The four crystal structures and their properties

Four ways to build a solid

How a solid behaves when it is heated depends on what it is made of and what holds those particles together. Salt, magnesium ribbon and the graphite in a pencil show no change when warmed in a test tube over a Bunsen flame. Grey-black iodine crystals turn within a minute into a purple vapour, which grows dark crystals again on the cool top of the tube.

There are four types of crystal structure. In an ionic crystal the particles are oppositely charged ions. In a metallic crystal they are positive ions in a sea of delocalised electrons. In a macromolecular crystal they are atoms joined by covalent bonds into one network running through the whole crystal. In a molecular crystal they are small molecules, each held together internally by covalent bonds but held to its neighbours only by intermolecular forces.

The first three are giant structures, with the strong bonding running unbroken from one side of the crystal to the other. To melt them, that bonding has to be overcome. A molecular crystal is different: melting separates the molecules from each other, and the covalent bonds inside each molecule stay intact. Intermolecular forces are far weaker than covalent bonds, so salt melts at 801 °C and iodine at 114 °C. Name the force that is overcome: 'the bonds break' for iodine does not earn the mark, because the covalent bonds are not the ones that break.

Giant structure

A structure in which the bonding extends throughout the whole crystal, with no separate molecules, as in ionic, metallic and macromolecular crystals.

Macromolecular crystal

A giant structure of atoms joined by covalent bonds in a network that extends throughout the crystal; also called giant covalent.

Molecular crystal

A regular arrangement of simple molecules held in place by intermolecular forces, with covalent bonds only within each molecule.

Intermolecular force

An attraction between separate molecules, much weaker than the covalent bonds within them.

Macromolecular: diamond, graphite and silicon dioxide

Diamond is pure carbon in which every atom is covalently bonded to four others at the corners of a tetrahedron, with bond angles of 109.5°. The network continues in three dimensions to the edges of the crystal. Melting it would mean breaking covalent bonds, four round every atom, and at atmospheric pressure diamond does not melt at all: heated in the absence of air it turns to graphite, and graphite only sublimes, above about 3600 °C. The same rigid network makes it the hardest natural substance, used to tip saws and drills for cutting stone.

All four outer electrons of every carbon atom in diamond are held in bonds between two particular atoms. None is free to move, so diamond does not conduct electricity.

Graphite is pure carbon too, arranged differently. Each atom is bonded to three others in a flat sheet of hexagons, with bond angles of 120°. The fourth outer electron of each atom is delocalised across the sheet, above and below the plane of the atoms. The sheets are stacked 0.335 nm apart, more than twice the C–C bond length of 0.142 nm within a sheet, and they are held to each other only by weak forces.

Describe the drawing

Two ball-and-stick structures of carbon. On the left, diamond: fourteen grey atoms in a three-dimensional network, one atom and its four neighbours picked out in cyan with the tetrahedral angle of 109.5° marked. On the right, graphite: three flat layers of fused hexagons seen at an angle, each atom bonded to three others at 120°, alternate layers offset, dashed amber lines between layers marking the weak forces and a spacing of 0.335 nm, an in-layer bond of 0.142 nm, and a faint cyan band above and below the top layer for its delocalised electrons.

That structure gives graphite its mixed properties. The covalent bonds within the layers are strong, so it does not melt below about 3600 °C. The forces between the layers are weak, so the layers slide over each other: graphite is soft, marks paper and is used as a dry lubricant. The delocalised electrons move freely along the layers, so graphite conducts electricity in the direction of the layers and is used for electrodes. It conducts poorly at right angles to them.

The wide spacing between the layers also makes graphite less dense than diamond: 2.26 g cm⁻³ against 3.51 g cm⁻³. Diamond and graphite are allotropes of carbon. A single layer of graphite is graphene, one atom thick, very strong in its own plane and a good conductor, since its delocalised electrons move across the sheet.

Silicon dioxide, the mineral quartz and the main part of sand, is macromolecular as well. Each silicon atom is bonded to four oxygen atoms tetrahedrally and each oxygen atom to two silicon atoms, so the formula SiO₂ is a ratio, as it is for an ionic compound. It melts at about 1710 °C, is hard, does not conduct and does not dissolve in water. Silicon itself has the diamond structure.

Allotrope

One of two or more different structural forms of the same element in the same physical state, such as diamond and graphite.

Molecular: iodine and ice

Iodine is made of I₂ molecules. Inside each molecule the two iodine atoms share a pair of electrons in a covalent bond. Between molecules there are only induced dipole–dipole forces, the weakest kind of intermolecular force, though with 106 electrons in every molecule they are strong enough to keep iodine solid at room temperature. The molecules pack in a regular lattice, so the solid is crystalline.

Warming iodine gently overcomes the forces between the molecules and the solid sublimes to a purple vapour. The I–I bonds are not affected: the vapour is still I₂, and it condenses back to the same crystals.

Describe the drawing

A thin rectangular box drawn in perspective with a cyan dumbbell-shaped I₂ molecule, two touching spheres, at each of its eight corners and the centre of each of its six faces. The molecules at the corners and on the top and bottom faces tilt one way; those on the four side faces tilt the other. A dotted amber line between two neighbouring molecules marks the induced dipole–dipole forces, and a magnified callout of one molecule marks the covalent bond between its two iodine atoms.

