Metals

Physical properties of metals: what makes a metal a metal

Metals are the elements on the left and in the middle of the periodic table, and they share a set of physical properties. They are good conductors of heat and good conductors of electricity, they are malleable (they can be hammered or rolled into shape without cracking) and ductile (they can be drawn out into wires), and they generally have high melting and boiling points.

The reason for all of this is the metallic structure. A metal is a lattice of positive ions surrounded by a sea of delocalised electrons. Because the electrons are free to move through the structure, they can carry charge and carry heat, which is why metals conduct. Because the layers of positive ions are held by a non-directional attraction, the layers can slide over one another when the metal is hit, so the metal bends rather than shatters.

Non-metals are the opposite in almost every way. They are poor conductors of heat and electricity (the one important exception is graphite, a non-metal that does conduct), they are brittle, so they crack or shatter when hammered, and they usually have low melting and boiling points, because the forces holding their molecules together are weak.

There are exceptions worth knowing. Mercury is a metal but is a liquid at room temperature, and carbon and silicon dioxide are non-metals with extremely high melting points because they are giant covalent structures.

Chemical properties: how metals react with oxygen, water and acids

With oxygen. Almost every metal reacts with oxygen to form a metal oxide. A reactive metal such as magnesium burns with a brilliant white flame: 2Mg + O2 -> 2MgO. Sodium burns with an orange flame: 4Na + O2 -> 2Na2O. Less reactive metals react more slowly, which is why copper slowly turns dull rather than bursting into flame. The oxides of metals are basic, so they neutralise acids.

With cold water. Only the most reactive metals react with cold water, and they produce a metal hydroxide and hydrogen. Sodium fizzes across the surface and may ignite: 2Na + 2H2O -> 2NaOH + H2. Calcium fizzes steadily and sinks: Ca + 2H2O -> Ca(OH)2 + H2. Potassium is even more reactive than sodium and catches fire with a lilac flame.

With steam. Metals that are not reactive enough to be attacked by cold water, such as magnesium, zinc and iron, will still react with steam, giving a metal oxide and hydrogen instead of a hydroxide. For example Mg + H2O -> MgO + H2, and 3Fe + 4H2O -> Fe3O4 + 4H2. The metal glows and hydrogen gas is collected.

With dilute acids. Metals above hydrogen in the reactivity series react with dilute hydrochloric acid to give a salt and hydrogen: Mg + 2HCl -> MgCl2 + H2, Zn + 2HCl -> ZnCl2 + H2 and Fe + 2HCl -> FeCl2 + H2. The more reactive the metal, the more vigorous the fizzing. Copper, silver and gold are below hydrogen, so they do not react with dilute acids at all.

• Metal + oxygen -> metal oxide, and the oxide is basic.
• Metal + cold water -> metal hydroxide + hydrogen (only potassium, sodium and calcium).
• Metal + steam -> metal oxide + hydrogen (magnesium, zinc, iron).
• Metal + dilute acid -> salt + hydrogen (only metals above hydrogen in the reactivity series).

Uses of metals: matching the property to the job

Every sensible use of a metal can be justified by one of its physical properties, and exam questions almost always want the property as well as the use. Aluminium is used in the manufacture of aircraft because it has a low density, so the aircraft is lighter and uses less fuel. It is also used in overhead electrical cables because of its low density together with its good electrical conductivity - a copper cable of the same current-carrying capacity would be far too heavy to hang between pylons.

Aluminium is used for food containers, drink cans and cooking foil because it is resistant to corrosion. It does react with oxygen, but the layer of aluminium oxide that forms is thin, tough and sticks firmly to the surface, so it seals the metal underneath and stops the reaction going further.

Copper is used in electrical wiring because of its good electrical conductivity and its ductility, which means it can be drawn into the thin flexible wires needed inside a cable. Copper is also a very good conductor of heat and is fairly unreactive, so it is used for water pipes and for the base of cooking pans. Iron and steel are used for bridges, buildings and car bodies because they are strong and cheap, but they must be protected from rusting.

Alloys: why a mixture beats a pure metal

An alloy is a mixture of a metal with other elements. Mixing is not a chemical reaction, so the metal keeps its own properties, but the mixture is usually far more useful than the pure metal. Brass is a mixture of copper and zinc, and stainless steel is a mixture of iron with other elements such as chromium, nickel and carbon.

In a pure metal the atoms are all the same size and pack into regular layers. When the metal is bent, whole layers of atoms slide over each other, which is why a pure metal such as copper is soft and easily shaped.

In an alloy the atoms of the different elements are different sizes. The larger and smaller atoms distort the regular layers, so the layers can no longer slide over each other. The alloy is therefore harder and stronger than the pure metal, and it is more useful for jobs where the metal has to keep its shape and resist wear.

You can recognise an alloy from a diagram of its structure: a pure metal shows a regular pattern of identical circles, while an alloy shows circles of at least two different sizes mixed together. Stainless steel is used for cutlery because it is hard and resists rusting, and brass is used for musical instruments and door fittings because it is hard, resists corrosion and looks attractive.

• Alloy = a mixture of a metal with other elements.
• Brass = copper + zinc; stainless steel = iron + chromium + nickel + carbon.
• Pure metal: identical atoms in layers that slide, so it is soft and malleable.
• Alloy: different sized atoms lock the layers, so it is harder and stronger.

The reactivity series and displacement reactions

The reactivity series puts metals in order of how vigorously they react. Learn it from most to least reactive: potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold. Carbon and hydrogen are included because they are the two useful reference points: a metal above carbon cannot be extracted with carbon, and a metal above hydrogen will react with dilute acids.

