Revision notes · Bonding, structure, and the properties of matter
Chemical bonds, ionic, covalent and metallic
Chemical bonds4.2.1.1
Definition: There are three main types of strong chemical bond: ionic, covalent and metallic.
The type of bonding in a substance depends on the elements involved.
| Bond type | Between | Formed by |
|---|---|---|
| Ionic | A metal and a non-metal | Transfer of electrons, forming charged ions |
| Covalent | Two or more non-metals | Sharing of electron pairs |
| Metallic | Metal atoms | Delocalised electrons shared between metal ions |
⚠️ Common mistake: Trying to work out bonding type from the compound's state or appearance — always look at what elements are involved (metal + non-metal = ionic; non-metals only = covalent).
🧠 Remember: Metal + non-metal = ionic. Non-metal + non-metal = covalent. Metal + metal = metallic.
Ionic bonding4.2.1.2
Definition: Ionic bonding is the strong electrostatic attraction between oppositely charged ions, formed when electrons are transferred from a metal atom to a non-metal atom.
Metal atoms lose electrons to form positive ions; non-metal atoms gain those electrons to form negative ions.
- •Atoms transfer electrons to achieve a full outer shell, like the nearest noble gas.
- •A metal atom loses electrons to form a positively charged ion (cation).
- •A non-metal atom gains electrons to form a negatively charged ion (anion).
- •The oppositely charged ions are then strongly attracted to each other by electrostatic forces.
⚠️ Common mistake: Forgetting that both atoms end up with a full outer shell after ionic bonding — this is the whole reason the electron transfer happens.
🧠 Remember: Ionic bonding: metal gives, non-metal takes — then opposites attract.
The formula of an ionic compound is worked out by balancing the total positive and negative charge so it adds up to zero.
| Ion | Charge |
|---|---|
| Group 1 metal (e.g. Na⁺, K⁺) | 1+ |
| Group 2 metal (e.g. Mg²⁺, Ca²⁺) | 2+ |
| Group 6 non-metal (e.g. O²⁻) | 2− |
| Group 7 non-metal (e.g. Cl⁻) | 1− |
| Hydroxide, OH⁻ | 1− |
| Nitrate, NO₃⁻ | 1− |
| Carbonate, CO₃²⁻ | 2− |
| Sulfate, SO₄²⁻ | 2− |
- •If the ions' charges already balance 1:1, just write them together, e.g. Na⁺ + Cl⁻ → NaCl.
- •If they don't balance, use enough of each ion to make the total charge zero, e.g. Mg²⁺ needs two Cl⁻ ions → MgCl₂.
- •If more than one of a compound (polyatomic) ion like OH⁻, NO₃⁻, CO₃²⁻ or SO₄²⁻ is needed, put it in brackets with the number outside, e.g. Mg²⁺ with two OH⁻ ions → Mg(OH)₂.
- •For Al³⁺ and O²⁻, the charges don't cancel with one of each — cross-multiply the charges as subscripts (2 Al, 3 O) → Al₂O₃.
⚠️ Common mistake: Putting brackets around a single ion, or around a simple ion (like Cl⁻) instead of a compound ion — brackets are only needed when more than one of a compound ion (e.g. OH⁻, NO₃⁻) is present.
🧠 Remember: Balance the charges to zero — the number of each ion is whatever makes positive and negative charge cancel out.
Ionic compounds4.2.1.3
Ionic compounds form a regular, repeating 3D arrangement called a giant ionic lattice.
- •In a giant ionic lattice, each ion is surrounded by ions of the opposite charge, held together by strong electrostatic forces in all directions.
- •Ionic compounds have high melting and boiling points, because a lot of energy is needed to overcome the many strong bonds in the lattice.
- •Ionic compounds conduct electricity when molten or dissolved in water, because the ions are then free to move and carry charge — but not when solid, as the ions are held in fixed positions.
⚠️ Common mistake: Saying an ionic compound conducts electricity as a solid — the ions can't move in a solid lattice, so it only conducts when molten or in solution.
🧠 Remember: Ionic compounds conduct only when the ions can move: molten or dissolved.
Covalent bonding4.2.1.4
Definition: Covalent bonding is a strong bond formed when two non-metal atoms share a pair of electrons.
- •Each shared pair of electrons counts towards the outer shell of both atoms, helping each achieve a full outer shell.
- •Covalently bonded substances can form small molecules (e.g. H₂O, CO₂) or giant covalent structures (e.g. diamond).
- •Covalent bonds themselves are very strong, but in small molecules, the forces between separate molecules are weak.
⚠️ Common mistake: Confusing the strength of the covalent bond itself with the strength of the forces between molecules — the bond is strong, but the intermolecular forces are usually weak, which is why many covalent substances have low melting points.
🧠 Remember: Covalent bond = strong. Forces between small molecules = weak.
Metallic bonding4.2.1.5
Definition: Metallic bonding is the strong attraction between positive metal ions and a 'sea' of delocalised electrons that can move freely throughout the structure.
- •Metal atoms lose their outer-shell electrons, which become delocalised (free to move) throughout the metal's giant structure.
- •The metal atoms become positive ions, held together by strong electrostatic attraction to the shared delocalised electrons.
- •Because the delocalised electrons can move, metals are good conductors of electricity and heat.
- •The layers of ions can slide over each other, which is why metals are malleable and ductile.
⚠️ Common mistake: Thinking metallic bonding involves electron transfer like ionic bonding — in metals, electrons become delocalised (shared throughout the structure), not transferred to another atom.
🧠 Remember: Metallic bonding = ions in a sea of delocalised electrons — that's why metals conduct and bend.
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