Metal Corrosion

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The structure of metals

The arrangement of the atoms

•Metals are giant structures of atoms held together by metallic bonds

•Metallic bonds - atoms surrounded by delocalised electrons ("sea of electrons")

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The structure of metals

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The structure of metals

• The electrons can move freely within the metallic bonds

• each electron becomes detached from its parent atom - the electrons are delocalised

• The metal is held together by the strong forces of attraction between the positive nuclei and the delocalised electrons

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The structure of metals

• The more electrons you can involve - the stronger the attractions

• Transition metals tend to have particularly high melting points and boiling points

• They can involve many delocalised electrons

• Metallic character increases as we move to the right of the periodic table

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The Corrosion of Metals

Most metals react with their surroundings to form oxides and hydroxides

E.g. copper forms copper hydroxide

iron rusts to form iron oxide

Sodium corrodes in air to form a layer of sodium oxide on the metal surface

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Only a few metals can resist corrosion:•Gold & Platinum (don’t react with oxygen)•Stainless steel (iron + carbon + chromium (form stable film against corrosion)

Metals That Don’t Corrode

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• Corrosion - is the changing of the surface of a metal element into a compound (oxide)• Silver + oxygen silver oxide

• Rusting - is the special name given to

the corrosion of iron• Iron + oxygen iron (III) oxide

Definitions

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What Happens When Iron Rusts?

• Rusting involves iron atoms (metal) losing electrons to

form ions. It occurs over two steps

Fe(s) Fe2+(aq) + 2e-

Fe2+(aq) Fe3+

(aq) + e-

• The electrons lost from the iron are accepted by the water

and oxygen

2H2O(l) + O2(g) + 4e- 4OH-(aq)

oxidation

oxidation

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Experiment: Conditions For Rusting

Experiment Air water electrolyte Did Rusting Occur?

1 yes no no no

2 no yes no no

3 yes yes Yes (small)

yes

4 yes yes yes yes

Dry air Boiled water + oil

water salty water

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Experimental Conclusions

1. Water and oxygen are both required for rusting

to take place

2. An electrolyte must also be present and the

speed of rusting is increased

3. An electrolyte is an ionic substance dissolved in

water and provides free ions to carry a current

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Corrosion and the Reactivity series

PotassiumSodiumLithiumCalciumMagnesiumAluminumZincIronTinLeadCopperMercurySilverGold

Increasing speed of corrosion

The uses of metals depends on their position in the reactivity series

1. Why don’t we make nails from potassium or gold?

2. What are the benefits and disadvantage of having silver jewelry?

3. Why are bridges not built from stainless steel?

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Corrosion and the Reactivity series

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Cell PotentialsWe can calculate the magnitude of electron flow by measuring the voltage

Anode (-) Cathode (+)

e-

Allows ion flow without mixing solutions

Allows ions to pass between solutions but doesn’t allow the solutions to mix

Figure 21.5

A voltaic cell based on the zinc-copper reaction

Displacing electrons from zinc to copper

Zn(s) Zn2+(aq) + 2e- Cu2+(aq) + 2e- Cu(s)

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Summary of Corrosion and reduction potential

•E cell = E anode – E cathode

•Anode (-)

•Cathode (+)

•MX (s) M+ (aq) + X- (aq)

•Oxidation: lose electrons at anode (-)

•M M+ + e-

•Reduction: gain electrons at cathode (+)

•X + e- X-

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Displacement Reactions

• Chemical reaction in which a less reactive element is replaced in a compound by a more reactive one

• For example, the addition of zinc metal to a solution of copper(II) sulphate displaces copper metal:

• Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

• The copper is taken out of the solution and is deposited as a solid

• In the electrochemical series an element can be displaced from a compound by any element above it in the series

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Corrosion and the Reactivity series

• The more negative the metal in the series the more reactive it is (its reaction is fast and more exothermic) - it wants to lose electrons to form an oxide

• Therefore the reverse reaction becomes difficult (oxide -> pure metal)

• Hard to extract a metal from its ore (stable)

• The pure metal is also more susceptible to corrosion with oxygen and water

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Electronegativity – the ability to gain or lose electrons

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Atomic radius

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Atomic radius

• Elements with a large atomic radius are high on the reactivity series (lose electrons easily)

• The number of protons increases across a period as does the effective nuclear charge

• Electrons within a shell cannot shield each other from the attraction to protons

• This causes the atomic radius to shrink (less reactive)

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Summary of trends in the Reactivity Series

High on activity series (-)•Large atomic radius•LHS of periodic table (lose electrons)

Low on activity series (+)•Smaller atomic radius•RHS periodic table (gain electrons)

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Protection Against Corrosion

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What Happens When Metals Corrode?

