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1. Corrosion

By the end of today’s lesson you will be able to explain what is meant by the term

corrosion and describe corrosion of iron.

• Corrosion is a chemical reaction where the surface of a pure metal changes from a pure element into a compound.

• This will cause the surface of the metal to be eaten away.

• Not all metals corrode with equal ease – some metals do not corrode at all.

• The more reactive a metal is the faster it will corrode.

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• Rusting is the term used to describe the corrosion of iron.

• No other metal apart from iron rusts. All reactive metals do corrode.

• Rust is the common name for the compound iron oxide.

Iron (III) Oxide

Property Test Result

Appearance (colour and state)

Effect on moist pH paper

Solubility in water

Flammability

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1. Collect a spatula and wooden splint, test tube, stopper, glass rod and test tube rack.

2. Rinse your test tube thoroughly, then place 2 cm depth of cold water in your test tube.

3. Add 1/4 spatula measure of iron (III) oxide powder into your test tube. Stopper and shake. – Does the powder dissolve?

4. Using a glass rod remove a drop of the liquid from your test tube and touch it against a piece of pH paper. – What affect has the iron (III) oxide had

on the pH of water?

Conditions for rusting

• Examine the tubes set up – look for signs of rusting. In which tube(s) did rusting occur?

• Which conditions are required for rusting to occur?

A B C

dryingchemical water

boiled water

oil(no water) (no oxygen)

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Detecting Rust• As iron metal rusts iron atoms are converted into

iron (II) ions.

• Fe(s) → Fe2+(aq) + 2e-

• We can detect the presence of iron (II) ions using ferroxyl indicator.

• Ferroxyl indicator changes from yellow to blue on contact with Fe2+

(aq) ions.

• The darker the shade of blue the more Fe2+(aq) ions

are present.

• Ferroxyl indicator also contains phenolphthalein.

• Phenolphthalein changes colour from colourless to pink on contact with hydroxide ions (OH-

(aq) ions).

• The darker the shade of pink achieved the more OH-

(aq) ions are present.

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1. Collect 5 test tubes and a test tube rack.

2. Place 2 cm depth of cold tap water into each test tube.

3. Add a piece of metal to each tube –iron, tin, copper, zinc and steel.

4. Add 3 drops of ferroxyl indicator to each test tube.

5. Leave for 5 minutes and examine the colour developed.

Metal Colour of Ferroxylindicator

iron

tin

copper

zinc

steel

Steel is an alloy of iron and carbon.

Steel rusts in the presence of oxygen and water

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Summary• During corrosion metal atoms are converted

into metal ions.• Rusting specifically refers to the corrosion of

iron.• Rust is the common name for iron (III) oxide.• Oxygen and water are both required for

rusting to occur.• Initial stages of rusting can be identified

using ferroxyl indicator. Ferroxyl indicator changes colour from yellow to blue if rusting is occurring.

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2. Changing the speed of rusting

By the end of today’s lesson you will be able to:• describe methods we can use alter the rate of iron rusting• state the % of oxygen found in air, and describe a suitable

experiment to illustrate this.

Speeding up RustingA B C D

sodiumchloridesolution

sulphurousacid

sugarsolution

water

Left for 10 minutes then 4 drops ferroxyl indicator added to each.

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• Which 2 nails rusted fastest?

• Which 2 nails rusted slowest?

• Which of the solutions are ionic, and which are covalent?

• Which solutions functioned as electrolytes?

• Which types of solutions speed up rusting?

Corrosive Chemicals

• What is the effect of using salt on icy roads has the rate cars rust at?

• Burning fossil fuels releases huge volumes of carbon dioxide and sulphur dioxide gases into the atmosphere.

• These gases combine with water to produce acidic solutions. The acidic solution falls to the earth as acid rain.

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• Write a balanced chemical equation for the formation of sulphurous acid (H2SO3) from sulphur dioxide and water.

• Write a balanced chemical equation for the formation of carbonic acid (H2CO3) from carbon dioxide and water.

• Acid rain is an ionic solution, it speeds up the rate of rusting.

• Acid causes the iron atoms to react to form iron (II) ions.

