and are not allowed in some school systems.
Mix the following dry ingredients in a plastic coffee can with an old spoon: 300 parts potassium chlorate, 60 parts aluminium fine granules, 2 parts charcoal.
If red colour is needed, add 500 parts strontium nitrate.
If green colour is needed, add 60 parts barium nitrate.
Mix 10% powdered dextrin in water and add the dextrin solution to the solution a little at a time while stirring to form a moist slurry.
Cut 15 cm lengths of heavy iron wire, e.g. coat hanger wire.
Prepare sparklers:
Dip the lengths of iron wire into the slurry leaving 5 cm undipped to use as a handle, then leave the sparklers to dry.
Store the sparklers in a cool, dry place.
If the ingredients are not sticking to the iron wire, increase the concentration of dextrin.
Before using sparklers, instruct the users to keep a safe distance apart and to drop used sparklers in a metal container of sand.
After use, the iron wire of used sparklers remains hot for some time.
The unusual "sparkling" burning of a sparkler is caused when larger iron particles from iron wire are ejected explosively, while smaller particles burn completely.
15.4.16 Potassium dichromate as an oxidizing agent
1. Add potassium dichromate solution and drops of dilute sulfuric acid to iron (II) sulfate solution.
The dichromate ion (Cr2O72+) is reduced to Cr3+ and the solution changes from orange to green.
The iron (II) ions (Fe2+) are oxidized to iron (III) ions (Fe3+).
Cr2O72+ (aq) + 14H+ (aq) + 6e-- --> 2Cr3+ (aq) + 7H2O (l)
6Fe2+ (aq) --> 6Fe3+ (aq) + e-
Cr2O72+ (aq) + 14H+ (aq) + 6Fe2+ (aq) --> 2Cr3+ + 7H2O (l) + 6Fe3+(aq).
2. Potassium dichromate (VI) solution, acidified with dilute sulfuric acid, is used as an oxidizing agent to oxidize primary alcohols to aldehydes and carboxylic acids, and oxidize secondary alcohols to ketones.
This reaction was used in an early form of the breathalyzer to detect alcohol.
Cr2O72- + 14H+ + 6e- --> 2Cr3+ +7H2O (l)
orange solution reduced to green solution.
3. Ethanol oxidized to ethanal, an aldehyde
Ethanol in excess, so not enough oxidizing agent, so only partial oxidation to an aldehyde
Use potassium dichromate (VI) solution + sulfuric acid + ethanol
Cr2O72- (aq) + 8H+ + 3C2H5OH (aq) --> 2Cr3+ (aq) + 7H2O (l) + 3CH3CHO (aq) (ethanal, acetaldehyde)
(Distil off the ethanal aldehyde as soon as it forms.)
4. Ethanol oxidized to ethanoic acid, a carboxylic acid
Oxidizing agent Cr2O72- in excess, so full oxidation to carboxylic acid.
Use potassium dichromate (VI) solution + sulfuric acid + ethanol
2Cr2O72- (aq) + 16H+ + 3CH2H5OH(aq) --> 4Cr3+ + 11H2O (l) + 3CH3COOH(aq) (ethanoic acid, acetic acid)
(Use a reflux condenser to form ethanoic acid, a carboxylic acid).
5. Secondary alcohols
propan-2-ol + potassium dichromate (VI) solution + dilute sulfuric acid --> water + propanone ketone
CH3CHOHCH + O (from oxidizing agent) --> (CH3)2CO + H2O
propan-2-ol (isopropanol) --> propanone (acetone) + water.
15.4.17 Tests for oxidizing agents
1. Tests for oxidizing agents by change in colour of iron (II) to iron (III)
Prepare iron (II) sulfate solution by dissolving iron filings in dilute sulfuric acid.
When the reaction stops, filter the solution.
The filtrate is acidified iron (II) sulfate solution that is green.
Add the test solutions and gently heat.
If the solution turns brown, Fe2+has changed to Fe3+, because of the presence of an oxidizing agent.
2. Tests for oxidizing agents by change in colour of iron with copper (II) sulfate
Add iron to copper (II) sulfate solution.
Note the colour change.
The copper ion is an oxidizing agent.
