Topic 9: Metals

Metals dominate the left and centre of the Periodic Table and share a family of useful physical and chemical properties. Understanding how reactive each metal is lets us predict its reactions and choose the right method to extract it from its ore. This topic links physical structure, chemical behaviour, and the real-world challenges of corrosion and metal use.

Physical Properties of Metals

Metals share a recognisable set of physical properties because of their structure: a lattice of positive ions surrounded by a 'sea' of delocalised electrons. They are good conductors of heat and electricity because those free electrons move and carry energy and charge. Most metals are shiny when freshly cut, malleable (can be hammered into shape) and ductile (can be drawn into wires), since layers of ions slide over one another without breaking the bonding. They generally have high melting and boiling points and high density, and many are sonorous (ring when struck). These properties contrast sharply with most non-metals, which tend to be dull, brittle, and poor conductors.

Alloys and Why They Are Harder

An alloy is a mixture of a metal with one or more other elements, usually other metals. Common examples include brass (copper and zinc), bronze (copper and tin), and steel (iron with carbon). Pure metals are often too soft for many uses because their identical atoms are arranged in neat layers that slide over each other easily. In an alloy the added atoms are a different size, so they distort the regular layers and stop them sliding so readily. This makes alloys harder and stronger than the pure metals they are made from, which is why steel is used in construction rather than pure iron.

The Reactivity Series

The reactivity series ranks metals in order of how readily they react, from most reactive to least. A common order is: potassium, sodium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold. Carbon and hydrogen are non-metals included as reference points. The more reactive a metal, the more strongly it tends to lose electrons and form positive ions. This single list is extremely powerful: it predicts how vigorously a metal reacts with water and acids, which metals displace others from solution, and which method must be used to extract a metal from its ore.

Reactions With Water, Steam, Acids, and Oxygen

A metal's reactions follow its place in the reactivity series. Very reactive metals such as potassium, sodium, and calcium react with cold water to give a metal hydroxide and hydrogen gas. Less reactive metals like magnesium, zinc, and iron do not react much with cold water but will react with steam to give the metal oxide and hydrogen. Metals above hydrogen in the series react with dilute acids to give a salt and hydrogen, the reaction becoming more vigorous the higher the metal. Copper, silver, and gold sit below hydrogen and do not react with dilute acids. Most metals also react with oxygen to form oxides, with reactive metals oxidising quickly and unreactive ones, like gold, resisting it.

Displacement Reactions

In a displacement reaction a more reactive metal takes the place of a less reactive metal in a compound, usually a salt in solution. For example, iron added to copper(II) sulfate solution produces iron(II) sulfate and copper metal, because iron is more reactive than copper. The blue colour of the copper sulfate fades and a coating of copper forms on the iron. The reverse does not happen: copper cannot displace iron. These reactions confirm the order of the reactivity series and are a useful experimental way to compare two metals, since only the more reactive one will displace the other.

Extraction of Metals

How a metal is extracted depends on its reactivity. Metals below carbon in the series, such as iron, zinc, and copper, can be reduced from their oxides by heating with carbon, because carbon is more reactive and pulls the oxygen away. Iron is extracted in a blast furnace, where iron ore (haematite), coke, and limestone are added; carbon monoxide formed from the coke reduces the iron oxide to molten iron. Metals more reactive than carbon, such as aluminium, cannot be reduced this way and must be extracted by electrolysis of the molten compound. Aluminium is obtained by electrolysing molten aluminium oxide dissolved in cryolite. Very unreactive metals like gold are found native and need no chemical extraction.

Uses of Metals

Metals are chosen for jobs that match their properties. Aluminium is low density and resists corrosion (a protective oxide layer forms), so it is used in aircraft, drinks cans, and overhead power cables. Copper is an excellent electrical conductor and is ductile, making it ideal for electrical wiring, and its good heat conduction suits it to cooking pans and pipes. Iron and its alloy steel are strong and used for buildings, bridges, and vehicles. Mild steel is used for car bodies, while stainless steel (containing chromium) resists rusting and is used for cutlery. Matching property to use is a key exam skill.

Rusting of Iron and Its Prevention

Rusting is the corrosion of iron, forming hydrated iron(III) oxide. It needs both water and oxygen present; experiments using boiled water and drying agents show iron does not rust if either is removed, and salty water speeds it up. Rust is flaky and does not protect the metal beneath, so corrosion continues. Prevention works by keeping water and oxygen away, for example by painting, oiling, greasing, or coating with plastic. Galvanising coats iron with zinc, which forms a barrier and also gives sacrificial protection: because zinc is more reactive, it corrodes in preference to the iron even if the coating is scratched. Attaching blocks of a more reactive metal such as zinc or magnesium is another form of sacrificial protection used on ships and pipelines.