Reactivity and extracting metals

When metals react with other substances the metal atoms form positive ions. The reactivity of a metal is related to its tendency to form positive ions — the more readily a metal loses its outer electrons, the more reactive it is.

Metals can be arranged in order of their reactivity in a reactivity series. The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids. The non-metals hydrogen and carbon are often included in the reactivity series.

You should be able to recall and describe the reactions, if any, of potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper with water or dilute acids, and where appropriate place these metals in order of reactivity.

The two general equations are:

• metal + water → metal hydroxide + hydrogen • metal + acid → salt + hydrogen

You should also be able to explain how the reactivity of metals with water or dilute acids is related to the tendency of the metal to form its positive ion. Both reactions require the metal atom to lose electrons and become a positive ion, so a metal that loses electrons readily reacts vigorously, and one that does not — copper — does not react at all.

A more reactive metal can displace a less reactive metal from a compound. That gives a second way of ordering metals, and you should be able to deduce an order of reactivity of metals based on experimental results.

Oxidation is the loss of electrons and reduction is the gain of electrons.

This is the wider definition, and it applies to reactions where no oxygen is involved at all. The word OILRIG is the usual memory aid: Oxidation Is Loss, Reduction Is Gain — of electrons.

You should be able to write ionic equations for displacement reactions, and identify in a given reaction, symbol equation or half equation which species are oxidised and which are reduced.

An ionic equation for a displacement reaction leaves out the spectator ions — the ones that are in solution before and after and take no part.

Metals react with oxygen to produce metal oxides. The reactions are oxidation reactions because the metals gain oxygen.

You should be able to explain reduction and oxidation in terms of loss or gain of oxygen. At this stage the definitions are simply:

• oxidation is the gain of oxygen • reduction is the loss of oxygen.

Both happen in the same reaction. When magnesium burns in air, 2Mg + O₂ → 2MgO, the magnesium is oxidised because it has gained oxygen — and the oxygen itself is not described as reduced here, because nothing has lost oxygen. It is in the extraction reactions below that the two appear together.

Unreactive metals such as gold are found in the Earth as the metal itself, but most metals are found as compounds that require chemical reactions to extract the metal.

Metals less reactive than carbon can be extracted from their oxides by reduction with carbon. Reduction involves the loss of oxygen.

Knowledge and understanding are limited to the reduction of oxides using carbon, and knowledge of the details of the processes used in the extraction of metals is not required.

You should be able to identify the substances which are oxidised or reduced in terms of gain or loss of oxygen, and interpret or evaluate specific metal extraction processes when given appropriate information.

In ZnO + C → Zn + CO, the zinc oxide has lost oxygen so it has been reduced, and the carbon has gained oxygen so it has been oxidised. Both happen at once, which is why these are called redox reactions.

Metals can be extracted from molten compounds using electrolysis. Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon, or if the metal reacts with carbon.

Large amounts of energy are used in the extraction process, both to melt the compounds and to produce the electrical current — which is why electrolysis is the expensive option, and why it is used only when reduction with carbon will not work.

Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and cryolite, using carbon as the positive electrode (anode).

You should be able to explain why a mixture is used as the electrolyte, and why the positive electrode must be continually replaced.

Why a mixture: aluminium oxide has a very high melting point on its own. Dissolving it in molten cryolite lowers the melting point of the mixture, so less energy — and less money — is needed to keep the electrolyte molten.

Why the anode is replaced: oxygen is produced at the positive electrode, which is made of carbon. At the high temperature of the cell, that oxygen reacts with the carbon to form carbon dioxide, so the anode gradually burns away and must be replaced.

The Earth's resources of metal ores are limited.

Copper ores are becoming scarce, and new ways of extracting copper from low-grade ores include phytomining and bioleaching. These methods avoid traditional mining methods of digging, moving and disposing of large amounts of rock.

Phytomining uses plants to absorb metal compounds. The plants are harvested and then burned to produce ash that contains metal compounds.

Bioleaching uses bacteria to produce leachate solutions that contain metal compounds.

The metal compounds can be processed to obtain the metal. For example, copper can be obtained from solutions of copper compounds by displacement using scrap iron, or by electrolysis.

You should be able to evaluate alternative biological methods of metal extraction, given appropriate information.

The advantages are that these methods work on low-grade ores that would not be worth mining conventionally, they avoid digging, moving and disposing of large amounts of rock, and they use less energy. The disadvantages are that both are slow — phytomining takes a growing season or more — and the metal compounds still have to be processed afterwards, which itself costs energy.