• write a balanced equations for the reaction of iron with sulphurous acid to form iron (II) sulphite.

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Chemistry of Rusting• Iron rusts to form iron (III) oxide in a series of chemical

reactions:

• Stage 1: OXIDATION – requires water• Fe (s) → Fe2+ (aq) + 2e- (loss of electrons)

• Stage 2: OXIDATION – requires oxygen• Fe2+ (aq) → Fe3+ (aq) + e- (loss of electrons)

• As these reactions occur, water and oxygen combine to form hydroxide ions – in a REDUCTION reaction.

• 2H2O (l) + O2 (g) + 4e- → 4OH- (aq) (gain of electrons)

• The rusting of iron is a REDOX reaction.

• The overall rusting reaction is:

• iron + oxygen → iron (III) oxide

• Write a balanced chemical equation for the overall rusting reaction.

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Summary

• Solutions containing ions, such as sodium chloride solution and sulphurous acid increase the rate of rusting.

• Acid rain increase the rate of rusting.• Rusting occurs in two stages:

– Fe (s) → Fe2+ (aq) + 2e-

– Fe2+ (aq) → Fe3+ (aq) + e-

• Both stages are examples of oxidation reactions.• As oxidation cannot occur in isolation, reduction must

also occur to create an overall REDOX reaction.

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3. Coatings

By the end of today’s lesson you will be able to describe the different methods of

coating iron and steel objects to prevent or slow down rusting.

• For an object to rust, oxygen and water must be able to access the surface of the iron object.

• If we prevent oxygen and/or water from reaching the surface of an iron object we will prevent rusting from occuring.

• Methods of protecting the surface of iron objects include greasing and oiling, painting, plastic coatings and metal coatings.

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Grease and Oil1. Quarter fill 2 test tubes with salty water.2. Add 4 drops of ferroxyl indicator to each.3. Dip one nail in a jar of grease.4. Put this nail in 1 test tube, and the

ungreased nail in the other test tube.5. Leave for 5 minutes and observe what

happens.

Draw a labelled diagram of your experiment shade the test tubes to reflect the colours at the end of the experiment.

• How does the grease prevent the nail from rusting?

• Why does this happen?

• Why do you think grease instead of paint is used to protect bicycle chains?

• What other iron objects are protected against rust in this way?

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Plastic Coatings

• How well would a plastic coating prevent iron from rusting?

• Why does this happen?

• Why do you think plastic instead of paint is used to protect wires?

Metal Coatings• Iron objects can be coated with a thin layer

of other metals to prevent the iron from rusting.

• Zinc, copper, silver, tin and gold can all be used to coat iron objects to protect them from rusting.

• Coating an iron object with zinc is known as galvanising.

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• How well would a zinc coating prevent iron from rusting?

• Why does this happen?

• Why do you think zinc instead of paint is used to protect dustbins that may contain hot ash?

Summary• Coating the surface of an iron object with a

physical barrier prevents rusting.• A physical barrier prevents oxygen and water

from reaching the surface of the iron object.• Grease, oil, paint, plastic and other metals

may all be used as a physical barrier.• If an iron object is coated with zinc, we say it

is galvanised.

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4. Electroplating

By the end of today’s lesson you will be able to • explain how to electroplate a metal object with a

thin layer of another metal• describe the chemical reactions that occur during

the electroplating process

• Electroplating is the use of electricity to coat an object with a thin layer of metal.

• Often iron objects are coated with a thin layer of chromium to make them shiny.

• Car parts are often coated with zinc by electroplating to prevent them from rusting.

• Why would a car not be made totally of zinc?

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Electroplating using Zinc• Scrub your iron nail with

detergent and iron wool.• Rinse with cold water.

• Soak nail in 3.0 mol l-1 sodium hydroxide solution for 3 minutes.

• Rinse with cold water.

• Soak nail in 3.0 mol l-1sulphuric acid for 3 minutes.

• Rinse with cold water.

• Assemble equipment shown. Leave running for 5 minutes.

Power Pack at 6 volts

(D.C.)

iron nail

50 cm3

plating solution

zinc strip

• What change did you see on the surface of the nail?

• What charge do zinc ions have?

• Why is the iron nail at the negative electrode?

