Topic 3.2 Properties of Solids: relate the macroscopic properties of a solid (melting point, hardness, conductivity) to its type (ionic, metallic, covalent network, molecular) and the forces holding its particles together.

What this topic is asking

The College Board (Topic 3.2) wants you to classify a solid as ionic, metallic, covalent network or molecular, identify the particles and forces in each, and use that to explain its macroscopic properties: melting point, hardness, brittleness and electrical conductivity. This builds directly on the bonding models of Unit 2 and the intermolecular forces of Topic 3.1.

The four types of solid

The single most useful idea is that melting a solid means supplying enough energy to overcome whatever force holds its particles in place. The nature and strength of that force differs hugely between the four types, which is why their melting points span hundreds of degrees.

Properties type by type

• Ionic (for example $\text{NaCl}$, $\text{MgO}$): the lattice of $+$ and $-$ ions gives high melting points; it is brittle because shifting a layer brings like charges together; it conducts only when the ions are free to move.
• Metallic (for example $\text{Cu}$, $\text{Fe}$): the electron sea lets cations slide past one another, so metals bend rather than shatter, and the mobile electrons carry charge and heat.
• Covalent network (for example diamond, $\text{SiO}_2$, graphite): a continuous web of covalent bonds gives extreme hardness and very high melting points. Graphite is the exception that conducts, because its layered structure leaves delocalised electrons free to move within sheets.
• Molecular (for example ice, dry ice, $\text{I}_2$): the molecules are held only by intermolecular forces, so little energy is needed to melt them and there are no mobile charges to conduct.

Reasoning from structure to property

The whole point is to predict properties rather than memorize them. Given a substance, ask what its particles are and what holds them together; the melting point follows from the strength of that force, and the conductivity follows from whether there are mobile charges. A covalent network melts higher than an ionic solid because breaking covalent bonds across an entire lattice costs more than separating ions; an ionic solid melts higher than a molecular solid because ionic attractions are far stronger than intermolecular forces. Conductivity is a separate question answered by mobile charge carriers: metals always have them, ionic solids have them only when molten or dissolved, and molecular and most network solids have none.

Try this

Q1. Classify (a) graphite, (b) solid argon, (c) iron. [3 points]

• Cue. (a) covalent network (conducting exception); (b) molecular (dispersion only); (c) metallic.

Q2. Explain why silicon dioxide ($\text{SiO}_2$, quartz) has a much higher melting point than carbon dioxide ($\text{CO}_2$), even though both contain similar atoms. [2 points]

• Cue. $\text{SiO}_2$ is a covalent network solid (bonds throughout the lattice must break to melt), while $\text{CO}_2$ is molecular (only weak dispersion forces between molecules).