Specification

The reactivity series of metals

The reactivity series of metals is a chart showing metals in order of decreasing reactivity. In general, the more reactive a metal is:

the more vigorous its reactions are
the more easily it loses electrons in reactions to form positive ions (cations)

The table summarises some reactions of metals in the reactivity series. Hydrogen is shown for comparison.

Reactions of metals with water

When a metal reacts with water, a metal hydroxide and hydrogen are formed

Metal + water → metal hydroxide + hydrogen

For example, sodium reacts rapidly with cold water, melting into a ball, and 'fizzing' about the surface:

sodium + water → sodium hydroxide + hydrogen

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

In general, the more reactive the metal, the more rapid the reaction is.

Reactions with steam

Metals that react slowly with cold water can react quickly with steam. In these reactions a metal oxide and hydrogen are produced.

Metal + steam → metal oxide + hydrogen

For example, magnesium reacts slowly with cold water. However, if steam is passed over hot magnesium, a vigorous reaction occurs:

Magnesium + steam → magnesium oxide + hydrogen

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Question

State the difference between the products formed when calcium reacts with cold water and when it reacts with steam.

Calcium hydroxide forms when it reacts with water, but calcium oxide forms when it reacts with steam.

Reactions of metals with dilute acids

When a metal reacts with a dilute acid, a salt and hydrogen are formed.

Metal + acid → salt + hydrogen

For example, magnesium reacts rapidly with dilute hydrochloric acid:

Magnesium + hydrochloric acid → magnesium chloride + hydrogen

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

The more reactive the metal, the more rapid the reaction is. A metal below hydrogen in the reactivity series will not react with dilute acids.

Question: Platinum is placed below gold in the reactivity series. Predict its reaction with dilute acids and explain your answer.

Platinum will not react with dilute acids. Metals below hydrogen in the reactivity series do not react with dilute acids, and both gold and platinum are placed below hydrogen.

Hydrogen is always given off when a metal reacts with water, steam or a dilute acid.

n the reactions of metals with water, steam and acids, the metals lose electrons and form cations. The metal is oxidised and the water is reduced.

A metal's relative tendency to form cations and its resistance to oxidation are both related to its position in the reactivity series. In general:

the higher up a metal, the greater the tendency to form cations

the lower down a metal, the greater its resistance to oxidation

Metals and displacement reactions

Displacement in solutions

A more reactive metal can displace a less reactive metal from its compounds. For example, magnesium is more reactive than copper. It displaces copper from copper sulfate solution:

magnesium + copper sulfate → magnesium sulfate + copper

Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)

In this displacement reaction:

magnesium becomes coated with copper
the blue colour of the solution fades as blue copper sulfate solution is replaced by colourless magnesium sulfate solution

Determining a reactivity series

A reactivity series can be deduced by carrying out several displacement reactions. A piece of metal is dipped into a salt solution. Different combinations of metal and salt solution are tested. The table shows the results of one of these investigations.

	Magnesium sulfate solution	Copper sulfate solution	Iron sulfate solution	Reactions
Magnesium	Not done	Brown coating	Black coating	2
Copper	No visible reaction	Not done	No visible reaction	0
Iron	No visible reaction	Brown coating	Not done	1

Question: Use the results in the table to deduce an order of reactivity, starting with the most reactive metal.; The order of reactivity is: magnesium > iron > lead. This is because magnesium could displace lead and iron, iron could only displace lead, but lead could not displace magnesium or iron.
Question

A metal cannot displace itself from a solution of one of its salts. There would be no reaction, so these combinations were not done.

Displacement reactions as redox reactions - Higher

A balanced equation for the reaction between magnesium and copper sulfate solution can be written in terms of the ions involved:

Mg(s) + Cu²⁺(aq) + SO₄^2-(aq) → Mg²⁺(aq) + SO₄^2-(aq) + Cu(s)

Sulfate ions, SO₄^2-, appear on both sides of the equation. They do not take part in the reaction and are called spectator ions. The equation can be rewritten without them:

Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)

This equation is an example of a balanced ionic equation. It can be split into two half equations:

Mg(s) → Mg²⁺(aq) + 2e^- (oxidation)

Cu²⁺(aq) + 2e^- → Cu(s) (reduction)

Notice that:

magnesium atoms lose electrons - they are oxidised
copper ions gain electrons - they are reduced

Reduction and oxidation happen at the same time, so the reactions are called redox reactions.

Displacement reactions are just one example of redox reactions. Electrolysis reactions are also redox reactions.

