Lattice Structures and Physical Properties

Why does salt melt at $801\ ^\circ\text{C}$ while iodine turns to vapour on gentle warming, and why does graphite conduct electricity when diamond, also pure carbon, does not? The answers lie in the structure of the solid: what particles it contains and what holds them together. This note describes the four types of crystalline lattice you must know, with the named examples, and shows how to explain and predict melting point, boiling point, electrical conductivity and solubility from structure, and how to work backwards from data to deduce the structure. These are high-frequency explain questions in Papers 1 and 2.

Crystalline solids and lattices

In a crystalline solid, the particles are arranged in a regular, repeating three-dimensional pattern called a lattice. The particles might be ions, atoms or molecules. The type of particle, and the forces between the particles, decide the properties.

The word "giant" means the bonding extends throughout the whole crystal: there are no separate molecules, and the formula (such as $\ce{NaCl}$ or $\ce{SiO2}$) gives only the ratio of particles.

Giant ionic lattices

Sodium chloride and magnesium oxide

In sodium chloride, each $\ce{Na+}$ ion is surrounded by six $\ce{Cl-}$ ions, and each $\ce{Cl-}$ ion by six $\ce{Na+}$ ions, in a cubic arrangement. The electrostatic attractions act in all directions, and there are millions of them in every crystal. Magnesium oxide has the same structure, with $\ce{Mg^2+}$ and $\ce{O^2-}$ ions.

Properties of ionic compounds

property	explanation
high melting and boiling points	a lot of energy is needed to overcome the many strong electrostatic attractions between oppositely charged ions throughout the lattice
do not conduct when solid	ions are held in fixed positions and cannot move
conduct when molten or in aqueous solution	the ions are free to move and carry charge
often soluble in water	water molecules are polar; their $\delta-$ oxygen atoms attract cations and their $\delta+$ hydrogen atoms attract anions, and the energy released helps pull the ions out of the lattice
insoluble in non-polar solvents	non-polar molecules cannot attract the ions strongly enough to separate them
hard but brittle	a sharp blow shifts a layer so that ions of the same charge line up and repel, and the crystal shatters

Magnesium oxide melts at $2852\ ^\circ\text{C}$, far higher than sodium chloride ($801\ ^\circ\text{C}$). Its ions carry double the charge ($2+$ and $2-$) and are smaller, so the attraction between them is much stronger. MgO is also much less soluble in water, and it is used to line furnaces.

Giant metallic lattices

Copper

Copper atoms are packed closely in regular layers. Each atom releases outer electrons into a sea of delocalised electrons that move freely throughout the lattice, leaving positive copper ions in fixed positions. The metallic bonding (the electrostatic attraction between the positive ions and the delocalised electrons) acts in all directions.

property	explanation
high melting point (copper: $1085\ ^\circ\text{C}$)	strong attraction between the positive ions and the delocalised electrons must be overcome
conducts electricity as solid and liquid	delocalised electrons are free to move through the lattice and carry charge when a voltage is applied
good thermal conductor	delocalised electrons carry kinetic energy quickly through the lattice
malleable and ductile	layers of positive ions can slide over one another without breaking the metallic bonding, because the delocalised electrons move with them and the bonding is not directional
insoluble in water and organic solvents	(some react with water, but none simply dissolve)

Copper is used for electrical wiring because it is an excellent conductor and very ductile (easily drawn into wires).

Giant molecular (giant covalent) lattices

In a giant molecular structure, every atom is joined to its neighbours by strong covalent bonds in a network extending through the whole crystal. To melt it, covalent bonds must be broken, so melting points are very high.

Diamond

Each carbon atom is covalently bonded to four other carbon atoms, arranged tetrahedrally ($109.5^\circ$; the carbon atoms are sp³ hybridised). The network extends in three dimensions.

• Very high melting point (diamond and graphite both sublime at around $3600\ ^\circ\text{C}$ at atmospheric pressure): many strong covalent bonds must be broken.
• Extremely hard: every atom is held rigidly in three dimensions by strong bonds. Used in cutting tools and drill tips.
• Does not conduct electricity: all four outer electrons of each carbon are used in localised covalent bonds, so there are no delocalised electrons or ions to move.
• Insoluble in all solvents: the strong covalent bonds cannot be broken by any interaction with solvent molecules.

Graphite

Each carbon atom is covalently bonded to three others in flat hexagonal layers (sp² carbon, $120^\circ$). Each carbon has one outer electron left over, which is delocalised across the layer. The layers are stacked, with weak id-id forces between them, and are relatively far apart.

• Very high melting point: the strong covalent bonds within the layers must be broken.
• Conducts electricity: the delocalised electrons can move along the layers when a voltage is applied. (It conducts poorly at right angles to the layers.) Used for electrodes.
• Soft and slippery: the weak forces between the layers let them slide over each other easily. Used as a lubricant and in pencils.
• Insoluble in all solvents.

Silicon(IV) oxide

In silicon(IV) oxide, $\ce{SiO2}$ (for example quartz and sand), each silicon atom is covalently bonded to four oxygen atoms arranged tetrahedrally, and each oxygen atom is bonded to two silicon atoms. That is why the ratio is $1 : 2$. The structure resembles diamond with an oxygen atom between every pair of silicon atoms.

• High melting point (about $1700\ ^\circ\text{C}$): strong Si–O covalent bonds must be broken.
• Hard; does not conduct (no delocalised electrons, no ions); insoluble in water.

Simple molecular lattices

In a simple molecular solid, small molecules are held in a regular lattice by weak intermolecular forces. Within each molecule the atoms are joined by strong covalent bonds, but these are not broken on melting or boiling.