• What will happen to the concentration of zinc ions as the experiments proceeds?

• Why is the positive electrode made of zinc metal?

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Positive zinc ions are attracted to the negative electrode (iron nail) and are converted into zinc atoms

To replace the zinc ions lost from solution, zinc atoms in the zinc electrodeare converted into zinc ions before being released into the solution.

• Reduction reactions refer to metals being extracted from their ores.

• Metal ions + electrons → metal atoms

• Oxidation reactions refer to metal atoms being converted into metal ions.

• Metal atoms → metal ions + electrons

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• Oxidation is Loss of electrons (OIL)

• e.g. Zn(s) → Zn2+(aq) + 2e-

• Reduction is Gain of electrons (RIG)

• e.g. Zn2+(aq) + 2e- → Zn(s)

• An oxidation reaction cannot occur without a corresponding reduction reaction – forming an overall REDOX reaction.

Summary• Electroplating means using electricity to plate

a metal from solution onto the surface of an object’s surface.

• During electroplating, a REDOX reaction occurs.

• Metal ions are converted into metal atoms (oxidation).

• Other metal atoms are converted into metal ions (reduction).

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5. Metals in Contact

By the end of today’s lesson you will be able to describe the affect different metals

have on the rate of rusting.

Nails in Contact

A B C

salty water+ 4 drops ferroxyl indicator

nail wrapped incopperwire

nail wrappedin zincstrips

• Scrub 3 nails using iron wool, then rinse in cold water. Set up as below.

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• Why did the nails have to have clean surfaces?

• Why was one nail left on its own?

• Which nail shows signs of rusting fastest?

• Which nail shows little sign of rusting at all?

• Shade your diagram to show the colours of the salty solution after 5 minutes.

Reactivity• Zinc is more reactive than iron.

• Copper is less reactive than iron.

• If zinc is in close contact with iron, it will prevent iron from rusting.

• If copper is in close contact with iron, it will speed up iron rusting.

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A B C

salty water+ ferroxyl indicator

nail wrapped inmagnesiumribbon

nail wrappedin tinfoil

• Scrub another 3 nails using iron wool, then rinse in cold water. Then set up as below.

• Is magnesium more or less reactive than iron? Why?

• Is tin more or less reactive than iron? Why?

• If you are given an unknown metal foil, and what to identify its place in the reactivity series, how could you go about this?

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Reactive Metals• Metals more reactive than iron can be used to

provide a physical and chemical barrier to prevent rusting.

• Galvanised buckets are made by coating iron buckets with a layer of zinc.

• When complete the zinc layer it provides a physical barrier to prevent oxygen and water from reaching the surface of the iron metal.

• If the layer of zinc is damaged, the remaining zinc provides chemical protection against rusting.

• Blocks of scrap magnesium are attached to steel pipes found underground.

• The magnesium blocks provide chemical protection against rusting.

• Similarly, scrap zinc blocks are attached to the hull of a steel ship.

• The zinc blocks provide chemical protection against rusting.

• The use of a more reactive metal to protect a ship is known as sacrificial protection.

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Summary• A more reactive metal will form ions more easily

than a less reactive metal.• More reactive metals corrode at a faster rate.• If two different metals are in physical contact

with each other, the more reactive metal will corrode and donate its electrons to the less reactive metal.

• The donation of electrons prevents the less reactive metal from corroding.

• This is sacrificial protection.

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6. Metals and Acids

By the end of today’s lesson you will be able to:•Describe the reaction between an acid and a metal

•State the affect physical factors have on the speed of a chemical reaction

Different Acids

Do all acids react with magnesium in the same way?

1. Place 1 cm depth of acid different acids into 4 test tubes.

2. Add a piece of magnesium to each separate test tube.

3. Leave for 30 seconds and observe any changes.

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• Which gas is produced when magnesium reacts with acids?

• What is the chemical test to identify this gas?

• What type of reaction have you observed in each case?

Different Metals

Do all metals react in the same way with hydrochloric acid?

1. Place 1 cm depth of hydrochloric acid in the bottom of 4 test tubes.

2. Add a piece of metal to each test tube.3. Observe the reaction in each tube.

Compare the reaction of different metals with hydrochloric acid.