Note that the reaction of metals with acids can also be described as a displacement reaction or a redox reaction. Only metals above hydrogen in the reactivity series will react and displace hydrogen from acids.

For example:

Magnesium + sulfuric acid → zinc sulfate + hydrogen

Written with the ions involved:

Zn(s) + 2H⁺(aq) + SO₄^2-(aq) → Zn²⁺(aq) + SO₄^2-(aq) + H₂(g)

Removing the spectator ions this becomes an ionic equation:

Zn(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂(g)

The half equations are:

Zn(s) → Zn²⁺(aq) + 2e^-

2H⁺(aq) + 2e^- → H₂(g)

Extracting iron and copper

Ores

Unreactivemetals such as gold are found in the Earth's crust as the uncombined elements. However, most metals are found combined with other elements to form compounds.

An ore is a rock that contains enough of a metal or a metal compound to make extracting the metal worthwhile:

low-grade ores contain a small percentage of the metal or its compound
high-grade ores contain a larger percentage

Most metals are extracted from ores found in the Earth's crust. It is more expensive and wasteful to extract a metal from a low-grade ore, but most high-grade ores have already been used.

Extraction methods

The extraction method used depends upon the metal's position in the reactivity series. In principle, any metal could be extracted from its compounds using electrolysis. However, large amounts of electrical energy are needed to do this, so electrolysis is expensive.

If a metal is less reactive than carbon, it can be extracted from its compounds by heating with carbon.

Metal oxide + carbon → metal + carbon dioxide

For example, molten copper can be produced from copper oxide by heating with carbon:

Copper oxide + carbon → copper + carbon dioxide

2CuO(s) + C(s) → 2Cu(l) + CO₂(g)

Copper oxide is reduced as carbon is oxidised, so this is an example of a redox reaction.

Remember:

oxidation is the gain of oxygen by a substance
reduction is the loss of oxygen by a substance
a redox reaction involves the loss and gain of oxygen

Note: the impure copper is purified by electrolysis.

The table summarises the extraction methods used for different metals.

Extracting iron

As iron is below carbon in the reactivity series it can be displaced from its compounds by heating with carbon. Iron is extracted from iron ore in a large container called a blast furnace. Iron(III) oxide is reduced to molten iron when it reacts with carbon. The overall reaction is:

Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)

The iron oxide is reduced and the carbon is oxidised.

These reactions happen because carbon is more reactive than iron, so it can displace iron from iron compounds. Extracting a metal by heating with carbon is cheaper than using electrolysis.

Question

Write a balanced equation for the reaction between tin(IV) oxide and carbon, forming molten tin and carbon dioxide.

SnO₂(s) + C(s) → Sn(l) + CO₂(g)

Extracting aluminium

Aluminium is more reactive than carbon so it must be extracted from its compounds using electrolysis. Even though aluminium is more abundant than iron in the Earth's crust, aluminium is more expensive than iron. This is mainly because of the large amounts of electrical energy used in the extraction process.

Electrolysis of aluminium oxide

The electrolyte

Aluminium ore is treated to produce pure aluminium oxide. The electrolytes used in electrolysis are ionic compounds:

in the molten state, or
dissolved in water

Aluminium oxide is insoluble in water, so it must be molten to act as an electrolyte. However, the melting point of aluminium oxide is high. A lot of energy must be transferred to break its strong ionic bonds, and this is expensive. To reduce costs, powdered aluminium oxide is dissolved in molten cryolite. This ionic compound melts at a lower temperature than aluminium oxide, reducing costs.

The electrolysis process

The diagram shows an electrolysis cell used to extract aluminium. Both electrodes are made of graphite, a form of carbon with a high melting point and which conducts electricity.

During electrolysis:

at the cathode, aluminium ions gain electrons and form aluminium atoms
at the anode, oxide ions lose electrons and form oxygen gas

The oxygen reacts with the carbon anodes, forming carbon dioxide. So the anodes gradually wear away. They must be replaced frequently, adding to the cost of producing aluminium.

Question: Higher - Explain, with the help of a half equation, how oxide ions are oxidised during the electrolysis of aluminium oxide.; The half equation is: 2O^2- → O₂ + 4e^-.
It shows that oxide ions lose electrons, and oxidation is loss of electrons.
Question: Explain, with the help of a half equation, how aluminium ions are reduced during the electrolysis of aluminium oxide.

The half equation is: Al³⁺ + 3e^- → AlIt shows that aluminium ions gain electrons, and reduction is gain of electrons.

Chemistry For GCSE Edexcel by KANAYATI ®

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11.Obtaining and Using Metals GCSE EDEXCEL CHEMISTRY