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Junior Part - Metals

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Metals - Ores

The Earth's crust is composed of a vast number of different compounds, containing both metal and non-metal elements, called ores. Some elements are present in greater amounts than others, and they all have a finite limit, i.e. when we have mined all there is no more will be available.

Here is a table showing the percentage abundance, in the Earth's atmosphere, oceans and crust, of some of the most common elements,

Element % abundance Element % abundance
O 49.4 K 2.40
Si 25.8 Mg 1.94
Al 7.57 H 0.88
Fe 4.70 Cl 0.19
Ca 3.39 C 0.09
Na 2.64 N 0.03

Metal elements are commonly found as metal oxides and the metal itself can be extracted using a one of a variety of chemical reactions. The process of changing the metal ions in a compound to the neutral metal is called reduction. At its simplest this means that oxygen has been removed from the compound; however, the term reduction also applies to the addition of electrons to metal ions (see electrolysis in middle part).

The first of the two most common methods of extraction are a reaction with carbon (or carbon monoxide) in a Blast Furnace. This method is used mainly to extract iron from iron oxide ores and zinc from zinc blende.

The other common method of extraction involves passing a great amount of electricity through the molten ore in process called electrolysis.

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Metals - The Reactivity Series

(1) Reaction of metals with water/steam :

Experimental sheet for the reaction of calcium and water.

Metals such as lithium, sodium,potassium, rubidium and caesium (the alkali metals) react violently with water, too violently to do experimentally, though they can be demonstrated:

The group II metals (also called alkaline earth metals) react less readily and can be used in the laboratory.

Whilst magnesium does react with water it reacts so slowly as to be barely noticeable. As with the alkali metals the reactivity of the group II metals increases as the group is descended, with calcium, strontium and barium all reacting at a reasonable rate.

Exemplar equation -

magnesium + water → magnesium hydroxide + hydrogen gas

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g)

However, if the water is turned into steam and passed over the metal the reaction becomes faster. The reaction with metals such as magnesium and zinc becomes noticeable as a white smoke is given off. This smoke is the metal oxide, instead of the metal hydroxide formed with water. Hydrogen gas is also produced as with water.

Exemplar equation -

zinc + steam → zinc oxide + hydrogen gas

Zn(s) + H2O(g) → ZnO(s) + H2(g)

(2) Reaction of metals with dilute hydrochloric acid :

Experimental sheet for the reaction of metals with dilute hydrochloric acid.

Metals such as sodium and potassium (the alkali metals) react too violently with dilute acids to do experimentally. Other metals, such as copper, silver and gold, will not react with dilute acids at all. Although a mixture of concentrated hydrochloric and nitric acid, known as aqua regia, is powerful enough to form salts with these unreactive metals.

Exemplar equation -

magnesium + hydrochloric acid → magnesium chloride + hydrogen gas

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

(3) Reaction with other metal ions :

Experimental sheet for the reaction of metals with metal ions.

When a metal is reacted with an aqueous solution of another metal ion a reaction will occur only if the metal is higher in the reactivity series than the metal ion. This provides a relatively easy way to obtain a Reactivity Series for any group of metals.

There is another way which involves forming electrical cells with different metals and measuring the potential differences produced, though this can give a slightly different series than the chemical reactions do.

Exemplar equation -

magnesium + silver nitrate → magnesium nitrate + silver

Mg(s) + 2AgNO3(aq) → Mg(NO3)2(aq) + 2Ag(s)

Try balancing these equations,

Li(s) + CuCl2(aq)LiCl(aq) + Cu(s)
Al(s) + CuSO4(aq)Al2(SO4)3(aq) + Cu(s)

(4) Relationship to extraction of metals :

The method used to extract a metal from its ore is linked very closely to that metal's position in the Reactivity Series. The more reactive a metal is the more it wants to form compound and therefore the harder it is to isolate the pure metal from its compounds.

The most reactive metals require electricity to extract the metal; less reactive metals can be reduced with coke in a Blast furnace; the least reactive metals are found as the pure metal and require very little purification.

(5) Summary of Reactivity Series :

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Metals - Iron

(1) Iron(II) and Iron(III) compounds :

Iron(II) and iron(III) compounds can be separated by their colours most of the time. Iron(II) compounds are generally green in colour and iron(III) compounds are generally redin colour (there are plenty of exceptions though).

There is also a chemical test that can distinguish between the two sets of compounds.

