Giant Covalent Structures Exam Questions

• In diamond, each carbon atom is bonded to 4 other carbon atoms by strong covalent bonds in a tetrahedral arrangement [1]
• This forms a rigid three-dimensional giant covalent structure with no weak points [1]
• In graphite, each carbon atom is bonded to 3 other carbon atoms in flat layers/sheets [1]
• The layers are held together by weak intermolecular forces, so they can slide over each other easily [1]

Examiner note: Do not say “graphite has weak bonds.” The covalent bonds within each layer of graphite are strong. It is the forces between layers that are weak. This distinction is critical.

• In graphite, each carbon atom forms only 3 covalent bonds [1]
• This leaves one free/delocalised electron per carbon atom [1]
• These delocalised electrons can move along the layers, carrying charge / conducting electricity. Diamond has no delocalised electrons because all four outer electrons are used in bonding [1]

Examiner note: “Graphite has free electrons” is the key phrase. Each carbon has 4 outer electrons but uses only 3 for bonding — the fourth is delocalised. Diamond uses all 4 in bonding, so none are free.

• Silicon dioxide has a giant covalent structure [1]
• Each silicon atom is bonded to four oxygen atoms and each oxygen to two silicon atoms by strong covalent bonds [1]
• A large amount of energy is needed to break the many strong covalent bonds throughout the structure [1]

Examiner note: SiO2 has the same type of giant covalent structure as diamond. Do not say “it has strong intermolecular forces” — there are no molecules in a giant covalent structure. The bonds broken on melting are the covalent bonds themselves.

(a) A single layer of carbon atoms arranged in hexagons [1], with each carbon bonded to three others by covalent bonds, with one delocalised electron per atom [1]

(b) It conducts electricity (because of delocalised electrons) [1]

OR: It is very strong / has a very high tensile strength [1]

Examiner note: Graphene is essentially one sheet from graphite. It has all the in-plane properties of graphite (conductivity, strength) but is only one atom thick, making it useful in nanotechnology and flexible electronics.

(a) Diamond: giant covalent [1]; Methane: simple molecular [1]

(b) Diamond has strong covalent bonds throughout the entire structure in 3D [1]. Although NaCl has strong ionic bonds, the covalent bonds in diamond are even stronger / diamond has more bonds per atom to break [1]

Examiner note: Both are giant structures with strong bonds, but diamond’s tetrahedral network of covalent bonds (each atom bonded to 4 others) is stronger than the ionic bonds in NaCl. Do not say NaCl has “weak bonds.”

(a) Spherical / ball-shaped / cage structure, made of hexagons and pentagons of carbon atoms [1]

(b) C60 is a simple molecule / has a finite number of atoms, whereas diamond is a giant covalent structure that extends in all directions / has millions of atoms bonded together [1]

(c) C60 is a simple molecular substance — only weak intermolecular forces between C60 molecules need to be overcome to melt it, which requires little energy [1]

Examiner note: Although C60 contains strong covalent bonds, it forms discrete molecules. It is therefore a simple molecular substance (albeit a large one) and has the typical low melting point of simple molecular substances.

Melting point:

• All three have very high melting points because they are giant covalent structures with many strong covalent bonds that require a lot of energy to break [1]
• Diamond has the highest melting point because of its tetrahedral structure with 4 bonds per atom [1]

Hardness:

• Diamond is the hardest because every carbon is bonded to 4 others in a rigid 3D network with no planes of weakness [1]
• Graphite is soft because the layers can slide over each other due to weak forces between layers [1]

Electrical conductivity:

• Diamond and SiO2 do not conduct electricity because all outer electrons are used in covalent bonding and none are free to move [1]
• Graphite conducts electricity because each carbon forms only 3 bonds, leaving one delocalised electron per atom that can carry charge along the layers [1]

Examiner note: This is a 6-mark extended response. Organise your answer by property (melting point, then hardness, then conductivity), covering all three substances for each. Writing about one substance at a time makes it easy to miss comparisons.