Bonding: Ionic, Covalent and Metallic
Why Atoms Form Bonds
Atoms are most stable when they have a full outer shell of electrons. Most atoms do not naturally achieve this on their own, so they combine with other atoms through chemical bonding. All chemical bonds involve electrostatic forces — attractions between positive and negative charges. The type of bond that forms depends on whether the atoms are metals or non-metals and on how many outer-shell electrons they have. There are three main types: ionic bonding, covalent bonding and metallic bonding.
Ionic Bonding
Ionic bonding forms between a metal and a non-metal. The metal atom transfers one or more electrons to the non-metal atom. The metal becomes a positively charged ion (cation) because it now has more protons than electrons. The non-metal becomes a negatively charged ion (anion) for the opposite reason. The resulting opposite charges attract each other strongly: this electrostatic force of attraction is the ionic bond.
In sodium chloride (NaCl): sodium (2, 8, 1) loses one electron to become \(\text{Na}^+\) (2, 8), and chlorine (2, 8, 7) gains that electron to become \(\text{Cl}^-\) (2, 8, 8). Both ions now have full outer shells. In magnesium oxide (MgO): magnesium (2, 8, 2) loses two electrons to form \(\text{Mg}^{2+}\), and oxygen (2, 6) gains two to form \(\text{O}^{2-}\). In any ionic compound, billions of ions pack together in a regular three-dimensional arrangement called a giant ionic lattice, where every ion is surrounded by oppositely charged neighbours in all directions.
In ionic bonding, electrons are transferred from metal to non-metal. Both ions achieve full outer shells. Ionic compounds form giant lattices held together by strong electrostatic attraction between opposite charges.
Covalent Bonding
Covalent bonding forms between non-metal atoms. Instead of transferring electrons, the atoms share pairs of electrons. Each shared pair counts as one covalent bond. By sharing, each atom achieves a full outer shell without losing or gaining electrons overall. A single shared pair is a single bond. Two shared pairs form a double bond (e.g. \(O_2\), \(CO_2\)). Three shared pairs form a triple bond (e.g. \(N_2\)).
In water (\(H_2O\)): the oxygen atom shares one pair of electrons with each of two hydrogen atoms, forming two single covalent bonds. In chlorine gas (\(Cl_2\)): two chlorine atoms share one pair of electrons. In carbon dioxide (\(CO_2\)): each oxygen forms a double bond with the central carbon atom. Covalent compounds include simple small molecules (like \(H_2O\), \(CO_2\), \(CH_4\)) and very large networks called giant covalent structures (like diamond and silicon dioxide).
Dot-and-cross diagrams are the standard way to represent electron arrangement in bonds. For covalent bonds, use dots for electrons from one atom and crosses for electrons from the other; draw the outer shells overlapping and place the shared pairs in the overlap. For ionic bonds, show the electron being transferred and enclose each ion in square brackets with its charge written outside. These diagrams are models — they show electron positions but not the three-dimensional shape, bond angles, or relative atomic sizes. Different models (ball-and-stick, space-filling) address these limitations but introduce others. No single model captures everything about a bond; choosing the right one depends on what you need to show.
In covalent bonding, non-metal atoms share electron pairs. Each shared pair is one bond. Both atoms count the shared electrons as their own, giving both a full outer shell.
Metallic Bonding
In a metal, atoms are arranged in a regular lattice structure. The outer-shell electrons are not bound to any individual atom — they are delocalised and free to move throughout the entire metallic structure. This creates a “sea” of mobile electrons surrounding a lattice of positive metal ions. The strong electrostatic attraction between the positive ions and the negative delocalised electrons holds the structure together. This is metallic bonding.
The delocalised electrons explain many properties of metals: they conduct electricity because electrons can move freely to carry charge; they conduct heat well because electrons can transfer energy; metals are malleable and ductile because the positive ions can slide past each other without breaking the bonds with the electron sea. The more delocalised electrons per atom (e.g. aluminium gives three, copper gives one), the stronger the metallic bonding and the higher the melting point.
Comparing Bond Types and Structures
The type of bond present in a substance determines its structure, which in turn determines its bulk properties. Understanding this link is essential:
Giant ionic lattice (e.g. NaCl, MgO) — high melting and boiling points because a large amount of energy is needed to overcome the many strong electrostatic attractions throughout the lattice. Conducts electricity when molten or dissolved in water (ions are free to move) but not when solid (ions are fixed in position).
Simple molecular structure (e.g. \(H_2O\), \(CO_2\), \(CH_4\)) — low melting and boiling points because the weak intermolecular forces between molecules require little energy to overcome. The covalent bonds within each molecule are strong, but they are not broken on melting. Does not conduct electricity (no free ions or electrons). The stronger the intermolecular forces (larger molecules), the higher the boiling point.
Giant covalent structure (e.g. diamond, graphite, silicon dioxide) — very high melting and boiling points because melting requires breaking many strong covalent bonds throughout the whole structure. Diamond does not conduct electricity (all electrons are used in bonding); graphite does conduct electricity because it has one delocalised electron per carbon atom.
Metallic structure — high melting and boiling points (especially for metals with more delocalised electrons per atom). Excellent electrical and thermal conductors due to delocalised electrons. Malleable and ductile because ions can slide.
The periodic table gives you the starting information. Metals on the left — metallic bonding. Non-metals on the right — covalent bonding between non-metals, ionic bonding when a metal and non-metal combine. Group 0 elements do not bond. From bond type, you can predict structure; from structure, you can predict properties. This logical chain — position in table → electron arrangement → bond type → structure → properties — is the engine of chemical prediction.