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• Which metal reacted fastest with hydrochloric acid?

• Which metal reacted slowest with hydrochloric acid?

• Does the order of the speed of the reactions match the reactivity series shown on page 7 of your data booklet?

• What would you expect to happen if you dropped a silver coin into acid?

TemperatureDoes the temperature of the reaction mixture affect the speed of a chemical reaction?

1. Heat half a beaker of water until boiling.2. Put two pieces of zinc in separate test

tubes, together with 1 cm depth of acid in each tube.

3. Stand one tube in the hot water, the other in a test tube rack.

4. Compare the speeds of reaction.

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• Which temperature speeds up the rate of a chemical reaction between a metal and acid?

• How could you use this property to help preserve food?

Catalysts

• A catalyst is a substance that speeds up a chemical reaction without itself being used up.

• A catalyst allows a chemical reaction to occur at a lower temperature.

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1. Place 1 cm depth of hydrochloric acid in 2 test tubes.

2. To each test tube add a piece of zinc.

3. Add 2 drops of copper sulphate solution to one test tube.

4. Compare the rates of reaction in each tube.

• Under which conditions did the zinc react fastest?

• How do you know the copper is not reacting?

• How could you recover the copper from the reaction mixture once all the zinc had reacted away.

• Why should it be possible to get all the catalyst back after a reaction?

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Summary• Reactive metals react with an acid to produce

hydrogen gas.• The chemical test for hydrogen is that it burns

with a “pop”.• Some metals are more reactive than others, and

react at a faster speed.• Increasing the temperature of an acid increases

the speed of a chemical reaction it has with a metal.

• Catalysts are chemicals which may speed up a chemical reaction, but are not themselves used up during a chemical reaction.

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7. Salts from Acids

By the end of today’s lesson you will be able to:•Describe how to make a salt from a metal and

acid, and write its formula.•Write a balanced equation for the reaction

between an acid and a metal.

Making Salts1. Place a 2 cm depth of hydrochloric acid in

a test tube.2. Add 4 pieces of magnesium. Leave to

react.3. When the reaction stops filter the

solution into an evaporating basin.4. Allow the basin to stand and after 2 days

look at the salt you have made.

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• Salts are formed when a metal reacts with an acid.

• The name of the salt produced is directly related to the names of the reactants.

• Hydrochloric acid produces chloride salts.• Sulphuric acid produces sulphate salts.• Phosphoric acid produces phosphate salts.• Nitric acid produces nitrate salts.

• Why can we not make copper chloride by reacting copper metal with hydrochloric acid?

• Why would it be too dangerous to make sodium chloride by reacting sodium with hydrochloric acid?

• Are salts ionic or covalent compounds? Why?

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Formulae of Salts

• Using valency rules work out the formulae of the following salts:

• zinc nitrate• aluminium nitrate• iron (II) phosphate• iron (III) sulphate

General Reaction

• When a reactive metal reacts with acid, a salt and hydrogen gas are produced.

• This can be shown as a general word equation:reactive metal + acid → salt + hydrogen

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• Copy and complete the following equations, then write balanced chemical equations for each.

• zinc + hydrochloric acid → ? + ?

• zinc + sulphuric acid → ? + ?

• ? + ? → aluminium + hydrogenchloride

• ? + hydrochloric → NO REACTIONacid

Summary• A metal reacts with an acid to produce a

salt and water.• The salt produced is determined by the

acid and metal reacting.• Sulphuric acid produces sulphate salts• Nitric acid produces nitrate salts• Hydrochloric acid produces chloride salts• Phosphoric acid produces phosphate salts

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8. Protective Methods

By the end of today’s lesson you will be able to describe physical, chemical and electrical

methods of protecting iron from rusting.

Electricity and Corrosion

4V d.c.+ -

salt solution + ferroxylindicator

Scrub three nails using iron wool, then set up the equipment shown below and leave for 5 minutes.

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• What happened to the nail connected to the positive terminal? How does this compare to the nail on its own?

• What happened to the nail connected to the negative terminal? How does this compare to the nail on its own?

• How could you use a power supply to protect an iron object from rusting?

