Types of Solids & Properties

Full Notes

Intermolecular Forces and Macroscopic Properties

The physical properties of liquids and solids – such as boiling point, melting point, and vapor pressure – are directly related to the strength of attraction forces between particles in them.

Vaporization and Vapor Pressure

Vaporization is the process by which particles in a liquid gain enough energy to escape into the gas phase.

In any liquid, some particles have enough kinetic energy to overcome the intermolecular attractions holding them in the liquid, allowing them to vaporize (evaporation).

The vapor pressure is the pressure exerted by the vapor that forms above a liquid in a closed container, due to these escaping particles.

Boiling Point

The boiling point is the temperature at which a liquid’s vapor pressure equals the atmospheric pressure acting on it.

At this point, molecules throughout the liquid—not just at the surface—have enough energy to overcome intermolecular forces and escape into the gas phase.

• Below the boiling point, vaporization (evaporation) only occurs at the surface.
• When vapor pressure equals atmospheric pressure, bubbles of vapor form within the liquid — this is boiling.
• Stronger IMFs mean a higher temperature is required for this, giving a higher boiling point.

Melting Point

Melting involves rearranging particles from a fixed solid structure into a more mobile liquid phase. Since particles stay in close contact in both phases, IMFs are not completely overcome. Still, stronger IMFs usually mean higher melting points, although the correlation can be less direct than with boiling points because particle arrangement and packing also matter.

Types of Solid

I. Ionic Solids

Made of positive and negative ions in a rigid lattice. Held together by strong electrostatic attractions (ionic bonding).

• High melting/boiling points: strong ionic bonds.
• Low vapor pressure: ions are tightly bound.
• Brittle: shifting layers aligns like charges causing repulsion, breaking the structure.
• Conduct electricity only when molten or in solution (ions mobile). Do not conduct as solids.

Example NaCl(s) forms a 3D ionic lattice of Na⁺ and Cl⁻ ions.

II. Covalent Network Solids

Composed of atoms bonded covalently in large 2D or 3D structures (typically nonmetals/metalloids).

• Very high melting points: covalent bonds must be broken.
• Hard and rigid: fixed bond angles in the network.
• Insoluble; non-conductive (graphite is an exception—conducts within layers due to delocalized electrons and is soft because layers slide).

Example Diamond (C) – 3D sp³ network; extremely hard and high melting point.

Example Graphite (C) – layered 2D networks; conducts along planes; soft due to weak forces between sheets.

Example SiO₂ – extended covalent lattice of Si–O bonds; very high melting point and hardness.

III. Molecular Solids

Made of discrete molecules held together by intermolecular forces (London dispersion, dipole–dipole, hydrogen bonding).

• Low melting/boiling points: IMFs are relatively weak.
• Non-conductive: no mobile charges.
• Soft or volatile depending on molecular structure and IMF strength.

Example Ice (H₂O) – molecules are held in a lattice by hydrogen bonds.

Example Iodine (I₂) – discrete I₂ molecules arranged in a crystal; solids held by dispersion forces between molecules.

IV. Metallic Solids

Metal atoms arranged in a lattice with delocalized “sea of electrons”.

• Excellent conductors of heat and electricity (mobile electrons).
• Malleable and ductile: layers of atoms can slide without breaking metallic bonding.

Example Copper (Cu) – metallic lattice with delocalized electrons.

Alloys

Mixtures of metals with other elements (metal or nonmetal).

• Substitutional alloys: similar-sized atoms replace host atoms in lattice (e.g., brass = Cu + Zn).
• Interstitial alloys: small atoms fit into holes in metal lattice (e.g., carbon in steel).

Effects of alloying: increased rigidity, reduced malleability and electrical conductivity generally maintained due to mobile electrons.

Example Brass (Cu + Zn) – substitutional alloy where Zn replaces some Cu atoms in the lattice.

Example Steel (Fe + C) – interstitial alloy with small C atoms occupying spaces between Fe atoms.

Large Biomolecules and Polymers

Polymers are large molecules made of repeating units formed from smaller molecules called monomers. Biological macromolecules (proteins, DNA) have specific 3D shapes defined by noncovalent interactions (hydrogen bonds, ionic attractions, London dispersion forces). These interactions determine folding, flexibility, and reactivity, strongly affecting function.

Example DNA double helix – two antiparallel strands held together by hydrogen bonding between base pairs.

Example Protein folding – interactions among amino-acid side chains (R groups) such as hydrogen bonding, ionic bonds, disulfide bridges, and hydrophobic interactions drive tertiary structure.

Matt’s exam tip

Use structure–property relationships to explain behavior:
Ionic = high melting points, brittle, conductive when molten/aqueous.
Covalent network = high melting points, rigid, non‑conductive (graphite conducts in-plane).
Molecular = low melting point, non‑conductive.
Metallic = conductive, malleable, shiny.

Summary

• Stronger interparticle forces → higher melting/boiling points and lower vapor pressure.
• Ionic solids: lattice of ions; high MP/BP; brittle; conduct only when ions are mobile (molten/aqueous).
• Covalent network solids: extended covalent bonding; very high MP; hard/rigid; typically non‑conductive (graphite conducts in layers).
• Molecular solids: discrete molecules held by IMFs; low MP/BP; non‑conductive.
• Metallic solids: cations in a sea of electrons; conductive, malleable, ductile; shiny.
• Alloys (substitutional/interstitial) adjust rigidity, malleability, and sometimes conductivity by changing how atoms pack.
• Biomolecules/polymers: noncovalent interactions (H‑bonding, ionic, dispersion, hydrophobic) control folding and function.