Covalent Bonding
Principles of Chemistry · 5 question types
Exam Frequency Analysis
Past paper frequency (2018 to 2024)
This topic accounts for approximately 8% of your exam marks.
Dot-and-cross diagrams and comparison of giant covalent vs simple molecular structures appear regularly.
What they are
- A (also called a macromolecule) contains a vast number of non-metal atoms joined by strong covalent bonds in one continuous lattice
- Two classic examples, both built from carbon only:
- Diamond and graphite are of carbon: different forms of the same element with different bonding arrangements
Why all giant covalent structures have high melting points
- There are many strong running throughout the lattice
- Melting or vaporising the solid requires breaking very large numbers of these bonds
- That takes a great deal of thermal energy, so the melting point is high
Diamond
- Each carbon atom forms 4 single covalent bonds to four other carbon atoms
- The result is a rigid 3D tetrahedral network extending throughout the crystal
- Every outer electron on every carbon is locked into a covalent bond → no free electrons
Properties of diamond:
- Very hard: the rigid 3D network of strong covalent bonds resists deformation. Diamond is used in cutting tools, drill bits and saw edges where extreme hardness matters
- High melting point: the many strong covalent bonds running through the lattice need a large amount of energy to break
- Does not conduct electricity: every outer electron is locked in a covalent bond, so there are no free electrons or ions to move and carry charge
Explain why diamond has a high melting point (3 marks)
Write (three marks): (1) diamond has a giant covalent structure — a continuous 3D lattice; (2) it contains many strong covalent bonds throughout the lattice that must be broken to melt it; (3) breaking so many strong bonds requires a very large amount of energy.
Watch out: never refer to intermolecular forces or ions in diamond. The mark scheme awards no marks for M2 or M3 if you mention intermolecular forces — diamond has no intermolecular forces, only covalent bonds.
Graphite
- Each carbon atom forms only 3 single with three other carbon atoms
- These bonded atoms arrange themselves into flat layers of hexagons (sometimes called sheets)
- Each carbon has 4 outer electrons in total; 3 are used in bonds and the 4th is a that can move freely along the layer
- The hexagonal layers stack on top of each other, held together only by weak between layers, not by covalent bonds
Properties of :
- Soft and slippery: the weak forces between layers let the layers slide easily over each other. Graphite is used as a dry lubricant and as the writing "lead" in pencils
- Conducts electricity (and heat): the delocalised electrons can move along the layers and carry charge. Graphite is one of the very few non-metals that conducts electricity well
- High melting point: within each layer the covalent bonds are strong, and a great deal of energy is needed to break enough of them for the structure to fall apart
Why diamond doesn't conduct but graphite does
- An electric current is a flow of mobile charged particles
- In , every carbon uses all 4 of its outer electrons in covalent bonds. There are no free electrons available to move, so diamond is an electrical insulator
- In graphite, every carbon uses only 3 of its 4 outer electrons in bonds. The 4th is delocalised and can move freely along the layer, carrying charge along with it. Graphite therefore conducts electricity
Explain why graphite conducts electricity (2 marks)
Write (two marks): (1) graphite has delocalised electrons; (2) those electrons are free to move throughout the structure.
Watch out: M2 depends on M1 — you must name electrons before saying they move. Saying "ions move" or "atoms move" scores zero for the whole question. "Carries charge" on its own is not enough for M2; the mark scheme requires you to say the electrons can move or flow.
Why diamond is hard but graphite is soft
What comes up: explain the difference in hardness between diamond and graphite.
Watch out: graphite's softness comes from layers sliding over each other because the forces between the layers are weak — not from "weak bonds". The covalent bonding within each layer (and throughout diamond's 3D lattice) stays strong.