Metallic bonding and types of crystal structureAQA A-Level Chemistry: Revision notes
Section 1
Metallic bonding
A metal is a giant lattice of positive metal ions surrounded by a 'sea' of delocalised electrons that came from the outer shells of the atoms. Metallic bonding is the electrostatic attraction between the positive ions and the delocalised electrons.
The strength increases with (i) a higher ionic charge, (ii) more delocalised electrons per atom and (iii) a smaller ionic radius. So Mg (2+, two electrons each) has a higher melting point (650 °C) than Na (1+, 98 °C).
Metals conduct because the delocalised electrons are free to move through the lattice. They are malleable and ductile because layers of ions slide over one another while the electrons keep the bonding intact.
In 'explain' answers say that the attraction is between the positive ions and the delocalised electrons. Never say that ions carry the current in a solid metal.
Section 2
The four types of crystal structure
Solids are classified by the particles in the lattice and the forces holding them together:
- Ionic (e.g. sodium chloride): oppositely charged ions, strong electrostatic attraction
- Metallic (e.g. magnesium, sodium): positive ions and delocalised electrons
- Macromolecular / giant covalent (e.g. diamond, graphite): atoms joined by a network of covalent bonds
- Simple molecular (e.g. iodine, ice): small molecules with covalent bonds inside and weak intermolecular forces between them
Section 3
Ionic and metallic structures: sodium chloride, magnesium and sodium
Sodium chloride is a giant ionic lattice in which each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺. Strong electrostatic attraction acts in all directions, so the melting point is high (801 °C). It is brittle because displacing a layer brings like charges together and the crystal splits.
It does not conduct as a solid because the ions are fixed, but conducts when molten or dissolved because the ions are free to move.
Magnesium and sodium are metallic lattices (see above): they conduct in both solid and liquid states.
Section 4
Macromolecular structures: diamond and graphite
Diamond: each carbon forms four covalent bonds in a tetrahedral arrangement (bond angle 109.5°). The strong bonds in all directions give a very high melting point and extreme hardness. There are no mobile charge carriers, so it does not conduct.
Graphite: each carbon forms three covalent bonds in flat layers of hexagonal rings (bond angle 120°). The fourth outer electron is delocalised, so graphite conducts along the layers. Weak intermolecular forces between layers let them slide, so graphite is soft and slippery. Its melting point is still very high (above 3500 °C) because the covalent bonds within the layers must be broken to melt it.
Do not say that graphite has a low melting point because its layers are weakly held. Melting breaks covalent bonds within the layers, so it is very high.
Section 5
Simple molecular structures: iodine and ice
Iodine is a lattice of I₂ molecules with covalent bonds inside each molecule and weak induced dipole-dipole forces between molecules. It has a low melting point (114 °C) and does not conduct, because there are no ions or delocalised electrons. When it melts or sublimes, only the intermolecular forces are overcome.
Ice is an open lattice of H₂O molecules held by hydrogen bonds. Each oxygen is bonded to two hydrogens by covalent bonds and to two other molecules by hydrogen bonds. The open structure makes ice less dense than liquid water.
When a simple molecular substance melts or boils, covalent bonds do not break. Never write that they do.
Section 6
Changes of state and diagrams
Melting and boiling need energy to overcome the forces holding particles together. During a change of state the temperature stays constant: the energy supplied increases the potential energy of the particles, not their kinetic energy.
The energy needed depends on the type of bonding: small for simple molecular (weak intermolecular forces), large for ionic, metallic and macromolecular (strong bonds). Boiling needs much more energy than melting because the particles must be separated completely.
When drawing structures, show enough particles to show the pattern: for NaCl alternate Na⁺ and Cl⁻ in a square grid (at least 3 × 3); for a metal a regular array of positive ions with electrons shown between them; for graphite at least three joined hexagons in a layer.
Section 7
Linking properties to structure
A standard exam method is to name the structure, name the particles, state the forces, then link them to the property.
- High melting point: strong bonds or forces in a giant structure (ionic, metallic, macromolecular)
- Low melting point: weak intermolecular forces (simple molecular)
- Conducts when solid: delocalised electrons (metals, graphite)
- Conducts only when molten or dissolved: mobile ions (ionic)
- Does not conduct: no mobile charged particles (diamond, iodine, ice)
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Metallic bonding and types of crystal structure
- A technician is testing four unlabelled solids at room temperature. Solid W conducts electricity, is malleable and melts at 650 °C. Solid X does not conduct as a solid, conducts when molten and melts at 801 °C. Solid Y does not conduct in any state and melts at 114 °C. Solid Z is very hard, does not conduct and melts above 3500 °C.Explain why solid X conducts electricity when molten but not when solid.2 marks
- A student compares sodium, which melts at 98 °C, with magnesium, which melts at 650 °C. Both metals conduct electricity as solids and both can be hammered into shape.Explain why sodium can be hammered into shape without breaking.2 marks
- Graphite and diamond are both forms of carbon. Graphite is used as electrodes and as the 'lead' in pencils, which leave a mark because layers rub off onto paper. Diamond is used on the cutting edges of drill bits.Explain why graphite conducts electricity but diamond does not.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).