Covalent Bonding & Molecular Structures
What Is Covalent Bonding?
When two non-metal atoms react, neither atom is strong enough to take electrons from the other. Instead, the atoms share one or more pairs of electrons so that both achieve a full outer shell. This shared pair of electrons is a covalent bond. Covalent bonding produces molecules — distinct groups of atoms bonded together.
A typical atom has a radius of about \(1 \times 10^{-10}\) m, while a small molecule such as water is roughly the same order of magnitude in size. You can represent covalent bonds using dot-and-cross diagrams, which show clearly which electrons belong to which atom and how they are shared.
A covalent bond is a shared pair of electrons between two non-metal atoms. Each atom contributes one electron to the shared pair. The atoms are held together because the shared electrons are attracted to both nuclei.
Simple Molecular Substances
You need to be able to draw dot-and-cross diagrams for six key molecules:
Hydrogen (\(\text{H}_2\)) — each hydrogen atom has 1 electron. They share one pair, forming a single bond. Hydrogen chloride (HCl) — hydrogen shares its 1 electron with one of chlorine’s 7 outer electrons, forming a single bond. Water (\(\text{H}_2\text{O}\)) — oxygen has 6 outer electrons and shares one pair with each of two hydrogen atoms, forming two single bonds.
Methane (\(\text{CH}_4\)) — carbon has 4 outer electrons and shares one pair with each of four hydrogen atoms, forming four single bonds. Oxygen (\(\text{O}_2\)) — each oxygen atom has 6 outer electrons. They share two pairs of electrons, forming a double bond. Carbon dioxide (\(\text{CO}_2\)) — each oxygen shares two pairs of electrons with the central carbon atom, forming two double bonds.
Step 1: Write the electronic configuration of each atom. Step 2: Identify how many electrons each atom needs to share to fill its outer shell. Step 3: Draw the outer shells overlapping, with shared electrons in the overlap region. Use dots for one atom’s electrons and crosses for the other. Step 4: Check that each atom has a full outer shell (2 for hydrogen, 8 for most other atoms).
For example, in water: oxygen (2, 6) needs 2 more electrons, hydrogen (1) needs 1 more. Two hydrogen atoms each share one electron with oxygen, giving oxygen 8 outer electrons and each hydrogen 2.
Properties of Simple Molecular Substances
Simple molecular substances have low melting and boiling points. This is because the covalent bonds within each molecule are strong, but the forces between molecules — called intermolecular forces — are very weak. When a simple molecular substance melts or boils, it is these weak intermolecular forces that are overcome, not the covalent bonds themselves.
Simple molecular substances also do not conduct electricity because they have no free electrons or ions to carry a charge. Examples include water, methane, carbon dioxide and oxygen.
When a simple molecular substance boils, it is the weak intermolecular forces between molecules that are overcome — not the strong covalent bonds within the molecules. This is why simple molecular substances have low melting and boiling points.
Giant Covalent Structures
Some covalent substances do not form small molecules. Instead, billions of atoms are bonded together in a continuous three-dimensional network called a giant covalent structure. These have very high melting points because many strong covalent bonds must be broken to separate the atoms.
Diamond and Graphite
Diamond and graphite are both forms of carbon — they are allotropes. In diamond, each carbon atom bonds to four others in a rigid tetrahedral arrangement. There are no free electrons, so diamond does not conduct electricity. Its hardness makes it ideal for cutting tools.
In graphite, each carbon atom bonds to only three others, forming flat layers. The fourth outer electron on each carbon atom is delocalised — free to move between the layers. These delocalised electrons allow graphite to conduct electricity, which is why graphite is used to make electrodes. The layers are held together by weak intermolecular forces, so they can slide over each other, making graphite soft and useful as a lubricant.
Fullerenes are carbon molecules formed into hollow shapes. The most famous is buckminsterfullerene, \(\text{C}_{60}\), which has 60 carbon atoms arranged in a sphere of hexagons and pentagons, resembling a football. Fullerenes can be used to deliver drugs to specific cells in the body and as catalysts.
Graphene is a single layer of graphite — just one atom thick. Each carbon atom is bonded to three others in a hexagonal pattern, with one delocalised electron per atom. Graphene is extremely strong, lightweight, and an excellent conductor of both electricity and heat. It has potential applications in electronics, composites and touchscreens.
Polymers
A polymer is a very long chain molecule made by joining together thousands of small molecules called monomers. For example, poly(ethene) is formed when many ethene molecules join together. The polymer chain consists of a long backbone of carbon atoms with hydrogen atoms attached. Polymers are simple molecular in the sense that their molecules — although very large — are held to each other by weak intermolecular forces.
Metallic Bonding
In metals, atoms are packed closely together in a regular arrangement. The outer electrons become delocalised — they are free to move throughout the entire structure. This creates a “sea” of delocalised electrons surrounding a lattice of positive metal ions. The strong electrostatic attraction between the positive ions and the delocalised electrons is a metallic bond.
Metallic bonding explains why metals are good conductors of electricity — the delocalised electrons can flow through the structure when a voltage is applied. Metals are also malleable because the layers of ions can slide over each other without breaking the metallic bonds, since the electron sea simply adjusts around them.
Metals conduct electricity because delocalised electrons can flow through the structure. Metals are malleable because the layers of ions can slide past each other without the bonding breaking down.
Dot-and-cross diagrams, ball-and-stick models, and 2D/3D representations are all useful for understanding bonding, but each has limitations. Dot-and-cross diagrams show electron arrangement clearly but not the 3D shape of a molecule. Ball-and-stick models show 3D shape but not electron distribution. 2D diagrams cannot fully represent 3D structures like diamond or giant ionic lattices. It is important to recognise that models are simplifications — they help us understand aspects of reality, but no single model captures everything.