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Have you ever wondered why hydrogen and helium, both light gases historically (and currently) used in balloons, behave so differently? Hydrogen is highly reactive and flammable, while helium is unreactive and stable. The key lies in the stability of noble gases, whose atoms have full outer electron shells.
To achieve the stability of a full outer electron shell, atoms of other elements tend to form chemical bonds, which determine the physical properties of the materials they form. In this article, we will explore the octet structure of a full valence shell, how different types of chemical bonding help atoms achieve stability, and how bonding affects the properties of materials.
Noble gases make up group VIII, and sit at the far-right corner of the periodic table. They are extremely stable, and usually do not need to form chemical bonds with other atoms to achieve such stability. This is because their outermost shell, known as the valence shell, is already full. A full valence shell is extremely stable, making it less reactive.
Other atoms are not so lucky, so they try to “copy” the noble gas by achieving a full valence shell too! There are several ways this can happen.
In all cases, chemical bonding helps atoms reach a more stable electronic configuration. These bonded atoms form substances and materials. These can be:
When atoms gain or lose electrons, they become charged particles called ions. An atom that loses electrons and becomes positively charged is called a cation, while one that gains electrons and becomes negatively charged is called an anion.
Atoms usually give up or gain just enough electrons to achieve a full valence shell. Most ions have charges of up to +3 or -3. Why is that so? Although it may sound easy to just gain or lose electrons, moving too many at once would take a lot of energy and effort. After all, you wouldn’t be so keen on letting go of more electrons than you need to, would you?
Also, not all ions are single atoms. Some groups of atoms can carry an overall charge too! These are called polyatomic ions. Some common polyatomic ions are shown below:
The law of conservation states that all matter must be conserved. When an atom wants to give away its valence electrons to become a stable cation, there must be a recipient that can accept those electrons to become a stable anion! This transfer of electrons form a cation and anion, which are electrostatically attracted to each other, forming an ionic bond. Ionic bonds typically form between a metal cation and non-metal anion. The extensive interactions between the positively charged and negatively charged ions form a giant lattice, in which the ions are held by electrostatic attraction.
Ionic compounds are known for their high melting and boiling points, brittleness and being a solid crystal at room temperature.
The ions packed together in the giant lattice are held in place by strong electrostatic attractions. To melt or boil an ionic compound, you would need to supply a large amount of energy to break these strong forces, resulting in very high melting and boiling points.
At the same time, the lattice is very rigid! If you hit or squeeze the crystal, the lattice structure cracks and shatters like glass. That’s why ionic compounds are both strong and brittle.
What happens if there are no metal atoms around to give up their valence electrons? In that case, non‑metal atoms have to work together differently, and share their electrons instead.
In a covalent bond, two non‑metal atoms share a pair of electrons, so that each atom “completes” its valence shell and becomes more stable. This sharing creates covalent compounds, where atoms are held together by shared electron pairs rather than transferred ones.
Deducing covalent structures is often the trickiest part of understanding covalent bonding! When you draw dot‑cross diagrams, you have to figure out which electrons come from which atom and how they are shared.
Practice makes perfect, but here are some useful tips to guide you:
Take methane as an example. We know that a carbon atom has 4 valence electrons, so it needs 4 more electrons to achieve a full valence shell. It can gain these 4 electrons by sharing 4 pairs of electrons. This happens when carbon forms 4 single covalent bonds with 4 hydrogen atoms, giving methane its familiar formula, CH₄.
Now what about double bonds? In some cases, each atom has enough valence electrons to contribute more than one electron to the shared pair. Instead of sharing just one pair of electrons, the atoms can share two pairs. This is called a double bond. In even more extreme cases, atoms can share three pairs of electrons, forming a triple bond!
Take oxygen, O₂, as an example. Each oxygen atom has 6 valence electrons and needs 2 more, so they can share two pairs of electrons, forming a strong double bond between them.
Simple covalent compounds have low melting and boiling points, and poor electrical conductivity. They are usually gases, liquids or low-melting solids at room temperature.
While the covalent bonding between atoms in a molecule is extremely strong, inter-molecular interactions between molecules are weak. This results in low melting and boiling points, and high volatility. Covalent compounds also lack free electrons or ions that can carry charge, and are hence poor conductors of electricity.
