Metals, Making Electricity and Corrosion

Metals The job that a metal is used for is

determined by its physical and chemical properties.

Physical properties of metals include :StrongMalleable - can be beaten into shapeDuctile – can be drawn into wires Conductors of heat and electricityShiny (some are shiny, others are dull

because they have already reacted with the air)

If the physical properties don’t quite match the requirements an alloy can be made.

An alloy is made by mixing metals with other metals or with non-metals by melting the mixture and then allowing it to cool.

Examples of alloys are:Solder, lead and tin, which has a lower mpBrass, copper and zinc, which is hard wearingSteel, iron and carbon, which is strong

Chemical properties are determined by how a substance behaves when it comes into contact with other substances.

The reactivity of a metal is determined by how it reacts in the presence of water, acid and oxygen.

Metal + Water Metal + Hydrogen

Hydroxide Only the most reactive metals will react with

water, some only react slightly.

Metal + Acid Salt + Hydrogen

Copper, mercury, silver and gold do not react with acid or water.

Potassium, sodium and calcium are too reactive to add to acid.

Metal + Oxygen Metal Oxide

Metals are placed in order of reactivity in the reactivity series, where the most reactive are nearest the top and the least reactive are found at the bottom.

The reactivity of a metal determines whether it will be found combined or uncombined in the earth’s crust.

Native metals are metals which are found uncombined in the earth’s crust. Silver and gold are examples of native metals.

Ores are naturally occuring compounds of metals from which metals can be extracted.

Extraction of metals The extraction method used is determined by the

reactivity of the metal in the compound. The more reactive the metal, the more difficult it is

to extract. The extraction methods used are:

Heating metal oxides – only very unreactive metals can be obtained this way.

Heating metal oxides with carbon – this method is used to extract metals below aluminium.

Heating with carbon monoxide – this method is used in the blast furnace to extract iron from iron oxide

Electrolysis – this method is used to extract reactive metals above zinc.

Chemical reactions can produce electricity. A cell (often referred to as a battery)

contains chemicals which react to make electricity.

The three main things needed for a cell are:Positive electrodeNegative electrodeElectrolyte – a substance which allows

electricity to flow (completes the circuit)

Making Electricity

Batteries run out when the chemicals they contain are all used up.

In rechargeable batteries the chemicals are not used up and can be regenerated by recharging the battery with electricity.

Making electricity with 2 different metals

Electricity can be produced when 2 different metals are dipped in an electrolyte and connected with a wire.

There is a flow of electrons in the wire

from the more reactive metal to the less reactive metal (and ions flow through the electrolyte)

The electrochemical series on page 7 of the data booklet lists the ion-electron equations for a number of reduction reactions.

If a voltmeter is placed in the circuit, a different voltage is obtained when different metals are used.

The bigger the gap between the metals in the electrochemical series, the larger the voltage obtained.

Displacement Reactions A displacement reaction occurs when a metal

higher in the electrochemical series is added to a solution containing a metal lower in the electrochemical series.

The metal lower in the series is displaced or ‘pushed out’ of solution and the more reactive metal takes its place.

(When looking at reactions of metals with acids only metals above hydrogen in the electrochemical series will displace it from solution)

Electricity from different metals in solutions of their own ions

The electrons flow from the more reactive metal to the less reactive metal.

The ion bridge is needed to complete the circuit (it allows the ions to flow)

The above reaction can be split into 2 half cells. The zinc electrode loses electrons to form zinc ions

(electrode gets lighter) Zn Zn2+ + 2e-

The copper ions in solution gain these electrons to form copper atoms (electrode get heavier)

Cu + 2e- Cu

The ion electron equations for these reactions are found on page 7 of the data booklet.

The reaction for the more reactive metal is reversed whereas the less reactive metal is as written.

Electricity from cells which contain a non-metal

The electrochemical series has some reactions that involve non-metals.

In this cell: The zinc loses electrons to

form zinc ions. Zn Zn2+ + 2e-

The iodine atoms gain electrons to form iodide ions.

I2 + 2e- 2I-

Reduction and Oxidation Reduction is the gain of electrons by a reactant Oxidation is the loss of electrons (or gain of

oxygen) A redox reaction is when both reduction and

oxidation occurs. A displacement reaction is an example of a redox

reaction. The more reactive metal loses electrons to change

from atoms to ions (oxidation) The less reactive metal gains electrons to change from

ions to atoms (reduction)

Corrosion Metals corrode over time. When this occurs

the surface of the metal changes from an element into a compound.

Corrosion of iron is also known as rusting.

For iron to rust it needs to be in contact with water and oxygen.

The presence of salt or an ionic solution speeds up the process.

When rusting occurs the iron atoms lose electrons to form iron ions.

Water and oxygen molecules gain these electrons to form hydroxide ions.

(see page 7 of the data booklet to find the ion electron equations)

Ferroxyl indicator is used to indicate rusting (it turns from green to blue).

Corrosion Prevention To prevent corrosion we need to put a barrier between

water and oxygen and the metal’s surface. Examples are :

Painting Oiling and greasing Coating with plastic Electroplating – coating the metal with a layer of another metal Cathodic protection – connecting the metal to be protected to the

negative electrode Galvanising – dipping the metal in molten zinc and allowing it to

harden Sacrificial protection – attaching a more reactive metal to a less

reactive metal. The more reactive metal sacrifices itself by giving away its electrons to the less reactive metal.

Physical vs Chemical protection Physical protection – if a break/gap appears then

the metal is no longer protected. (If a less reactive metal is used to electroplate a

more reactive metal if a scratch occurs the metal being protected corrodes faster as it gives its electrons away.)

Chemical protection – if a break/gap appears then the metal is still protected.

(Examples are electroplating with a more reactive metal or sacrificial protection)