Molecular solids share a set of properties. Their melting and boiling points are low, because only weak intermolecular forces are overcome. They do not conduct electricity as solids or liquids, because the molecules are uncharged and there are no delocalised electrons. They are usually soft. Their solubility depends on the forces they can form with a solvent: iodine dissolves poorly in water but readily in hexane, giving a violet solution.

Ice is a molecular crystal too, but the forces between its molecules are hydrogen bonds, the strongest kind of intermolecular force. In ice every water molecule is hydrogen bonded to four others, two through its own hydrogen atoms and two through its lone pairs. The four neighbours sit at the corners of a tetrahedron, and the result is an open lattice of puckered hexagonal rings with a lot of empty space inside them.

Describe the drawing

Two panels of twelve water molecules each, cyan oxygen atoms with white hydrogens. Ice, on the left: the molecules form a hexagonal ring round an empty central channel, with more rings starting outside it; every oxygen has two covalent O–H bonds and two dashed amber hydrogen bonds, one connection in each panel pointing out of the page as a wedge or a hatched wedge. Liquid water, on the right: the same twelve molecules packed into a smaller area at random angles, some where the channel was, with only a few hydrogen bonds.

When ice melts, some of the hydrogen bonds break, the open lattice collapses and molecules move into the spaces. Liquid water is therefore denser than ice, 1.00 g cm⁻³ against 0.917 g cm⁻³ at 0 °C. Ice floats, and a pond freezes from the top down, the ice insulating the water below it. For almost every other substance the solid is denser than the liquid and sinks in it.

Sublimation

The change of state from solid directly to gas, without passing through the liquid.

Energy at a change of state

Melting and boiling are endothermic for every substance, because energy is needed to overcome the forces between the particles. A beaker of crushed ice heated steadily stays at 0 °C until the last of the ice has gone; the temperature then rises, and stops again at 100 °C while the water boils. At each pause the energy going in does not raise the temperature. Freezing and condensing release the same amounts of energy, which is why steam at 100 °C scalds worse than water at 100 °C.

The energy needed per mole is the enthalpy of fusion for melting and the enthalpy of vaporisation for boiling. Both follow the strength of the forces that are overcome, and vaporisation is always the larger. On melting the particles stay in contact and most of the attractions survive, loosened; on boiling every particle is separated from all its neighbours. For water the two values are 6.01 and 40.7 kJ mol⁻¹. Sodium chloride, a giant structure, needs 28 kJ mol⁻¹ just to melt, more than four times water's enthalpy of fusion.

Enthalpy of fusion

The enthalpy change when one mole of a solid is converted into a liquid at its melting point.

Enthalpy of vaporisation

The enthalpy change when one mole of a liquid is converted into a gas at its boiling point.

Comparing the four structures

Side by side, the four structures differ first in their particles and in the force that has to be overcome to separate them. The melting point and the conductivity follow from those two.

The energy needed to melt follows the force. Ionic and macromolecular crystals need a great deal, metals usually a lot, molecular crystals little. Metals have a feature of their own: melting only loosens the lattice, since the ions in the liquid are still held by the delocalised electrons, so a metal's boiling point lies far above its melting point. Magnesium melts at 650 °C and boils at 1090 °C.

To identify an unknown solid from data, start with conductivity. A solid that conducts is metallic, or graphite. One that conducts only when molten or dissolved is ionic. If it conducts in neither state, the melting point separates macromolecular, which is very high, from molecular, which is low.

A 'high' melting point covers a wide range, and the ranges for the giant structures overlap. Magnesium melts at 650 °C, below sodium chloride, but tungsten, also metallic, melts at 3422 °C, far above it. Strength depends on the charges and sizes of the particles in each substance, so two substances are compared one against the other, never by the name of the bonding type alone. The gap between any giant structure and a small molecular one, by contrast, is always large.

AQA asks for each structure to be drawn with a stated number of particles, and the drawing must show the bonding, not only the arrangement. For sodium chloride, draw a cube of eight ions with Na⁺ and Cl⁻ alternating at the corners and a key or a charge on each. For a metal, draw a regular arrangement of ions, each with its charge, with electrons in the spaces labelled delocalised. For diamond, one carbon bonded to four others; for graphite, at least two layers of hexagons with the gap between them labelled. For iodine or ice, separate molecules, with the forces between them drawn differently from the bonds within them.

Other boards: graphene, fullerenes and 'giant covalent'

OCR, Edexcel and Cambridge say giant covalent where AQA says macromolecular; the two terms describe the same structure, and a scheme on either board accepts its own word. OCR and Edexcel also name graphene, the single layer of graphite, and ask why it is strong and conducts.

Cambridge adds a third allotrope of carbon, which is molecular. Buckminsterfullerene, C₆₀, is a closed cage of sixty carbon atoms arranged in twelve pentagons and twenty hexagons, the pattern on a football. Each carbon is bonded to three others, as in graphite, and the remaining electrons are delocalised over the cage. Between one C₆₀ molecule and the next there are only induced dipole–dipole forces: strong for a molecule with 360 electrons, but nothing like the covalent network of diamond.

Solid C₆₀ is therefore soft, turns to vapour at a far lower temperature than diamond or graphite, and dissolves in non-polar solvents such as methylbenzene to give a purple solution. It is a very poor conductor: its delocalised electrons can move round one cage but cannot cross the gap to the next.

Fullerene

A molecular allotrope of carbon in which each atom is bonded to three others in a closed cage or tube, such as C₆₀.