• Why should we always use a d.c. power supply?

Cathodic Protection• If a metal object is connected to the negative

terminal of a power supply, then it is protected from corrosion.

• Note a complete circuit is required to allow the current to flow.

• The metal body of a car and oil rigs in the north sea are protected from rusting using cathodic protection.

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power supply

scrap steel

-

+

Sea water contains ionsit acts as an electrolyteto complete the circuitand allow electricalcurrent to flow.

Scratched Coatings

scratched tin-plated

iron

scratched galvanised

iron

scratched painted

iron

salty water+

ferroxyl indicator

salty water+

ferroxyl indicator

salty water+

ferroxyl indicator

Set up the equipment shown below and leave for 5 minutes.

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• If iron is coated completely with another substance this will prevent oxygen and water from reaching the metal’s surface.

• This prevents rusting.

• Scratched less reactive metal coatings INCREASE the rate of rusting.

• Scratched more reactive metal coatings DECREASE the rate of rusting.

Methods of protection

Method Advantage DisadvantageEasily sprayed or

brushed onEasily chipped or

scratchedGood for moving parts Messy, wears off.

Tough and flexible Not heat resistant

No coating needed Uses up a more reactive metal

No coating needed Needs a power supply

Coating which prevents rust even when

scratched

Not suitable for foods containing an acid.

Unreactive metal not easily chipped

Causes faster rusting if the coating is broken.

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Summary• Electricity can be used to protect an iron

object from rusting – this is cathodic protection.

• The negative terminal of a power supply donates electrons to the metal and prevents metal ion formation – preventing corrosion.

• If electrons are removed from a metal, such as when connected to the positive terminal of a power supply, the rate of corrosion is increased.

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9. Electrochemical Series

By the end of today’s lesson you will be able to:• describe the affect differences in reactivity have on the

ability to form ions• explain the different voltages produced when metal pairs are

used to create a cell• describe an experiment to predict the reactivity of a metal

Metal pairs• Pairs of metals in physical contact with each other,

and in the presence of an electrolyte will corrode at different rates.

• The more reactive metal can form ions more easily.

• It will undergo OXIDATION to form ions and release electrons.

• The electrons released are donated to the less reactive metal to prevent it from corroding.

• Any ions of the less reactive metal that may be present gain electrons and are REDUCED to form atoms.

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V

X Y

salt solution

Metal X

Metal Y

Current Direction

Corroding Metal

Corrosion Equation

Cu Zn

Cu Mg

Cu Pb

Zn Mg

Zn Pb

Pb Mg

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Voltages

V

salt bridge

CopperXcopper sulphate

solutionsalt solution

• What general rule can you establish regarding voltage and differences in reactivity?

• For each metal pair, in which direction does the electrical current flow?

Metal Voltage (V)

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Practical Problem• A dull grey metal reacts with acid to produce

hydrogen gas.

• Insufficient acid is available to show it is more or less reactive than zinc and aluminium.

• Describe an experiment to show that unknown metal is less reactive than aluminium and more reactive than zinc.

Summary• Whenever an oxidation reaction occurs, a

corresponding reduction reaction must also be occurring – creating a REDOX reaction.

• A metal higher up the electrochemical series undergoes an oxidation reaction more easily, and hence corrodes more easily.

• By examining the direction electrical current flows between two metals forming a simple cell, we can determine which is higher up the electrochemical series.

• The larger the difference in positions on the electrochemical series, the higher the voltage produced.

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10. Displacement Reactions

By the end of today’s lesson you will be able to predict displacement reactions and write

balanced chemical equations for them.

• Displacement reactions occur when a pure element higher in the reactivity series is added to a solution of a compound containing an element lower in the reactivity series.

• e.g. Magnesium metal can displace hydrogen gas from an acidic solution. An acidic solution contains a compound made up of hydrogen atoms.

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SolutionMetal

Mg Zn Fe Cu

MgSO4

ZnSO4

CuSO4

AgNO3

• Magnesium metal will displace copper from a solution of copper sulphate.

• Word Equation:

• Balanced Chemical Equation:

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• Write balanced chemical equations for all the combinations that would cause a displacement reaction to occur.