When an aqueous solution of iron ions is reacted with an aqueous solution of hydroxide ions (e.g. sodium hydroxide or ammonium hydroxide) a coloured precipitate will be formed. A green precipitate is formed with iron(II) ions and a red-brown precipitate is formed with iron(III) ions.

The precipitates are insoluble iron hydroxides - Iron(II) hydroxide (green) and iron(III) hydroxide (red-brown).

General ionic equations are given below for the two reactions -

Fe2+(aq) + 2-OH(aq) → Fe(OH)2(s)

Fe3+(aq) + 3-OH(aq) → Fe(OH)3(s)

(2) The Blast Furnace :

The raw materials needed for the blast furnace are coke (carbon), limestone (calcium carbonate, CaCO3), heated air and the iron ore itself (e.g. haematite, which is iron(III) oxide, Fe2O3).

The solid raw materials are fed into the top of the furnace and the hot air is pumped in from near the bottom. This means that the hottest part of the reaction occurs near the bottom of the Blast furnace.

The coke burns in heated air to partially oxidise to carbon monoxide. This carbon monoxide is the primary reducing agent for the iron ore (though the coke itself will also reduce the iron ore). The coke also reduces the carbon dioxide, produced in various reactions in the furnace, to produce more carbon monoxide.

2C(s) + O2(g) → 2CO(g)

C(s) + CO2(g) → 2CO(g)

Thus the iron ore is reduced to iron metal which is produced at such a temperature that it is molten. The carbon dioxide gas is simply vented off.

Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)

There are quite a lot of impurities present in the ore, mainly sand (silicon dioxide, SiO2). This is removed by the limestone. When the limestone is heated in the furnace it decomposes to lime (calcium oxide) and oxygen. This calcium oxide reacts with the sand to produce a slag, calcium silicate, CaSiO3.

CaCO3(s) → CaO(s) + CO2(g)

CaO(s) + SiO2(s) → CaSiO3(l)

The molten iron sinks to the bottom of the furnace with the liquid slag floating on top of it. It is then a simple matter to allow the slag to flow away and obtain the iron from beneath it.

(3) Rusting :

Rusting is a very specific reaction of iron, where the metal is turned into hydrated iron(III) oxide, known as rust.

For iron to rust the metal must be in contact with air and water. Without either one of these the metal will remain intact and not corrode.

There are a number of separate stages involved, such as the oxidation of iron to iron(II) ions and then the oxidation of iron(II) ions into the iron(III) oxide,

Equation -

4Fe2+(aq) + O2(g) + 2H2O(l) → 4Fe(OH)3(s) or 2Fe2O3.3H2O(s)

These processes involve the movement of electrons and as such the rusting process can be speeded up by having salt present in the water, as this allows greater conductivity.

Rust prevention -

Anything that prevents air and/or water from contacting the surface of the iron will prevent the metal from rusting; as does anything that reacts faster then the iron, i.e. is higher in the reactivity series.

In the latter case the oxygen and water will preferentially react with the more reactive metal, to form metal oxides, before the iron can react. This allows the overall strength of the metal to remain the same.

Common methods for rust prevention involve painting the metal; covering the metal in oil or grease; coating the metal in zinc, called galvanising; attaching blocks of a more reactive metal, such as magnesium or zinc, called sacrificial protection; alloying the iron with other metals, such as chromium and nickel, or changing the carbon content of the iron to create steel.

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Metals - Alloys

Metals are mixed together to create alloys. These alloys have better physical properties then the individual metals, such as higher melting points, greater mechanical strength or increased resistance to corrosion.

Steel -

Steel is an alloy created by blowing oxygen through molten iron produced in the Blast furnace. The oxygen reacts with the carbon impurity in the iron, turning it into carbon dioxide, which is vented off. The amount of carbon can be very carefully controlled, giving a wide range of different steels. For example, mild steel has about 0.25% carbon and is hard and strong; whilst high carbon steel has about 1.5% carbon and is harder but more brittle.

Other metals can also be added to the steel as it is made to make an alloy such as stainless steel. For this chromium and nickel are added which form unreactive oxides on the surface of the iron and prevent the rusting process starting.

Others -

A few other common alloys are bronze, which is a mixture of copper and tin; brass, which is a mixture of copper and zinc; solder and pewter, which are mixtures of tin and lead.

Even gold used for jewelry is alloyed with other metals such as zinc orb nickel to produce normal gold as well as white gold.

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written by Dr Richard Clarkson : © Saturday, 1 November 1997

updated : Wednesday, 11th July, 2012

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