Iodine

Iodine crystals contain $\ce{I2}$ molecules arranged in a regular lattice, held together by id-id forces. Because each molecule has many electrons (106), these forces are strong enough to make iodine a solid at room temperature, but still weak compared with bonds.

• Low melting point ($114\ ^\circ\text{C}$); on gentle heating, iodine sublimes, turning straight to a purple vapour.
• Does not conduct: no ions and no delocalised electrons.
• Only slightly soluble in water (it is non-polar), but dissolves well in non-polar solvents such as hexane, giving a purple solution, and in aqueous potassium iodide (by forming the $\ce{I3-}$ ion).

Buckminsterfullerene, C60

$\ce{C60}$ is a molecule of 60 carbon atoms arranged in a hollow sphere of 20 hexagons and 12 pentagons, like a football. Each carbon is bonded to three others. In the solid, the $\ce{C60}$ molecules are packed in a regular lattice held by id-id forces.

• Much lower melting / sublimation point than diamond or graphite: only the weak forces between molecules are overcome; no covalent bonds break.
• Soft compared with diamond.
• Poor electrical conductor: although some electrons are delocalised over each molecule, they cannot move from one molecule to the next.
• Insoluble in water, but soluble in some non-polar organic solvents such as methylbenzene.

Ice

In ice, each water molecule is hydrogen bonded to four others in a tetrahedral arrangement, forming an open lattice. Ice is a simple molecular solid: melting breaks only some of the hydrogen bonds, not the O–H bonds. Its melting point ($0\ ^\circ\text{C}$) is low compared with giant structures but high for such a small molecule, because hydrogen bonds are the strongest intermolecular forces. The open structure is why ice is less dense than water.

Properties of simple molecular substances

property	explanation
low melting and boiling points	only weak intermolecular forces need to be overcome; covalent bonds are not broken
do not conduct in any state	molecules are neutral; there are no ions or delocalised electrons that can move through the substance
soluble in non-polar solvents (if non-polar)	similar weak forces between solute and solvent molecules
soluble in water only if polar enough to hydrogen bond with water (for example ethanol, ammonia)

Summary of properties by structure

Deducing structure from data

Worked examples

Practice

Answers

1. A giant ionic lattice in which each $\ce{Na+}$ is surrounded by six $\ce{Cl-}$ and each $\ce{Cl-}$ by six $\ce{Na+}$, held by strong electrostatic attractions in all directions. In the solid the ions are fixed in position and cannot move. In the molten state the lattice has broken down and the ions are free to move and carry charge.
2. Both are giant ionic lattices with the same arrangement. The $\ce{Mg^2+}$ and $\ce{O^2-}$ ions have double the charges of $\ce{Na+}$ and $\ce{Cl-}$ and are smaller, so the electrostatic attraction between them is much stronger. Much more energy is needed to overcome these attractions.
3. A giant metallic lattice of positive copper ions in regular layers, surrounded by a sea of delocalised electrons; metallic bonding is the electrostatic attraction between the ions and the delocalised electrons. When a force is applied, layers of ions slide over each other; the delocalised electrons move with them, so the bonding is maintained and the metal changes shape without breaking.
4. A giant molecular (covalent) lattice. Each silicon atom is covalently bonded to four oxygen atoms arranged tetrahedrally, and each oxygen is bonded to two silicon atoms. Each Si "owns" a half-share of four O atoms, so the ratio is $1 : 4 \times \tfrac{1}{2} = 1 : 2$.
5. Lubricant: the layers are held together only by weak id-id forces, so they slide over each other easily. Electrodes: one electron per carbon atom is delocalised, so graphite conducts electricity; its very high melting point and lack of reactivity also help.
6. $\ce{C60}$ is a simple molecular substance: sublimation overcomes only the weak id-id forces between $\ce{C60}$ molecules. Diamond is giant molecular: sublimation requires breaking many strong C–C covalent bonds throughout the lattice, which needs far more energy.
7. Ice is a simple molecular solid: melting overcomes only some intermolecular forces (hydrogen bonds), which need much less energy than the strong ionic bonds throughout the NaCl lattice. Compared with $\ce{H2S}$, the forces between water molecules are hydrogen bonds, which are much stronger than the pd-pd and id-id forces between $\ce{H2S}$ molecules (sulfur is not electronegative enough to form hydrogen bonds), so more energy is needed to melt ice.
8. Giant molecular (giant covalent) structure, for example silicon. Very high melting point and hardness mean strong covalent bonds throughout a giant lattice must be broken; insolubility fits a covalent network; poor conductivity rules out a metal or an ionic solid that dissolves. (Silicon melts at $1414\ ^\circ\text{C}$; it is a semiconductor, conducting slightly.)
9. Aluminium chloride: electronegativity difference $3.0 - 1.5 = 1.5$, intermediate: covalent bonding, simple molecular (as $\ce{Al2Cl6}$ molecules). Low sublimation point because only weak intermolecular forces are overcome; no ions, so the liquid does not conduct. Aluminium oxide: difference $3.5 - 1.5 = 2.0$, large: ionic bonding, giant ionic lattice of $\ce{Al^3+}$ and $\ce{O^2-}$. High melting point because strong electrostatic attractions between highly charged ions must be overcome; the molten oxide conducts because the ions are free to move.
10. Shared property: soft and slippery (layers held by weak forces slide over each other) or very high melting point (strong covalent bonds in the layers). Insulator: in boron nitride the electrons that would be delocalised are held as a lone pair on each nitrogen atom, because nitrogen is much more electronegative than boron. The electrons stay localised on the nitrogen atoms and are not free to move along the layers, so there are no mobile charge carriers.