Bulk Properties of Chemical Structures

Solid ionic compounds do not conduct electricity because the ions are locked firmly in place. However, when melted or dissolved in water, the lattice breaks down and the ions become mobile charge carriers that are free to move and carry an electrical current.

These crystals are highly brittle and shatter easily. If a force makes the regular layers shift, ions with the same charge are pushed next to each other, and the resulting electrostatic repulsion blasts the crystal apart.

Simple Molecules

Why is the oxygen we breathe a gas, while the water we drink is a liquid, even though both are made of small groups of atoms? This comes down to the difference between forces inside and outside the molecule.

Simple molecules consist of a small, fixed number of non-metal atoms. Within the molecule, atoms are strongly joined by covalent bonds, but between different molecules, there are only weak intermolecular forces.

Simple molecules have very low melting and boiling points, meaning most are gases or liquids at room temperature. When they change state, you only need enough thermal energy to overcome the weak intermolecular forces—the strong covalent bonds between the atoms do not break.

As molecules get larger (like long-chain alkanes), these weak intermolecular forces increase in strength, resulting in higher boiling points. Simple molecules are entirely neutral and do not conduct electricity, as they have no free electrons or ions to carry a charge.

Giant Covalent Structures

You can easily rub a pencil tip across paper, but a diamond drill bit can cut through solid rock. Both are made entirely of carbon atoms, but their internal structures give them radically different properties.

A giant covalent structure is a massive 3D lattice where every single atom is connected by strong covalent bonds. Unlike simple molecular substances, individual molecules do not exist here.

Because of the vast number of covalent bonds, massive amounts of energy are required to break them. This gives them incredibly high melting and boiling points.

Diamond and graphite are allotropes of carbon. Silicon dioxide ($SiO_2$), also known as silica, forms a giant lattice very similar to diamond.

Property	Diamond	Graphite
Hardness	Extremely hard (used in cutting tools)	Soft and slippery (used as a lubricant)
Conductivity	Insulator (does not conduct electricity)	Good conductor (contains delocalised electrons)
Arrangement	3D tetrahedral lattice (4 bonds per carbon atom)	2D hexagonal layers (3 bonds per carbon atom)

Graphite's unique softness comes from the fact that its 2D layers are held together only by weak intermolecular forces. This allows the layers to easily slide over one another.

Polymers

From plastic water bottles to the DNA inside your cells, long-chain molecules form the building blocks of the modern world. A polymer is a macromolecule built from thousands of small repeating units called monomers.

The atoms within the polymeric chains are held together by strong covalent bonds. Because these molecules are extremely large, they have many points of contact, creating intermolecular forces strong enough to make most polymers solid at room temperature.

Polymers typically melt over a range of temperatures because their chains vary in length.

Thermosoftening polymers melt when heated because their tangled chains are only held by intermolecular forces, allowing them to be reshaped. Thermosetting polymers do not melt because their chains are locked rigidly together by strong covalent cross-links.

Worked Example

How do we represent the polymerisation of a monomer into a polymer chain?

Step 1: Draw the repeating unit. Identify the small repeating section, known as the repeat unit, derived from the original monomer. In this case, the $C=C$ double bond breaks open to form a $C-C$ single bond.

Step 2: Draw the polymer bracket notation. Draw brackets around the repeat unit. The bonding "arms" must go straight through the brackets to show the continuous chain.

Step 3: Add the multiplier. Add an $n$ at the bottom right outside the bracket to represent the large, variable number of repeating units.

Metallic Bonding

Why does a metal spoon feel instantly cold when you pick it up, but get blisteringly hot if left in a pan of soup? Metals are excellent conductors because of their unique internal structure.

Metals form a giant metallic lattice composed of positive metal ions surrounded by a "sea" of freely moving delocalised electrons. Metallic bonding is the strong electrostatic attraction between these positive metal ions and the negative delocalised electrons.

Most metals have high melting and boiling points. This is because the metallic bonds are very strong and extend in all directions throughout the giant lattice, requiring a large amount of thermal energy to overcome the attraction.

Because the delocalised electrons are mobile, they can move rapidly throughout the entire structure to carry both electrical charge and thermal energy.

Metals are highly malleable, meaning they can be bent into shape without breaking. This happens because the positive metal ions are arranged in regular layers that can slide over each other while maintaining the strong metallic bonds.

Alloys are mixtures of a metal with other elements. The different-sized atoms disrupt the regular layers, preventing them from sliding easily and making alloys much harder than pure metals.

Bulk Properties and Scale

A single atom of gold is not shiny, and a single molecule of water is not wet. Characteristics like density, melting point, and conductivity are bulk properties.

Bulk properties arise purely from the collective behaviour of huge numbers of atoms, ions, or molecules acting together in a structure. Individual simple molecules and atoms are incredibly small—typically around $10^{-10}\text{ m}$ in radius—and simply do not possess bulk properties on their own.