1. Molecules come in all shapes and sizes

Covalent bonds join atoms together in different ways, forming small and big molecules. While this article is about the big, to appreciate the uniqueness and enormity of giant molecular structure, let’s first recap on the small.

Discrete molecules of oxygen (© 2014-2019 OUP and Nick Greeves).

Oxygen gas in the air has a simple molecular structure. Each diatomic molecule is really small, with just two atoms per molecule.

The simple molecular structure of sucrose, C12H22O11. Each discrete molecule is highlighted in brown. The dotted lines between the molecules represent intermolecular forces of attraction.

There are bigger simple molecules, like the happiness-inducing sucrose found in all things sweet. Each sucrose molecule, C12H22O11 has 45 atoms.

A molecule of the protein hexokinase. Shown beside are the relatively tiny molecules of oxygen and sucrose.

Even bigger? We have the protein hexokinase, one of the many enzymes in charge of converting your sucrose-rich brown sugar milk tea into energy! Each molecule of the protein has 16000 atoms.

Diamond (left) is a giant molecule of countless carbon atoms represented by black balls. It dwarves a sm0ll discrete molecule of hexokinase protein (right). What an absolute unit!

What about the hugest?! Unlike simple molecules of oxygen, sucrose, and hexokinase protein, diamond has a giant molecular structure. A one-carat diamond weighing 0.2 g has 10,000,000,000,000,000,000,000 carbon atoms, all covalently bonded to each other to form one giant molecule that extends vastly.

Giant molecular structures are an extensive network of atoms joined together by strong covalent bonds.

So how do the many carbon atoms join up? Let’s zoom in on a carbon atom of diamond to find out.

2. In diamond, each carbon atom bonds covalently with four other atoms

A simplified 2D diagram showing a carbon atom of diamond bonded to four surrounding carbon atoms. The surrounding carbon atoms have incomplete octet and hence can continue to bond with even more carbon atoms, extending the structure perpetually.

Carbon is a non-metal from Group IV, with 4 valence electrons. It is short of 4 electrons to complete its octet and achieve noble gas electronic configuration.

Focusing on the central carbon atom, it achieves the noble gas electronic configuration by forming 4 single covalent bonds with 4 surrounding carbon atoms. The surrounding carbon atoms can still form more single covalent bonds with yet more carbon atoms. This will extend the network of covalent bonds between many carbon atoms, hence establishing a giant molecular structure.

As all the electrons of every carbon atom are involved in covalent bonding, there is no free-moving electron to conduct electricity.

With neither free-moving electrons or ions, diamond does not conduct electricity.

3. A giant molecule of extremes: high melting point, insane hardness, and impossible to dissolve

A 3D diagram of the giant molecular structure of diamond, which represents the carbon atoms as black balls and the strong covalent bonds as the lines. The grey lines represent covalent bonds that extend the structure beyond what is shown here.

Unlike simple molecules, diamond has a very high melting point. To melt diamond is to break the covalent bonds between carbon atoms. They are very strong, hence requiring a large amount of energy to break.

Diamond is hard, in both senses of the word. It is hard to own with its sky-high price. It is also the hardest material, owing to the strong covalent bonds between carbon atoms that take extraordinary force to break. This physical property makes diamond an excellent material for cutting tools.

In other news, it is insoluble in both water and organic solvents. It’s weird to see your precious diamond dissolve huh.

Diamond has a high melting point, is hard, and cannot dissolve in water. This is because of the strong covalent bonds between its carbon atoms that take a large amount of energy to break.

4. A cousin of diamond: graphite as another allotrope of pure carbon

Diamond (left) and graphite (right)

Besides diamond, carbon atoms can be arranged differently to give another allotrope: graphite. The different structures give rise to vastly different physical properties.

However, graphite and diamond have the same chemical composition and properties. They comprise carbon atoms that can oxidise at high temperature to form carbon dioxide gas.

Allotropes are different structural forms of an element.

5. Delocalised, wild and free: graphite conducts electricity

Each layer as an extensive network of covalent bonds between carbon atoms. Between layers are weak intermolecular forces of attraction.

In graphite, each carbon atom uses just 3 of its valence electrons to form 3 single covalent bonds with 3 other carbon atoms. The remaining electron is delocalised and can move about freely.

Due to the presence of delocalised, free-moving electrons, graphite can conduct electricity.

6. Hippity hoppity, graphite is so slippery

While the covalent bonds within a layer are strong, we can easily overcome the weak intermolecular forces of attraction between layers by applying a small force. Therefore, graphite is slippery.

This is why we make pencil “lead” with graphite. When we write, the layers of graphite come off easily, leaving a thin layer of black graphite we call pencil mark. Furthermore, we make lubricants with graphite, like the one to oil bicycle chain.

7. Copycat properties: high melting point and impossible to dissolve

Within each extensive layer of carbon atoms in graphite, the strong covalent bonds require a large amount of energy to break, hence conferring graphite a high melting point.

Also, like diamond, graphite is insoluble in water and organic solvents. This is during paper chromatography, we mark the starting line with pencil. The graphite marking will not dissolve in the solvent to affect the results.