Bonding
Why Atoms Bond
Atoms are most stable when they have a full outer shell of electrons. Most atoms do not naturally have a full outer shell, so they bond with other atoms to achieve one. There are three main types of chemical bond: ionic bonding, covalent bonding and metallic bonding. All chemical bonds involve electrostatic forces — the attraction between positive and negative charges. The type of bond that forms depends on whether the atoms involved are metals or non-metals and how many electrons are in their outer shells.
Ionic Bonding
Ionic bonding occurs between metals and non-metals. The metal atom transfers one or more electrons to the non-metal atom. The metal atom loses electrons and becomes a positively charged ion (cation). The non-metal atom gains electrons and becomes a negatively charged ion (anion). The oppositely charged ions are then held together by strong electrostatic forces of attraction — this is the ionic bond.
For example, in sodium chloride (NaCl): sodium (2, 8, 1) loses one electron to become \(\text{Na}^+\) (2, 8), and chlorine (2, 8, 7) gains one electron to become \(\text{Cl}^-\) (2, 8, 8). Both ions now have full outer shells, making them stable. In magnesium oxide (MgO), magnesium (2, 8, 2) loses two electrons to become \(\text{Mg}^{2+}\), and oxygen (2, 6) gains two electrons to become \(\text{O}^{2-}\). In an ionic compound, billions of ions are arranged in a regular three-dimensional structure called a giant ionic lattice, where each ion is surrounded by oppositely charged ions in all directions.
In ionic bonding, electrons are transferred from metal to non-metal. The resulting ions have opposite charges and attract each other. Ionic compounds form giant ionic lattices.
Covalent Bonding
Covalent bonding occurs between non-metal atoms. Instead of transferring electrons, the atoms share one or more pairs of electrons. Each shared pair of electrons is one covalent bond. By sharing, both atoms achieve full outer shells. A single shared pair is a single bond, two shared pairs form a double bond (e.g. \(O_2\)), and three shared pairs form a triple bond (e.g. \(N_2\)).
For example, in a molecule of water (\(H_2O\)), the oxygen atom shares one pair of electrons with each of two hydrogen atoms, forming two covalent bonds. In chlorine gas (\(Cl_2\)), two chlorine atoms each share one pair of electrons, forming a single covalent bond. Covalent compounds include small simple molecules (like \(H_2O\), \(CO_2\), \(CH_4\)) and enormous giant covalent structures (like diamond, graphite and silicon dioxide).
Dot-and-cross diagrams show how electrons are arranged in a bond. Use dots for the electrons from one atom and crosses for the electrons from the other. For ionic compounds, show the electron transfer and put each ion in square brackets with its charge. For covalent compounds, show the shared pairs in the overlap between the two atoms. These diagrams are a useful model but have limitations: they do not show the three-dimensional shape of molecules, the relative sizes of atoms, or the distances between atoms. Ball-and-stick models and space-filling models address some of these limitations but introduce others. No single model is perfect — each emphasises different aspects of bonding.
Metallic Bonding
In metals, the atoms are arranged in a regular structure. The outer-shell electrons are delocalised — they are free to move through the whole structure rather than being attached to individual atoms. This leaves behind a lattice of positive metal ions surrounded by a “sea” of delocalised electrons. The strong electrostatic attraction between the positive ions and the negative delocalised electrons is what holds the metal together.
Types of Structure
Substances can have different types of structure, which affects their properties:
Giant ionic lattices (e.g. NaCl) — regular arrangement of alternating positive and negative ions, held by strong ionic bonds in all directions. Simple molecular structures (e.g. \(H_2O\), \(CO_2\)) — small molecules with strong covalent bonds within molecules but weak intermolecular forces between molecules. Giant covalent structures (e.g. diamond, graphite) — huge networks of atoms all bonded by strong covalent bonds. Polymers — very long chain molecules made of repeating units joined by covalent bonds. Metallic structures — regular lattice of positive ions in a sea of delocalised electrons.
Chemical bonding is always about electrostatic forces: ionic = attraction between opposite ions; covalent = shared electron pairs attracting both nuclei; metallic = positive ions attracted to delocalised electrons.
The Periodic Table helps predict bonding type. Elements on the left (metals) tend to lose electrons. Elements on the right (non-metals) tend to gain or share electrons. An element’s position in the table — its group and period — tells you its electronic configuration, which determines how it bonds. Group 1 elements have one outer electron and form +1 ions; Group 7 elements have seven outer electrons and gain one to form −1 ions. Elements in Group 0 have full outer shells and do not readily bond.