Crystalline solids and lattice structures

· Crystalline solid = particles arranged in a regular repeating lattice.
· The type of particles and forces/bonds between them determine physical properties.
· Main structure types: giant ionic, simple molecular, giant molecular / giant covalent, and giant metallic.
· Exam answers must link structure → bonding/forces → property, not just state the property.

Giant ionic lattices

· Examples: sodium chloride, NaCl and magnesium oxide, MgO.
· Structure: regular 3D lattice of oppositely charged ions.
· Bonding: strong electrostatic attractions between positive and negative ions in all directions.
· High melting and boiling points because many strong ionic bonds must be overcome.
· Do not conduct electricity when solid because ions are fixed in position.
· Conduct electricity when molten or aqueous because ions are mobile.
· Often soluble in water if attractions between ions and water molecules overcome ionic lattice attractions.
· MgO has a higher melting point than NaCl because Mg²⁺ and O²⁻ have higher ionic charges, giving stronger electrostatic attractions.

Sodium chloride has a giant ionic lattice made from alternating Na⁺ and Cl⁻ ions. Strong electrostatic attractions act in all directions, explaining its high melting point. Ions are fixed in the solid but mobile when molten or dissolved. Source

Simple molecular lattices

· Examples: iodine, I₂, buckminsterfullerene, C₆₀, and ice, H₂O.
· Structure: made of separate molecules arranged in a lattice.
· Inside each molecule: strong covalent bonds.
· Between molecules: weak intermolecular forces such as van der Waals’ forces or hydrogen bonding.
· Low melting and boiling points compared with giant structures because only intermolecular forces are overcome.
· Usually poor electrical conductors because there are no mobile ions or delocalised electrons.
· Iodine: simple molecular solid; melting involves overcoming instantaneous dipole–induced dipole forces, not breaking I–I covalent bonds.
· C₆₀: molecules are covalently bonded internally, but molecules are held together by weak van der Waals’ forces.
· Ice: water molecules are held by hydrogen bonds in an open structure, so ice is less dense than liquid water.

Ice is a simple molecular solid with water molecules arranged in an open hydrogen-bonded lattice. The open structure explains why ice is less dense than liquid water. Melting ice breaks hydrogen bonds between molecules, not O–H covalent bonds. Source

Giant molecular / giant covalent structures

· Examples: diamond, graphite, and silicon(IV) oxide, SiO₂.
· Structure: atoms joined by strong covalent bonds in a giant lattice.
· Very high melting points because many strong covalent bonds must be broken.
· Usually insoluble because covalent networks are too strong to separate into solvent particles.

Diamond and graphite are both giant covalent forms of carbon but have different structures. Diamond has a 3D tetrahedral network, while graphite has layers. Their different bonding explains their different hardness and electrical conductivity. Source

Diamond, graphite and silicon(IV) oxide

· Diamond: each carbon atom forms 4 covalent bonds in a tetrahedral 3D network.
· Diamond is very hard, has a very high melting point, and does not conduct electricity because it has no mobile charged particles.
· Graphite: each carbon atom forms 3 covalent bonds in hexagonal layers.
· Graphite has delocalised electrons between carbon atoms within layers, so it conducts electricity.
· Graphite is soft/slippery because layers are held together by weak van der Waals’ forces and can slide.
· Silicon(IV) oxide, SiO₂: each silicon is covalently bonded to oxygen atoms in a giant covalent network.
· SiO₂ has a high melting point and does not conduct electricity because there are no mobile ions or electrons.

Silicon(IV) oxide forms a giant covalent lattice of silicon and oxygen atoms. Many strong Si–O covalent bonds must be broken to melt it. This explains its high melting point and lack of electrical conductivity. Source

Giant metallic structures

· Example: copper, Cu.
· Structure: positive metal ions in a regular lattice surrounded by delocalised electrons.
· Bonding: strong electrostatic attraction between positive ions and delocalised electrons.
· High melting and boiling points because metallic bonding is strong.
· Good electrical conductivity because delocalised electrons can move and carry charge.
· Malleable and ductile because layers of metal ions can slide while metallic bonding is maintained.

Metallic bonding consists of positive metal ions attracted to a sea of delocalised electrons. The mobile electrons explain electrical conductivity. The non-directional bonding helps explain malleability and ductility. Source

Comparing physical properties

· Highest melting points usually: giant covalent and many giant ionic / metallic structures.
· Lowest melting points usually: simple molecular substances because only intermolecular forces are overcome.
· Electrical conductors when solid: metals and graphite.
· Electrical conductors when molten/aqueous: ionic compounds.
· Electrical non-conductors: simple molecular substances, diamond, and SiO₂.
· Soluble in water: many ionic compounds and some polar simple molecular substances.
· Insoluble in water: giant covalent structures, most metals, and many non-polar simple molecular substances.

Deducing bonding and structure from data

· High melting point + conducts only when molten/aqueous → likely giant ionic.
· Low melting/boiling point + does not conduct → likely simple molecular.
· Very high melting point + does not conduct → likely giant covalent, e.g. diamond or SiO₂.
· Very high melting point + conducts when solid → likely metallic or graphite.
· Soft/slippery solid + conducts electricity → likely graphite.
· Always explain conductivity using mobile charged particles: mobile ions or delocalised electrons.
· Always explain melting point using the strength and number of bonds/forces overcome.

Common exam traps

· Do not say ionic solids conduct electricity: solid ionic compounds do not conduct because ions are fixed.
· Do not say iodine melts by breaking covalent bonds: melting iodine overcomes intermolecular forces.
· Do not say graphite conducts because it has ions: graphite conducts due to delocalised electrons.
· Do not describe C₆₀ as giant covalent: buckminsterfullerene is simple molecular.
· Do not describe SiO₂ as simple molecular: silicon(IV) oxide is giant covalent.
· Do not explain high melting point by saying “strong structure”; specify strong ionic / covalent / metallic bonding.