The blue colour is removed as copper forms.
Cu2+ (aq) + Fe (s) --> Fe2+ (aq) + Cu (s).
3. Tests for oxidizing agents by change in colour of zinc with copper (II) sulfate
In this reaction, the copper ion Cu2+ attracts electrons better than the zinc ion, Zn2+.
The Zn is oxidized to zinc ions and the copper is reduced to copper metal.
Red copper precipitates and the solution lose its blue colour.
Add pieces of zinc to copper (II) sulfate solution.
The zinc corrodes and goes into solution.
Red copper precipitates and the solution lose its blue colour.
Add excess zinc so that all the copper precipitates.
Decant the solution and evaporate to leave zinc sulfate crystals.
Add excess zinc so that all the copper precipitates.
Decant the solution and evaporate to leave zinc sulfate crystals.
Zn (s) + CuSO4 (aq) --> Cu (s) + ZnSO4 (aq)
Zn (s) + Cu2+ (aq) --> Cu (s) + Zn2+.
15.4.18 Solids that conduct electricity
1. Use a 6 V dry cell or lead cell accumulator and a 1.5 V light bulb.
Fix electrodes from old 6 V dry cells in a cork to keep them at a constant distance apart.
Test the conductivity of solids by making a good contact between the surfaces of the test solid and the two electrodes.
Test metals and non-metals, e.g. scissors, nails, plastic, paper, naphthalene, wax, sugar, sodium chloride, and water.
Record which substances are conductors and non-conductors, insulators.
2.Test conductivity of glass.
Test the conductivity of a glass rod at room temperature.
Heat the glass rod until it becomes very hot and begins to soften.
Test the hot soft part with the conductivity apparatus.
Molten glass can be a good conductor of electricity.
15.4.19 Liquids that conduct electricity
1. Test melted substances.
If you heat the following substances, heat very gently and cautiously, because they may ignite and burn: sulfur, wax, naphthalene, polyethylene, tin, lead, a low melting point salt, e.g. lead bromide, m.p. 488oC, or potassium iodide, m.p. 682oC.
To test the conductivity of the melt, dip the electrodes in the melt and wait for the electrodes to reach the same temperature as the melt.
Make sure that the electrodes are in contact with the liquid melt and not the solidified melt.
Scrape and clean the electrodes between each test.
2. Test methylated spirit, acetone, vinegar, sugar solution, copper (II) sulfate solution, sodium chloride solution, and other substances dissolved in water.
Clean and dry the electrodes between each test.
3. Test demineralized water.
Put the electrodes into a container of deionized water.
The light bulb does not light.
Slowly add small crystals of sodium chloride to the demineralized water.
Observe the light bulb as the salt dissolves.
4. Test tap water.
Note whether you get the same result as for deionized water.
15.5.1 Rusting
Rusting is an electrochemical process that needs water and oxygen.
At the anode:
Fe (s) --> Fe2+ (aq) + 2e-
At the cathode:
O2 (aq) + 2H2O (l) + 4e --> 4OH- (aq)
OR
1/2O2 + H2O + 2e- --> 2OH-
The Fe(OH)2 solution oxidizes to rust (Fe2O3.xH2O, hydrated iron oxide).
Both air and water are necessary for rusting of iron.
When in moist air, iron is very liable to form rust, most of which is Fe2O3.xH2O or Fe(OH)3.xH2O.
Rust forms on the surface, because of the action of water and oxygen on it.
Show that oxygen occupies about one fifth of the atmosphere by volume based on the decrease in the air volume during rusting.
15.5.2 Conditions necessary for rusting
See diagram 3.3.6: Conditions necessary for rusting
Prepare clean nails and put them in test-tubes as follows:
Test-tube 1: Put the nail in the test-tube and half cover with demineralized water.
The nail is in contact with air and water.
Test-tube 2: Put anhydrous calcium chloride or silica gel in the bottom of the test-tube.
Put a nail in the test-tube and fix put a plug of cotton wool at the mouth.
The nail is in contact with air, but not with moisture.
Test-tube 3: Boil water for several minutes to expel dissolved air then pour the water into the test-tube while still hot.
Put the nail in the test-tube and half cover with hot water.