• Corrosion is a chemical reaction• It involves the metal atoms losing electrons

Fe(s) Fe2+ + 2e-

• Metals corroding are examples of oxidation reactions

Cu(s) Cu2+ + 2e-

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Methods of Protection

There are two main methods of protecting metals

from corrosion:

Physical protection - placing a barrier to water and

oxygen on the surface of the metal

Chemical Protection - providing the metal with a

source of electrons to prevent oxidation

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Physical Protection

Plastic coating

Oil and grease

Paint

Tin plating

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Questions

What are the advantages and disadvantage of each type of physical protection:

a) Plasticb) Oil and greasec) Paintd) Tin-plating

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Physical Protection

Electro-plating - coating the iron with a less

reactive metal• Gold, silver, chromium, and copper

are a few of the metals that can be used

as coatings

•The metal item to be plated is used as

the negative (-) electrode

•The item is placed in a solution of the

metal coating

•The metal in solution is reduced from

ions to atoms and deposited on the

metal item

gold electrode

Power pack

+ -

gold ion solution

item to be coated

Au2+(aq) + 2e- Au(s) 29

Problems with Electroplating

electrolyte solution

gold iron

Electron flow• When the electroplating is broken an electrochemical cell is set up

• If iron is higher in the electrochemical series than the coating metal electrons flow away from the iron

• Rusting is speeded up

V

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Physical Protection

Galvanizing - coating the iron with zinc

• The iron or steel is dipped in

molten zinc

• The zinc provides a physical

coating on the surface

• If the zinc coating is damaged

sacrificial protection occurs. This

also is a form of chemical protection

(see later)

• This is expensive and requires

special equipment to achieve31

When a metal corrodes it loses electrons

Mg Mg2+ + 2e-

Chemical protection supplies the metal with a flow

of electrons by two methods:

1. Direct current power supply

2. Metal higher in the reactivity series

Chemical Protection

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Direct Electrical Protection

This involves connecting the iron to the negative terminal of a battery or power supply

Connecting the negative terminal of a car battery to the car body slows corrosion

Ocean liners use direct electrical protection when docked

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Zn2+ + 2e-

Zn

Sacrificial Protection

What happens when you connect a more

reactive metal to a less reactive metal in a

simple cell?

zinc

electrolyte solution

iron

V

Electron flow

The electrons flow from the metal higher up the electrochemical series to the metal lower

zinc iron

The flow of electrons prevents the iron from rusting

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Sacrificial Protection

Mg

steel pipeline

• To protect underground pipes from rusting the pipe is connected to scrap magnesium by a wire

• The magnesium sacrificially corrodes giving electrons to the iron pipe

• Rusting is slowed down

• The scrap magnesium needs regularly replaced 35

Questions

1. What happens to a metal when it corrodes?

2. How can we prevent this loss of electrons from a corroding metal?

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Questions1. Name the two types of corrosion

prevention.2. What two chemicals are required

for rusting to occur?3. What else is required for rusting

to occur?

Clue: It contains charged particles that allow

electrons to be lost more easily from the

iron 37

Questions

1. What metals are used to plate steel and

iron?

2. Where do most of these metals sit in the

electrochemical series?

3. What terminal on the power pack is the

metal to be coated connected?

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Questions1. What metal is used in galvanising?2. What type of protection does this offer

the iron/steel?3. Give 3 examples of galvanising being

used to prevent rusting.4. Which metal is higher in the

electrochemical series - iron or zinc?

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Questions

1. When does electro-plating prevent rusting?

2. When does electro-plating cause rusting to occur faster?

3. When rusting occurs what metal is losing electrons?

4. What metal is being protected from corrosion?

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Questions

1. What do you use to provide direct-electrical

protection to a metal surface?

2. Are there any disadvantages to this method?

3. Give two examples of where this is used

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Questions

1. Where must the metal used for sacrificial

protection be on the E.C. series?

2. What is the next most suitable metal for

sacrificial protection after zinc?

3. Which metal would provide the best

protection out of these two?

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Ferroxyl indicator + iron nail

Ferroxyl indicator + iron nail wrapped in magnesium

Ferroxyl indicator + iron nail wrapped in copper

Testing For Protection Against Rusting

Set up the following experiments:

1. Half fill the petri dish with warm agar solution.

2. Add 5 drops of ferroxyl indicator and 3 drops of

phenolpthalein indicatior into the dish.

3. Gently stir.

4. Place a nail in each dish as shown below.

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Ferroxyl indicator + iron nail

Ferroxyl indicator + iron nail wrapped in magnesium

Ferroxyl indicator + iron nail wrapped in copper

Testing For Protection Against Rusting

• After leaving the experiment for 20-30

minutes, draw a before and after diagram. What

did you observe happening?

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Ferroxyl indicator + iron nail

Ferroxyl indicator + iron nail wrapped in magnesium

Ferroxyl indicator + iron nail wrapped in copper

Testing For Protection Against Rusting

Set up the following experiments:

1. Half fill the petri dish with warm agar solution.

2. Add 5 drops of ferroxyl indicator and 3 drops of

phenolpthalein indicatior into the dish.

3. Gently stir.

4. Place a nail in each dish as shown below.

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