Chemistry: Materials and their structure

Everything is made of atoms. At the centre of each atom is the nucleus, containing positive protons and neutral neutrons. Whizzing around the nucleus in specific energy levels, or shells, are negative electrons. The number of protons, called the atomic number (Z), defines the element. For example, any atom with 6 protons is carbon. The mass number (A) is the total count of protons and neutrons. In a neutral atom, the number of electrons equals the number of protons. The arrangement of these electrons determines an element's chemical properties and its place in the Periodic Table. The number of electron shells tells you the period (row) it's in, and the number of electrons in the outermost shell tells you its group (column). For example, Sodium (Na) with electron configuration 2.8.1 is in Period 3 and Group 1.

Atoms react to achieve a stable electron configuration, usually a full outer shell of electrons, like a noble gas. In ionic bonding, this happens by transferring electrons from a metal atom to a non-metal atom. The metal atom loses electrons to form a positively charged ion (a cation). The non-metal atom gains these electrons to form a negatively charged ion (an anion). The resulting oppositely charged ions are then strongly attracted to each other by electrostatic forces. This is the ionic bond. This process doesn't just form one pair of ions; it creates a huge, regular, repeating structure of alternating positive and negative ions called a giant ionic lattice.

When non-metal atoms react together, they also want to achieve a stable, full outer shell of electrons. Instead of transferring electrons, they share them. A covalent bond is a shared pair of electrons between two atoms. The positive nuclei of both atoms are electrostatically attracted to the shared pair of negative electrons, holding the atoms together strongly. Substances made of covalent bonds can either be simple molecules (like water, H₂O, or carbon dioxide, CO₂) or giant structures. In simple molecules, the covalent bonds within the molecule are very strong, but the forces of attraction between the molecules (intermolecular forces) are very weak.

The type of bonding and structure a substance has determines its physical properties, like melting point and electrical conductivity. By knowing the structure, you can predict the properties.

Ionic Compounds: They have a giant ionic lattice structure. The electrostatic forces of attraction between the oppositely charged ions are very strong and require a large amount of energy to overcome. This results in high melting and boiling points. In the solid state, the ions are fixed in the lattice and cannot move, so they do not conduct electricity. However, when molten or dissolved in water, the ions are free to move and carry charge, so they do conduct electricity.

Simple Covalent Molecules: These have weak intermolecular forces between the molecules. Very little energy is needed to overcome these weak forces. Consequently, they have low melting and boiling points (many are gases or liquids at room temperature). They do not have any free-moving charged particles (no ions or delocalised electrons), so they do not conduct electricity in any state.

Some substances consist of billions of atoms joined together by a network of strong covalent bonds. These are called giant covalent structures or macromolecules. Because all the atoms are linked by strong covalent bonds, these substances have very specific properties. The two most famous examples are allotropes of carbon: diamond and graphite. Allotropes are different structural forms of the same element.

Diamond: Each carbon atom is covalently bonded to four other carbon atoms in a rigid tetrahedral arrangement. This creates a very strong, hard, 3D structure. It has a very high melting point and does not conduct electricity as there are no free electrons.

Graphite: Each carbon atom is covalently bonded to three other carbon atoms, forming hexagonal rings in flat layers. The layers are held together by weak intermolecular forces, so they can slide over each other, making graphite soft and slippery. Each carbon atom has one delocalised electron which is free to move along the layers, allowing graphite to conduct electricity.

Silicon Dioxide (Silica): Found in sand and quartz, it has a structure similar to diamond. Each silicon atom is bonded to four oxygen atoms, and each oxygen to two silicons. It is also hard and has a very high melting point.

Metals have a unique type of bonding that explains their characteristic properties. A metal consists of a giant, regular lattice of positive metal ions. The outer shell electrons from each metal atom are lost and become 'delocalised', meaning they are free to move throughout the entire structure. This creates a 'sea' of delocalised electrons surrounding the fixed positive ions. A metallic bond is the strong electrostatic attraction between the positive metal ions and the delocalised sea of negative electrons. This strong attraction means metals have high melting and boiling points. The mobile electrons explain why metals are excellent conductors of electricity and heat. The regular lattice structure explains why metals are malleable (can be hammered into shape) and ductile (can be drawn into wires); the layers of ions can slide over each other without breaking the metallic bond.