Put petroleum jelly or olive oil on the surface of the hot water.
The petroleum jelly melts and forms an airtight layer then solidifies as the water cools.
Half the nail is in contact with water, but not with air.
Test-tube 4: Dissolve sodium chloride crystals in water in the test-tube.
Put the nail in the test-tube and half cover with salt solution.
The nail is in contact with air, water and sodium chloride.
Leave the test-tubes in a rack for several days and note the conditions for rusting.
Most rust forms in test-tube 4, then test-tube 1, then test-tube 2 or test-tube 3.
Conclusions:
Air and water are necessary for rusting.
Sodium chloride solution increases rusting.
2. Use four test-tubes fitted with corks each containing two identical clean nails.
Use rainwater.
Half the nail is in contact with water and half the nail is in contact with air.
This the control test-tube.
Put anhydrous calcium chloride or silica gel in the test-tube.
Plug the test-tube with cotton wool.
The nail is in contact with air, but is not in contact with moisture.
Pour water into the test-tube and boil for some minutes to expel all the dissolved air.
Pour oil on the surface of the water to form an airtight layer.
The nail is contact with water, but is not in contact with air.
Use salt water.
Half the nail is in contact with the salt water and half the nail is in contact with air.
The nail is in contact with air and salt water.
Observe more rusting in test-tube 1.4 than test-tube 1.1.
Observe no rusting in test-tubes 1.2. and 1.3.
3. Label three test-tubes 1, 2 and 3.
Roll up some iron wool, steel wool, to make three little balls that will fit inside the test-tubes.
Put one ball in each test-tube.
Leave test-tube 1 as a control.
Add water to test-tube 2 until the ball is half-covered.
Shake the test-tube to wet all the steel wool.
Add salt water to test-tube 3 until the steel wool is half-covered.
Shake the test-tube to wet all the steel wool.
Put the three test-tubes where they won't be disturbed.
Observe them after a week.
15.5.3 Electrode potential order of metals
See diagram 3.86: Electrode potential apparatus
A metal tested, B copper sulfate in filter paper, C copper foil.
1. Electrode potentials of metals are calculated from comparisons with the hydrogen cell under standardized conditions.
However, you can use a copper and copper (II) sulfate solution as a standard.
Lay filter paper soaked with copper (II) sulfate solution on clean copper foil.
Use a short length of wire and crocodile clips to connect the copper foil to the positive terminal of the 1 to 5 V voltmeter.
Similarly connect the specimen metal to the negative terminal of the voltmeter.
Clean the surface of the specimen metal and press it firmly on absorbent paper.
Record the voltage for this metal.
2. Test the following metals: magnesium, tin, lead, iron, zinc aluminium, and silver.
After testing a metal, clean the copper again with a fine emery cloth and replace the absorbent paper, then test another metal.
The metal surfaces must be clean and the absorbent paper must contain enough copper (II) sulfate solution for a steady reading on the voltmeter.
If the voltage starts at a high value and then falls as a deposit forms on the metal, record the highest value.
3. Test aluminium after dipping it in concentrated hydrochloric acid then press it on absorbent paper to remove the layer of aluminium oxide.
The voltage reading will start at a low value then increase as remaining aluminium oxide dissolves.
Record the maximum value.
15.5.4 Metals can prevent rusting
Test-tube 1: Put a nail in the test-tube and half cover with demineralized water.
Test-tube 2: Put a nail in the test-tube and half cover with tap water.
Test-tube 3: Wrap a piece of zinc foil around one end of a nail.
Put the nail in the test-tube and half cover with tap water.
Test-tube 4: Wrap a piece of tin foil around one end of a nail.
Put the nail in the test-tube and half cover with tap water.
Test-tube 5: Wrap a piece of copper wire around one end of a nail.
Put the nail in the test-tube and half cover it with tap water.
Note whether zinc, copper or tin best prevents rusting.
15.5.5 Need for oxygen for rusting
1. Compare the heights of water in the two measuring cylinders in the previous experiment.
The water level is higher in the cylinder containing the rusted steel wool.
The height of water rises until the original volume of air in the cylinder decrease by one fifth.
This proportion represents how much oxygen is in the air.