The series is really a list of how readily each metal loses electrons to form a positive ion. Potassium loses its outer electron almost instantly, so it is extremely reactive, while gold holds on to its electrons so tightly that it hardly reacts at all.

This is why displacement reactions happen. If a more reactive metal is placed in a solution containing the ions of a less reactive metal, the more reactive metal takes the place of the less reactive one. Zinc displaces copper from copper(II) sulfate: Zn + CuSO4 -> ZnSO4 + Cu. Iron does the same: Fe + CuSO4 -> FeSO4 + Cu. Copper displaces silver from silver nitrate: Cu + 2AgNO3 -> Cu(NO3)2 + 2Ag.

In terms of electrons, the more reactive metal is oxidised (it loses electrons and becomes a positive ion) and the ions of the less reactive metal are reduced (they gain those electrons and become the metal). The ionic equation for the zinc and copper(II) sulfate reaction is Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s).

One metal looks like an exception. Aluminium appears unreactive, and does not seem to react with water or dilute acid as its position in the series would suggest. The reason is that a thin, tough layer of aluminium oxide forms immediately on the surface and sticks tightly to it, protecting the metal underneath. Aluminium is in fact a very reactive metal.

• Zinc + copper(II) sulfate -> zinc sulfate + copper: a pink-brown deposit forms and the blue colour fades.
• Copper + silver nitrate -> copper(II) nitrate + silver: silvery crystals of silver appear.
• Copper + zinc sulfate -> no reaction, because copper is less reactive than zinc.
• Aluminium looks unreactive only because of the protective oxide layer on its surface.

Corrosion of iron: rusting and how to stop it

Rusting is the corrosion of iron and steel. It needs both oxygen and water to happen - take away either one and the iron will not rust. The product is hydrated iron(III) oxide, Fe2O3.xH2O, the familiar orange-brown flaky solid. Salt speeds rusting up, which is why cars and railings near the coast or on salted winter roads corrode so quickly.

Rust is a poor protector. It is flaky and crumbly, so it falls away and exposes fresh iron underneath, and the corrosion eats into the metal. This is completely different from the oxide layer on aluminium, which is thin, tough and firmly attached, so it seals the metal and stops the reaction.

The simplest way to prevent rusting is a barrier method: cover the iron so that water and oxygen cannot reach it. Common barriers include painting (bridges, gates and car bodies), greasing or oiling (tools, machine parts and bicycle chains) and coating with plastic (dishwasher racks and garden furniture). Electroplating with chromium or nickel also works as a barrier.

Galvanising means coating iron or steel with a layer of zinc. It works in two ways. The zinc layer is a barrier, keeping out water and oxygen. But zinc is also more reactive than iron, so if the coating is scratched and the iron is exposed, the zinc corrodes in preference to the iron. This is sacrificial protection: the zinc loses electrons and is oxidised instead of the iron, so the iron is protected even where the coating has been broken. The same idea is used with blocks of magnesium bolted to the hulls of ships and to underground steel pipelines.

• Rusting needs oxygen and water together; salt makes it faster.
• Rust is hydrated iron(III) oxide, Fe2O3.xH2O.
• Barrier methods: painting, greasing, oiling, coating with plastic, electroplating, galvanising.
• Sacrificial protection: a more reactive metal corrodes instead of the iron, because it loses electrons more readily.

Extraction of metals: heating with carbon or electrolysis

How a metal is extracted from its ore depends on where it sits in the reactivity series. The more reactive the metal, the more strongly it is held in its compound, and the harder it is to get it out. Metals above carbon in the series - potassium, sodium, calcium, magnesium and aluminium - are extracted by electrolysis of their molten compounds, because no ordinary reducing agent is strong enough. Metals below carbon - zinc, iron and copper - are extracted by reduction with carbon, which is much cheaper. Copper, silver and gold are so unreactive that they are sometimes found in the ground as the metal itself.

Iron from hematite. The blast furnace is charged with iron ore (hematite, which is iron(III) oxide, Fe2O3), coke (carbon) and limestone (calcium carbonate). Coke burns in the hot air blast to give carbon dioxide and release a great deal of heat: C + O2 -> CO2. Higher up the furnace the carbon dioxide is reduced by more coke to carbon monoxide: C + CO2 -> 2CO. The carbon monoxide then reduces the iron(III) oxide to molten iron: Fe2O3 + 3CO -> 2Fe + 3CO2. The limestone removes the sandy impurity by thermal decomposition, CaCO3 -> CaO + CO2, and the calcium oxide produced reacts with silicon dioxide to form molten slag: CaO + SiO2 -> CaSiO3. The molten iron runs to the bottom of the furnace and the slag floats on top of it.

Aluminium from bauxite. The main ore of aluminium is bauxite, which is purified to give aluminium oxide. Aluminium oxide melts at over 2000 °C, which would be far too expensive, so it is dissolved in molten cryolite. Cryolite lowers the melting point to about 950 °C, so the electrolysis can run at a much lower temperature and use far less energy. The electrolyte is electrolysed between carbon electrodes: at the cathode Al3+ + 3e- -> Al, and at the anode 2O2- -> O2 + 4e-. The carbon anodes have to be replaced regularly because the oxygen produced at the anode reacts with the hot carbon and burns it away as carbon dioxide.

Notice the pattern: the method of extraction is chosen by the position of the metal in the reactivity series, and that is what the examiner is testing.