The covalent bonds between atoms in a simple covalent compound are already very strong. Now imagine an entire giant structure made of atoms held together by covalent bonds. That is a giant covalent compound.
Because the covalent bonds extend throughout the entire structure, giant covalent compounds form regular repeating lattices instead of small individual molecules. Since a single covalent bond is already very strong, a giant network of these bonds makes the substance extremely hard to break apart. As a result, giant covalent compounds usually have very high melting points and are extremely hard.
Two well‑known examples of giant covalent compounds made of carbon are diamond and graphite. Both are elements composed entirely of carbon, but because their atoms are arranged differently in the covalent lattice, they are called allotropes, different structural forms of the same element in the same physical state.
In diamond, each carbon atom forms a single covalent bond with four neighbouring carbon atoms, creating a rigid, three‑dimensional tetrahedral network. This strong, continuous bonding in all directions makes diamond extremely hard and gives it a very high melting point.
In graphite, each carbon atom forms single covalent bonds with three neighbouring carbon atoms, lying in flat, hexagonal layers. The fourth valence electron from each carbon atom becomes delocalisedbetween the layers, much like the “sea of electrons” in metals (which we will cover later). This delocalised electron cloud allows graphite to conduct electricity, a rare property among covalently bonded materials. The flat, hexagonal layers also allow each layer to slip over each other, giving graphite its lubricating properties.
Giant covalent structures can also be made from more than one type of atom. A classic example is silicon dioxide (SiO₂), a prominent component of sand. Like diamond, silicon and oxygen atoms are linked by extensive covalent bonds, forming a giant repeating lattice with a tetrahedral arrangement.
Each silicon atom has four valence electrons and needs four more to complete its valence shell, so it forms single covalent bonds with four oxygen atoms. At the same time, each oxygen atom has six valence electrons and needs two more, so it forms single covalent bonds with two silicon atoms. This continuous network of strong covalent bonds gives silicon dioxide the same key features as diamond, namely its high hardness, high melting point, and high boiling point.
Polymers, even though they also involve extensive covalent bonds, are technically not giant covalent compounds. Rather, they are very large, individual covalent molecules (macromolecules), while giant covalent structures are rigid, continuous 3D networks of atoms.
Polymers are long chains made up of repeating units called monomers, which are small covalent molecules. These monomers are joined together by covalent bonds to form a polymer chain.
One common example is polyethene (also known as polyethylene), which is made from repeating ethene (ethylene) units linked into a long chain. While the atoms within each chain are held together by strong covalent bonds, the intermolecular forces between neighbouring chains are much weaker. This is why polymers like polyethene usually have lower melting points and greater flexibility compared to giant covalent structures.
Many everyday synthetic materials are made from polymers, including polyethene (PE), polypropylene (PP), and polyvinyl chloride (PVC).
Metal atoms are really good at letting go of their valence electrons to achieve a full valence shell, because they usually only need to give up a few electrons. But what happens when there aren’t any non‑metal atoms around to receive them?
In a metal, metal atoms team up with each other. Each metal atom releases its valence electrons, forming positive metal cations arranged in a lattice, while the freed electrons move around the whole structure like a “sea”. This forms metallic bonds, a strong electrostatic attraction between positively charged metal cations and a shared sea of delocalised electrons.
Metallic bonding explains many of the key properties of metals, such as high melting and boiling points, electrical conductivity, malleability and ductility.
The strong electrostatic attraction between the cations and the sea of electrons gives metals high melting and boiling points. The delocalised electrons can move easily through the structure, carrying charge and allowing metals to conduct electricity very well.
Unlike ionic or covalent bonds, metallic bonds are non‑directional, meaning the electrons and cations can slide past one another without breaking the overall bonding. This is why metals are malleable and ductile. You can observe this in action when bending or twisting a metal wire without it snapping.
In pure metals, the metal cations are the same size, so they can slide smoothly over each other, making the metal very malleable. However, in alloys, which is a mixture of different types of metals, the cations are different sizes and don’t slide as easily. This makes alloys stronger and less malleable than pure metals, which is why we use alloys like steel in buildings and bridges.
The chemical structure and the types of bonds between atoms can tell you a lot about a material’s properties! Now that you’ve learned about chemical bonds and how they give rise to the different materials we use every day, take a moment to think about your favourite common objects, like glass, metal, plastic, or even salt, and how their bonding shapes their properties.
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Prepared by: Michelle
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