The lost oxygen is combined with the iron of the steel wool to form rust.
2. Moisten inside a test-tube with water.
Put iron filings in the bottom of the test-tube and insert a piece of cotton wool to keep them in place.
Invert the test-tube in a beaker one third full of water.
The water levels inside and outside the test-tube should be the same.
Mark the original water level on the outside surface of the test-tube.
After two days, the iron rusts and the water level rise inside the tube until it is steady.
About one fifth of the original air in the test-tube is used up.
This suggests that when iron filings rust, oxygen is used.
3. Repeat the experiment with magnesium ribbon replacing iron wire.
The water height inside the graduated cylinder will go down to give an increase in the air volume.
This result comes from the hydrogen gas formed in the reaction of magnesium with water.
15.5.6 Oxygen gas combines with iron during rusting
See diagram: 3.2.42: Steel wool rusting
Moisten inside a test-tube with water.
Sprinkle iron filings into the test-tube then rotate it horizontally to make the iron filings spread and stick to the walls.
Invert the test-tube in a container one third full of water.
Support the test-tube so that the water levels inside and outside the test-tube are the same.
Mark the levels on the test-tube with a grease pencil and leave for a few days.
The iron will rust and the water level will rise inside the test-tube, finally becoming steady.
Add water to the container until the levels inside and outside the test-tube are the same and mark the new level.
About one fifth of the air volume has been used up, suggesting that oxygen has been used up in the rusting of iron.
Test the remaining gases with a lighted splint.
The lighted splint is extinguished.
15.5.7 Rusting of iron wire
See diagram 3.3.1: Rusting iron wire
Rusting is the corrosion of iron and iron-based alloys, eg. steel.
The corrosion product of iron is termed rust, or iron oxides to give rust its chemical name.
Rusting of iron is a multi-step process where metallic iron changes to Fe(OH)3.xH2O.
Rusting needs air and water and increases if the water contains salts.
Prevent rusting by painting outside surfaces or by oiling machinery surfaces or by absorbing moisture with silica gel to protect delicate machinery, e.g. cameras.
Experiments
1. When iron rusts, it changes from Fe to Fe2O3.xH2O.
Weigh some dry iron filings.
Leave in moist air for two days.
Note any increase in weight as rust forms.
2. Fill a 30 mL wide necked bottle with a big ball of polished thin iron wire (about 0.6 g).
Add water to soak the iron wire and then pour the water out.
Close the mouth of the bottle with a rubber stopper fitted with a 40 cm straight glass tube.
Invert the bottle and clamp it on an iron stand with the end of the glass tube under the water in a beaker.
Mark the original water level on the outside of the glass tube.
Note the water height every hour.
The water level rises slowly in the first five hours and then goes up at a faster speed of about 0.5-0.6 cm an hour.
After one day, rising of the water level slows again.
3. Polish 0.4 g (about 130 cm long) of thin iron wire (or thin wire gauze) and curl it into a small ball.
Push the ball into the bottom of a 10 mL graduated cylinder.
Add water to immerse the iron wire and cover the mouth of the cylinder with a slice of glass.
Holding the glass slice, invert the cylinder and adjust the water height to a certain mark (say, "9.0 mL") by carefully moving the glass slice.
Stand the inverted graduated cylinder over a dish containing water and remove the glass slice.
After two days, much reddish brown rust forms on the surface of the iron wire and the water level rises to show a one fifth decrease, (about 1.8 mL if the original water level is adjusted to "9.0") in the air volume inside the cylinder.
4. Mass of iron and its temperature increases during rusting
See diagram 3.2.41: Counterbalanced iron
Balance a piece of iron or steel wool on a knife edge with a brass weight.
Leave in moist air for a few days and note the effect of rusting on the weight of the iron.
During rusting, iron changes to iron (III) oxide hydrate, hydrated ferric oxide.
4Fe + xH2O + 3O2 --> 2Fe2O3.xH2O.
5. Weigh a piece of steel wool, add water and allow it to rust.
Dry it carefully and weigh the steel wool again.
The rusted steel wool is heavier than the unrusted steel wool.