Summary• During a displacement reaction, atoms of a

more reactive displace metal ions of a less reactive metal from solution.

• e.g. magnesium atoms will displace copper ions from solution.

• When a displacement reaction occurs, we will see atoms of the less reactive metal form as a solid sample.

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11. Batteries

By the end of today’s lesson you will be able to:• state the energy change inside a battery.

• explain the terms cell, battery, wet cell and dry cell.• describe the features batteries require in different

situations.

Energy Conversion

• When both metals are placed in the acid, an electrical current flows through the voltmeter.

• The metals and acid have formed a cell.

• Chemical energy is converted into electrical energy.

• The hydrochloric acid acts as an electrolyte –completing the circuit.

V-ve +ve

hydrochloric acid

zinc copper

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Wet Cells

• Scientists developed wet cells to produce an electrical current.

• A wet cell contains chemicals dissolved in water.

• The chemicals in a wet cell react together to produce an electrical current.

• If no more chemicals are left to react, then the cell will not produce any electrical current. We say the battery is flat.

manganesedioxide

cotton woolplug

ammoniumchloride solution

zinc strip carbon rod

V

bubbles of gas

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• What energy change is occurring in the wet cell?

• What evidence is there of a chemical change?

• What does the term flat cell mean? Why would this happen?

Dry Cells• Scientists developed dry cells as a portable

source of electricity.• They are usually small, lightweight and do

not spill when turned upside down.• The chemicals inside the dry cell are

present as a paste – usually with water.• The paste acts as an electrolyte allowing

the electrical current to flow.• Longlife cells contain special chemicals that

take longer to be used up before the cell goes flat.

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A zinc carbon dry cell:

A typical dry cell produces 1.5 volts.

Batteries

• A battery is formed when two or more cells are joined together in series.

• Batteries have a wide variety of uses today. There are many different types of battery – each made up of different materials.

• The materials used to make the battery will influence the physical properties of the battery.

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• Look at the poster “Batteries in Action”

• Which examples use wet cells?• Which examples have large

batteries?• Which examples are easy to carry?• Which could not use wet cells?

Summary• In a cell, a chemical reaction occurs that

converts chemical energy into electrical energy.

• Both wet and dry cells contain an electrolyte – completing the circuit and allowing electricity to flow.

• The chemicals contained in a cell dictate its use and application.

• Batteries are a group of cells joined together in series.

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12. Making Electricity

By the end of today’s lesson you will be able to:• explain why a lead/acid battery is said to be

rechargeable• describe the processes occurring in a zinc-copper cell

• Zinc/Carbon and Longlife batteries will eventually go flat – once all the chemicals they contain have been used up.

• Some special types of battery can have the chemicals they contain regenerated.

• If the chemicals are regenerated then more chemical reactions can occur to produce electricity.

• Batteries that can have their chemicals regenerated are called rechargeable batteries.

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Lead-Acid Battery• The dilute sulphuric

acid acts as an electrolyte.

• An electrolyte is a solution containing ions, that are able to move.

• The movement of ions allows the transfer of charge – the conduction of an electrical current.

dilute sulphuric acid

leadplates

• What signs were there of a chemical change during the charging phase?

• Did the bulb stay lit for long? Why?

• Why should lead-acid batteries be fixed firmly in place?

• Why would it not work if you used a covalent solution?

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Car Battery• A car battery is a group of lead-acid cells

joined together in series.• The acid acts as an electrolyte.• The lead plates make the battery very

heavy.• One car battery contains 6 cells. Each cell

produces 2 volts.• Altogether, a car battery produces

12 volts.

Zinc - Copper Cell

V

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• What happens to the copper ions?

• Where do the electrons come from?

• Where do the electrons end up?

• How could you represent this as an equation?

Summary• Once the chemicals inside a battery have been

used up, no more electricity is produced. The battery is said to be flat.

• Rechargeable batteries contain chemicals that may be regenerated or recharged.

• Once regenerated, the chemicals can react and produce electricity once more.

• Lead-acid batteries found in cars are rechargeable.

• Cells are examples of REDOX reactions, one chemical is oxidised whilst the other is reduced.