Bonding, structure & states

An element contains one type of atom; a compound has different atoms chemically bonded; a mixture is not chemically bonded and can be separated physically. Atoms react to gain a full outer shell, the stable arrangement of a noble gas. The type of bond depends on the atoms involved.

Ionic bonds form when a metal transfers electrons to a non-metal, making oppositely charged ions that attract. Predict charges from the Group: Group 1 → $1+$, Group 2 → $2+$, Al → $3+$, Group 16 → $2-$, Group 17 → $1-$. Know compound ions such as ${\text{CO}}_3^{2-}$ and ${\text{OH}}^{-}$. When a metal has more than one charge, a Roman numeral shows it, e.g. iron(III) chloride is ${\text{FeCl}}_3$.

Ionic compounds form giant lattices: high melting points, and they conduct electricity only when molten or dissolved (ions free to move).

A covalent bond is a shared pair of electrons between non-metals. Substances are either small molecules (water, ammonia, methane), with low melting points and no conductivity, or giant covalent structures (diamond, graphite, silicon dioxide) with very high melting points.

Metals are a giant lattice of positive ions in a sea of delocalised electrons. The free electrons explain why metals conduct electricity and heat, and the strong attraction gives high melting points.

Between molecules there are weak intermolecular forces; melting and boiling a molecular substance overcomes these (not the covalent bonds), so it happens at low temperatures. Always relate structure and bonding to properties such as melting point and conductivity.

In a solid particles are packed and vibrate in place; in a liquid they touch but move; in a gas they are far apart and fast. Changes of state (melting, boiling, freezing, condensing) change the spacing and movement; the energy needed depends on the strength of the bonding/intermolecular forces.

• Saying ionic solids conduct: they only conduct when molten or in solution.
• Thinking boiling a molecular substance breaks covalent bonds (it breaks the weaker intermolecular forces).
• Not balancing ion charges when writing a formula.