6. Invert a measuring cylinder and place it below the surface of the water.
Use a bent rubber tube to withdraw air from the cylinder until the volume of air is exactly 100 mL when the water level inside and outside the cylinder are the same.
Push some steel wool up to the closed end of the cylinder.
Leave the apparatus for a few days and note how much gas is left in the cylinder.
Push a lighted taper into the gas and note what happens.
Remove the steel wool and examine it.
When iron rusts, an increase in weight occurs.
The iron combines with oxygen from the air.
The remaining gas extinguishes a lighted taper placed in it.
7. Almost fill a thermos flask with loosely packed, damp, steel wool.
Place a one-hole stopper fitted with a thermometer into the mouth of the flask.
The bulb of the thermometer bulb must be touching the steel wool.
Record the initial temperature of the apparatus and note the daily temperatures for five days.
The rusting of steel wool is an exothermic reaction.
15.5.8 Rusting of steel wool
See diagram 3.3.3: Rusting steel wool
1. Rusting forms red hydrated, Fe2O3.
Wrap a thermometer bulb in wet steel wool results and note the temperature rise.
Fe2O3 is a red pigment.
4Fe (s) + 3 O2 (g) --> 2 Fe2O3 (s).
2. Use two measuring cylinders.
Push steel wool into the bottom of one measuring cylinder and leave the other as a control.
Pour 50 mL water into each measuring cylinder.
Hold a piece of cardboard over the mouth of each measuring cylinder and invert it over a shallow dish containing water.
Remove the cardboard.
Adjust the height of the water in each inverted measuring cylinder by blowing in air with a bent pipette so the height of water in two measuring cylinders is the same.
Leave the experiment for several days.
3. Repeat the experiment with salty water.
The rusting occurs more quickly, not because the sodium chloride takes part in the reaction, but because it gives the water more conduction.
Similarly, the presence of sulfur dioxide in the air in cities and industrial sites increases the rate of rusting.
Fe + 1/2O2 + H2 (from water) -->Fe(OH)2 [iron (II) hydroxide]
4Fe(OH)2 + O2 --> 2Fe2O3.3H2O + H2O [iron (III) oxide]
The Fe(OH)2 in solution is oxidized to Fe2O3.
4. Fit a small wide mouth bottle with a rubber stopper and a glass tube about 3 m long.
Fit the bottle with a rubber stopper and a glass tube about 3 m long.
Use a bundle of steel wool that is big enough to fill the bottle.
Remove any oil from the steel wool by washing it in petrol then leaving it to dry.
Put the steel wool in the bottle and insert the stopper fitted with a glass tube.
Invert the bottle and support it with the end of the tube under water.
Record the water level in the tube each hour.
5. Moisten some steel wool with iron chloride solution to accelerate rusting.
Wrap the bulb of a thermometer in the steel wool.
Hang in a draught free place.
Note the temperature changes as rust forms.
6. Roll some steel wool into a ball and weigh it.
Use tongs to hold the ball of steel wool over a sheet of paper.
Heat the steel wool over a burner until red-hot.
Remove the burner and blow gently on the red hot steel wool until it stops burning.
Weigh the burned steel wool and any fragments that have fallen on to the sheet of paper.
The weight is greater, because the iron oxide that forms is heavier than the steel wool.
15.5.9 Rusting paper clips
Observe a paper clip used to clip together an old pile of paper.
Note any rust marks.
Rusting starts where the paper clip is closest to the paper.
This occurs because there is the least exposure to oxygen gas to allow the chromium layer to produce protective oxides.
15.5.10 Stainless steel
Stainless steel contains 12-5% Cr to produce the stainless chromium oxide film on the surface to prevent corrosion.
So stainless steel must be kept clean to maximum availability of oxygen to the chromium atoms.
15.6.0 Cathodic protection
Electrochemical prevention of rusting
Cathodic protection (CP) is a method of corrosion control that can be applied to buried and submerged metallic structures.
It is normally used in conjunction with coatings and can be considered as a secondary corrosion control technique.
The primary corrosion control method on any given structure is a coating, which can be between 50 and 99 % efficient, depending upon age and installation.
A properly designed and maintained cathodic protection system will take up the remainder resulting in a 100 % efficient corrosion protection system.
