Metals

9.1 Properties of metals

Chemical properties

Metals form positive ions by losing electrons. The general reactions are:

How vigorously a metal reacts depends on its position in the reactivity series (section 9.4). Metals low in the series, such as copper, react with none of these under normal conditions.

9.2 Uses of metals

The use of a metal depends on its physical properties.

9.3 Alloys and their properties

Alloys are harder and stronger than the pure metals. In a pure metal the atoms are all the same size, in regular layers, and the layers can slide over each other easily, so the metal is soft and bends. In an alloy, the atoms (or ions) of the other element are a different size, which disrupts the regular layers, so the layers can no longer slide over each other as easily.

Stainless steel is used for cutlery because it is hard and resists rusting.

9.4 Reactivity series

The reactivity series lists metals in order of their tendency to form positive ions (to lose electrons).

A way to remember the order: Please Stop Crying My Aunt Carol Zebra Is Hungry Cats Should Go. (Carbon and hydrogen are included for comparison; they are not metals in this sense.)

Reactions with water and steam. Potassium, sodium and calcium react with cold water to give the hydroxide and hydrogen. Magnesium reacts with steam to give magnesium oxide and hydrogen. The unreactive metals below do not react.

$\ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)}$

Reactions with dilute hydrochloric acid. Magnesium, zinc and iron react to give a salt and hydrogen (magnesium most vigorously), but copper, silver and gold do not react, because they are below hydrogen in the series.

Displacement reactions

The more reactive metal forms its positive ion (loses electrons) and the less reactive metal ion is reduced to the metal.

$\ce{Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s)}$

$\ce{Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)}$

The blue colour of copper(II) sulfate solution fades when a more reactive metal displaces copper, and a pink-brown solid of copper forms.

$\text{X}$ is more reactive than copper and less reactive than zinc, so it is between zinc and copper. (Iron is an example.)

The more bubbles, the more reactive the metal, so the order is $\text{Q}$ (most), then $\text{P}$, then $\text{R}$ (least).

The apparent unreactivity of aluminium

Aluminium is high in the reactivity series, but it seems unreactive. This is because it is covered with a thin, strong layer of aluminium oxide that stops the acid or water reaching the metal. If the layer is removed, aluminium reacts vigorously.

9.5 Corrosion of metals

Rusting is the corrosion of iron and steel.

The experiment shows both are needed: a nail in air and water rusts; a nail in boiled water (no dissolved air) covered with oil does not rust, and a nail in dry air (with calcium chloride to absorb water vapour) does not rust.

Preventing rust

Barrier methods work by excluding oxygen or water:

• painting (cars, bridges, ships),
• greasing or oiling (machine parts, bicycle chains),
• coating with plastic (fridge shelves, wire fences),
• coating with another metal such as tin or chromium.

Sacrificial protection. A more reactive metal such as zinc or magnesium is attached to the iron. The more reactive metal loses electrons in preference to the iron, so it corrodes instead of the iron:

$\ce{Zn -> Zn^2+ + 2e-}$

The electrons flow to the iron, which keeps it from forming $\ce{Fe^2+}$ ions. The zinc or magnesium is used up and must be replaced. Blocks of zinc or magnesium are bolted to the hulls of ships and to underground pipes.

Galvanising is coating iron with zinc. It protects by both methods: the zinc layer is a barrier, and if it is scratched, the zinc still protects the iron by sacrificial protection, because zinc is more reactive than iron.

9.6 Extraction of metals

A metal is found in the Earth as an ore, a rock that contains the metal compound in a large enough amount to be worth extracting. The method depends on the metal's position in the reactivity series.

The more reactive the metal, the more difficult it is to extract, as its compounds are more stable.

Iron: the blast furnace

The raw materials are iron ore (hematite, $\ce{Fe2O3}$), coke (carbon), limestone ($\ce{CaCO3}$) and hot air.

1. Coke burns in the hot air, giving carbon dioxide and heat energy:

$\ce{C(s) + O2(g) -> CO2(g)}$

2. The carbon dioxide is reduced to carbon monoxide by more hot coke:

$\ce{CO2(g) + C(s) -> 2CO(g)}$

3. Carbon monoxide reduces iron(III) oxide to iron. The iron(III) oxide loses oxygen, so it is reduced:

$\ce{Fe2O3(s) + 3CO(g) -> 2Fe(l) + 3CO2(g)}$

4. Limestone removes impurities. It decomposes with heat (thermal decomposition) to give calcium oxide:

$\ce{CaCO3(s) -> CaO(s) + CO2(g)}$

5. The basic calcium oxide reacts with acidic impurity silicon(IV) oxide, $\ce{SiO2}$, to form slag, calcium silicate:

$\ce{CaO(s) + SiO2(s) -> CaSiO3(l)}$

The molten iron is denser, so it collects at the bottom, and the slag floats on top of it. Both are tapped off separately. The slag is used for road building.

Aluminium: electrolysis

Aluminium is extracted from aluminium oxide, $\ce{Al2O3}$, purified from the ore bauxite. It is more reactive than carbon, so it is made by electrolysis.

• Aluminium oxide has a very high melting point ($2050\,^\circ\text{C}$), so it is dissolved in molten cryolite. This lowers the temperature of the electrolyte to about $950\,^\circ\text{C}$ and saves energy.
• The electrodes are carbon (graphite).
• At the cathode (negative), aluminium ions are reduced: $\ce{Al^3+ + 3e- -> Al}$. Molten aluminium sinks to the bottom and is tapped off.
• At the anode (positive), oxide ions are oxidised: $\ce{2O^2- -> O2 + 4e-}$.

The oxygen formed at the anode reacts with the hot carbon anodes, so they burn away to carbon dioxide and must be replaced regularly:

$\ce{C(s) + O2(g) -> CO2(g)}$

Aluminium is more reactive than carbon, so carbon cannot take oxygen away from aluminium oxide. Electrolysis is needed.