What you'll learn
Metallic bonding and the properties of metals completes the set of three bonding types in AQA GCSE Chemistry, and it is the most economical of the three to learn: a single structural picture explains every property metals possess. Metals consist of positive ions arranged in layers, surrounded by a sea of electrons that have left their parent atoms and are free to move throughout the whole structure. Conductivity, malleability, ductility, high melting points and the behaviour of alloys all follow from that one arrangement. By the end of this guide you should be able to describe metallic bonding precisely, explain each property of metals in terms of the structure, explain why alloys are harder than pure metals, compare metals with ionic and covalent substances, and identify a metal from experimental data.
Key terms and definitions
Metallic bond — the strong electrostatic attraction between positive metal ions and the delocalised electrons surrounding them
Delocalised electron — an electron no longer associated with a particular atom and free to move throughout the structure
Giant metallic structure — the regular arrangement of layers of positive metal ions in a sea of delocalised electrons
Malleable — able to be hammered or rolled into shape without breaking
Ductile — able to be drawn out into a wire
Alloy — a mixture of a metal with one or more other elements
Lattice — a regular repeating arrangement of particles
Thermal conductivity — the ability to transfer energy by heating
Pure metal — a metal consisting of atoms of one element only, arranged in regular layers
Core concepts
The structure of a metal
In a metal, the atoms are arranged in a giant structure of regular layers. The electrons in the outer shell of each atom are delocalised, meaning they are no longer attached to any particular atom and are free to move throughout the whole structure.
Because each atom has lost its outer electrons to this shared pool, what remains at each lattice point is a positive ion. The structure is therefore described as a lattice of positive ions in a sea of delocalised electrons.
The metallic bond is the strong electrostatic force of attraction between these positive ions and the delocalised electrons. Note carefully that it is not an attraction between ions, as in ionic bonding, nor a shared pair, as in covalent bonding. Stating it as the attraction between positive metal ions and delocalised electrons is what earns the mark.
The bond acts throughout the structure in all directions, which is why metals form giant structures rather than molecules.
High melting and boiling points
Metals generally have high melting and boiling points because the metallic bonds are strong, and a large amount of energy is needed to overcome the many strong electrostatic attractions between the positive ions and the delocalised electrons.
The strength of the bond varies with the charge on the ion and the number of delocalised electrons contributed. A metal whose atoms release two or three electrons produces ions of higher charge and a denser electron sea, giving stronger attraction and a higher melting point. This is why magnesium melts at a far higher temperature than sodium, and it is a point examiners use to distinguish strong candidates.
Mercury is the familiar exception, being liquid at room temperature, and the alkali metals have relatively low melting points because each atom contributes only one electron.
Electrical conductivity
Metals conduct electricity because the delocalised electrons are free to move through the structure. When a potential difference is applied, these electrons drift through the metal, carrying charge from one end to the other.
Two points distinguish this from ionic conduction and are worth stating. First, metals conduct in the solid state, whereas ionic compounds do not, because in a metal the electrons move while the ions stay put. Second, the metal is not changed by conducting, whereas an ionic compound is decomposed by electrolysis.
Thermal conductivity
Metals are good conductors of thermal energy for the same reason. The delocalised electrons gain kinetic energy at the hot end of the metal, move rapidly through the structure, and transfer that energy to ions elsewhere by collision.
The vibration of the ions themselves also transfers some energy, but electron movement is the faster and more important mechanism, which is why good electrical conductors are also good thermal conductors.
Malleability and ductility
The layers of positive ions in a pure metal can slide over one another when a force is applied. As they slide, the delocalised electrons move with them, so the metallic bonding is maintained throughout and the structure holds together in its new shape rather than breaking.
This is why metals can be hammered or rolled into sheets, which is malleability, and drawn out into wires, which is ductility.
The contrast with ionic compounds is instructive and frequently examined. In an ionic lattice, sliding one layer brings ions of like charge alongside one another; they repel and the crystal shatters. In a metal, the electron sea is uniform, so no such repulsion arises and the layers simply move.
Alloys
Pure metals are often too soft for practical use precisely because their regular layers slide so easily. An alloy solves this.
An alloy is a mixture of a metal with one or more other elements, usually other metals or carbon. The added atoms are of a different size from the host atoms, so they distort the regular arrangement of the layers.
Because the layers are no longer regular, they cannot slide over one another easily, and the alloy is therefore harder than the pure metal.
Examples worth knowing include steel, which is iron alloyed with carbon and sometimes other metals; bronze, which is copper and tin; brass, which is copper and zinc; and gold alloys, in which pure gold is mixed with copper, silver or zinc to make jewellery hard enough to wear.