Cathodic protection reduces the rate of corrosion of a metallic surface by making the metal the cathode.
This can be achieved by attaching galvanic (sacrificial) anodes or impressed current anodes.
A more electronegative metal, e.g. zinc, is attached as a "sacrificial anode" that goes into solution instead of the iron.
Also, you can apply direct current to make the iron into a cathode.
Cathodic protection can protect iron ships and bridges from corrosion.
Other metals, including aluminium alloys, stainless steels, brasses and bronzes all require some level of protection against corrosion.
Wooden sailing ships were protected from fouling organisms by the release of copper ions from copper sheathing of the ship's bottom.
However, copper sheathing on an iron bottom ship produced an electrochemical cell in the sea water that corroded the iron.
This could be prevented by attaching blocks of zinc the bottom to give cathodic protection to the copper.
Experiments
1. A "tin can" is made by covering sheets of iron with tin plate to exclude oxygen.
If the "tin can" is scratched and it is wet, the iron corrodes very rapidly, because an electrochemical cell is set up.
Wrap a piece of aluminium foil around the lower part of a nail.
Put the nail and metal in a test-tube.
Add tap water to cover the lower part of the nail.
Use these metals: control (no metal) magnesium ribbon, zinc foil, copper wire, tin foil.
Put the test-tubes in a test-tube rack put stoppers on the test-tubes and leave them undisturbed for several days.
If a very small amount of sodium chloride is added to each test-tube, rusting can occur within an hour.
Rusting first starts in the test-tubes containing copper or tin, then it starts in the control.
Iron is more active than copper or tin, so forms the positive ion Fe2+ to react with negative ions in solution to form precipitates of rust on the nail.
No rusting occurs in the test-tubes containing magnesium ribbon or zinc, but they form ions that react with negative ions to form white precipitates.
2. Use a magnesium pencil sharpener with the iron blade still attached by a screw.
Add two drops of universal indicator to a dilute solution of sodium chloride in a beaker.
The neutral solution should be green.
Add the magnesium pencil sharpener.
Hydrogen gas forms at the steel blade.
The solution becomes a basic purple.
A white precipitate of magnesium hydroxide forms.
During the next few days, the magnesium becomes corroded (sacrificed), leaving the steel blade protected from rusting.
At the magnesium anode oxidation occurs.
Mg --> Mg2+ + 2e-
At the steel blade cathode reduction occurs.
2H2O + 2e- --> H2 + 2OH-.
15.7.1 Potential difference from combining half-cells, zinc and iron
Measure the potential difference of a zinc half-cell connected to an iron half-cell.
Use a strip of zinc metal in a zinc chloride solution and an iron nail in iron (II) sulfate solution.
Connect the two half-cells with a strip of filter paper soaked in potassium chloride solution to act as a salt bridge.
Complete the circuit by connecting leads from each metal to a voltmeter and read the voltmeter.
Electrons flow with potential difference of 0.32 V.
Zn (s) --> Zn2+ (aq) + 2e-(E0= +0.76 V)
Fe2+ + 2e- --> Fe (aq) (E0 = -0.44 V)
Zn (s) + Fe2+ --> Zn2+ + Fe (s) (E0= + 0.32 V).
15.7.2 Potential difference from combining half-cells, Zn and Cu, Zn and Pb
If Zn E0 = -0.76 V set up cells to measure the E0 values of copper [copper in copper (II) sulfate solution] and lead [lead in lead (II) nitrate solution].
15.7.3 Differences in potential on an iron nail
Soak 1 gm agar in 100 mL water for two hours then boil until dissolved.
Add phenolphthalein indicator and add acid or alkali until pH = 8.\
Add drops of freshly prepared potassium ferricyanide solution and pour into a Petri dish.
Add a very clean nail and place the petri dish on an overhead projector for some hours.
A pink colour forms around the shaft of the nail, because of hydroxide ions and blue-green colour forms at the head of the nail, because of Fe2+ ions.
The stressed head shows positive potential and the unstressed shaft shows negative potential.
At the anode: Fe (s) --> 6 Fe2+ (aq) + 2e-
At the cathode: O2 (aq) + 2H2O (l) + 4e-- --> 4OH- (aq).