Chemical Bonding

Why Do Atoms Bond?

Atoms form chemical bonds to achieve a more stable electron configuration — specifically, a full outer electron shell. This stable state is called the noble gas configuration. Most atoms have partially filled outer shells and are energetically "uncomfortable" — bonding with other atoms lowers their energy and increases stability.

Ionic Bonding

Ionic bonding occurs between a metal and a non-metal. The metal atom has few valence electrons (usually 1–3) and loses them to the non-metal, which has many valence electrons (usually 5–7) and gains them. This transfer creates charged particles called ions:

• The metal loses electrons → becomes a positively charged cation (e.g. Na⁺, Mg²⁺, Ca²⁺)
• The non-metal gains electrons → becomes a negatively charged anion (e.g. Cl⁻, O²⁻, F⁻)

The oppositely charged ions attract each other strongly — this electrostatic attraction is the ionic bond. The ions arrange themselves into a regular, repeating 3D structure called an ionic lattice (or crystal lattice).

Properties of ionic compounds:

• High melting and boiling points — the electrostatic forces holding the lattice together are very strong and require a lot of energy to overcome.
• Hard but brittle — the lattice is rigid, but if a layer is displaced by a blow, like-charged ions align and repel, causing the crystal to split.
• Conduct electricity when dissolved in water or melted — ions are free to move and carry charge, but solid ionic compounds cannot conduct (ions are locked in place).
• Soluble in water (many ionic compounds) — water molecules are attracted to the ions and pull them into solution.

Covalent Bonding

Covalent bonding occurs between two non-metals. Instead of transferring electrons, the atoms share pairs of electrons. Each shared pair of electrons counts toward the full outer shell of both atoms simultaneously.

• Single bond: 1 shared pair of electrons (e.g. H–H in H₂, H–Cl in HCl)
• Double bond: 2 shared pairs of electrons (e.g. O=O in O₂, O=C=O in CO₂)
• Triple bond: 3 shared pairs of electrons (e.g. N≡N in N₂)

Properties of covalent (molecular) compounds:

• Lower melting and boiling points — the forces between separate molecules (intermolecular forces) are weak; only a little energy is needed to separate them.
• Do not conduct electricity — molecules have no overall charge and no free ions or electrons to carry current.
• Can exist as gases, liquids, or solids at room temperature (H₂O is liquid; CO₂ is gas; wax is solid).
• Generally insoluble in water (non-polar covalent compounds), though polar ones like HCl dissolve readily.

Metallic Bonding

In metallic bonding, metal atoms release their valence electrons into a shared pool — a "sea of delocalised electrons" — that flows freely throughout the structure. The remaining positive metal ions (cations) sit in a regular lattice, held in place by their attraction to the surrounding electron sea.

Properties explained by metallic bonding:

• Excellent electrical conductors — the delocalised electrons can move freely through the lattice when a voltage is applied.
• Excellent heat conductors — free electrons rapidly transfer kinetic energy through the metal.
• Malleable and ductile — layers of ions can slide past each other without breaking, because the electron sea is not directional.
• Lustrous (shiny) — free electrons reflect light at the surface.
• High melting points (generally) — strong attraction between positive ions and electron sea.

Comparing Bond Types

Lewis Dot Structures (Electron Dot Diagrams)

A Lewis dot structure shows the valence electrons of each atom as dots arranged around the element symbol. When atoms bond, shared electron pairs are shown between the two symbols, and lone (unshared) pairs are shown on the outer sides. This is a useful way to visualise how many bonds an atom forms and how electrons are distributed.

• H₂: H· ·H → H:H (1 shared pair = single bond)
• Cl₂: Each Cl has 7 valence electrons; they share 1 pair → :Cl:Cl: with 3 lone pairs on each Cl
• H₂O: O has 6 valence electrons; shares 1 pair with each H → H:O:H with 2 lone pairs on O