Note that an alloy is a mixture, not a compound. The proportions can be varied, which is exactly why alloys are so useful: the composition is adjusted to give the properties a particular application needs.
Identifying a metal from data
Questions often provide a table of properties and ask you to deduce structure and bonding. The diagnostic for a metal is conduction in the solid state.
A substance with a high melting point that conducts as a solid is metallic. A substance with a high melting point that conducts only when molten or dissolved is ionic. A substance with a high melting point that does not conduct at all is giant covalent, with graphite the exception. A substance with a low melting point that does not conduct is simple molecular.
Metals are also typically malleable and insoluble in water, both of which support the identification.
Worked examples
Example 1: Explaining conductivity (4 marks)
Explain why copper conducts electricity in the solid state but sodium chloride does not.
In copper, the outer electrons of the atoms are delocalised and free to move throughout the giant metallic structure. When a potential difference is applied these electrons move through the metal, carrying charge, so solid copper conducts.
In solid sodium chloride, the charged particles are ions held in fixed positions within a giant ionic lattice by strong electrostatic forces. Although the ions are charged, they cannot move, so no charge can flow and the solid does not conduct. Sodium chloride conducts only when melted or dissolved, once the ions are free to move.
Example 2: Explaining the hardness of an alloy (4 marks)
Explain why steel is harder than pure iron.
Pure iron consists of atoms of a single size arranged in regular layers. Because the layers are regular, they can slide over one another readily when a force is applied, which makes pure iron relatively soft.
Steel is an alloy of iron with carbon. The carbon atoms are a different size from the iron atoms, so they distort the regular arrangement of the layers. The distorted layers cannot slide over one another easily, so a greater force is needed to change the shape of the metal and steel is harder than pure iron.
Example 3: Comparing melting points (3 marks)
Sodium melts at 98 degrees Celsius while magnesium melts at 650 degrees Celsius. Suggest why.
Each sodium atom contributes one outer electron to the sea of delocalised electrons and becomes an ion with a charge of plus 1. Each magnesium atom contributes two electrons and becomes an ion with a charge of plus 2.
Magnesium therefore has more highly charged positive ions and a greater number of delocalised electrons, so the electrostatic attraction between the ions and the electron sea is stronger. More energy is needed to overcome this stronger metallic bonding, so magnesium has the higher melting point.
Common mistakes and how to avoid them
The most frequent error is describing a metal as containing atoms rather than positive ions. Once the outer electrons are delocalised, what remains at each lattice point is an ion, and the mark scheme is strict about this.
Students often write that electrons are free without using the word delocalised, or without saying that they move through the whole structure. Both elements are usually required.
Another common slip is explaining malleability by saying that metals are soft. The explanation is that the layers of ions slide over one another while the metallic bonding is maintained.
Many candidates say alloys are harder because they contain two metals. The reason is that the different-sized atoms distort the layers so they cannot slide.
Finally, candidates sometimes describe an alloy as a compound. It is a mixture, which is why its composition and properties can be adjusted.
Exam technique for "Metallic bonding and properties of metals"
Every property question here follows the same chain: describe the structure, name the feature responsible, then link it to the property. Delocalised electrons explain conduction; sliding layers explain malleability; strong attraction explains melting point.
When asked to compare a metal with an ionic compound, use conductivity in the solid state as the key difference and explain why in terms of what is free to move — electrons in the metal, nothing in the ionic solid.
For alloy questions, always mention the difference in atom size and the consequent distortion of the layers. Saying only that the layers cannot slide omits the reason.
In data questions, check solid-state conductivity first. It identifies a metal immediately and separates it from every other structure type.
Quick revision summary
A metal is a giant structure of positive ions arranged in layers within a sea of delocalised electrons, and the metallic bond is the strong electrostatic attraction between the two, acting throughout the structure. Melting and boiling points are high because many strong attractions must be overcome, and are higher for metals whose atoms release more electrons and form more highly charged ions. Metals conduct electricity in the solid state because the delocalised electrons move through the structure carrying charge, and they conduct thermal energy by the same mechanism. They are malleable and ductile because the layers of ions slide over one another while the metallic bonding is maintained, unlike ionic lattices where sliding brings like charges together and shatters the crystal. Alloys are mixtures in which differently sized atoms distort the regular layers so they cannot slide easily, making the alloy harder than the pure metal; steel, bronze and brass are the standard examples. Conduction in the solid state is the diagnostic that distinguishes a metal from ionic, giant covalent and simple molecular substances.