Ionic, Covalent and Metallic Bonding
Why atoms bond
Atoms bond to achieve a full outer shell of electrons (a stable arrangement, like the noble gases). There are three types of chemical bond: ionic, covalent and metallic. Which one forms depends on whether the atoms are metals or non-metals.
Ionic bonding (metal + non-metal)
An ionic bond forms when a metal transfers electrons to a non-metal.
- The metal loses electrons to become a positive ion (cation).
- The non-metal gains electrons to become a negative ion (anion).
- The oppositely charged ions are held together by strong electrostatic forces of attraction.
Example: sodium chloride — Na loses 1 electron (Na⁺), Cl gains 1 electron (Cl⁻).
Ionic compounds form giant lattices and have:
- High melting/boiling points (strong forces need lots of energy to break).
- Conduct electricity when molten or dissolved (ions become free to move) — but not when solid.
Covalent bonding (non-metal + non-metal)
A covalent bond is a shared pair of electrons between two non-metal atoms. Each shared pair counts towards both atoms' outer shells.
- Simple molecular substances (e.g. H₂O, CO₂, O₂): low melting/boiling points because the weak forces between molecules (not the strong bonds within them) are easily broken. Usually don't conduct electricity (no free charges).
- Giant covalent structures (e.g. diamond, graphite, silicon dioxide): billions of atoms covalently bonded → very high melting points.
Metallic bonding (metal + metal)
Metals consist of a lattice of positive ions in a "sea" of delocalised (free) electrons. The attraction between the positive ions and the delocalised electrons is the metallic bond.
This explains metal properties:
- Good conductors of electricity and heat (delocalised electrons carry charge/energy).
- Malleable (layers of ions can slide over each other).
- Generally high melting points.
Summary table
| Bond | Between | How | Conducts? |
|---|---|---|---|
| Ionic | metal + non-metal | transfer electrons → ions | When molten/dissolved |
| Covalent | non-metal + non-metal | share electrons | Usually no (graphite is an exception) |
| Metallic | metal + metal | sea of delocalised electrons | Yes (solid and molten) |
Worked example
Why does sodium chloride conduct electricity when molten but not when solid?
- In solid NaCl the ions are locked in the lattice and can't move. When molten (or dissolved), the ions are free to move and carry charge. ✓
Common mistakes
- Saying ionic compounds conduct as a solid — only when molten or dissolved.
- Confusing the weak forces between molecules with the strong covalent bonds within them (it's the between-forces that break on melting).
- Mixing up which bond forms — check metal/non-metal combination.
Exam tips
- Link properties to structure and bonding every time (e.g. high melting point ↔ strong forces).
- For simple molecules, stress it's the intermolecular forces that are weak, not the covalent bonds.
- Explain metallic conductivity/malleability using delocalised electrons and sliding layers.
Key facts to remember
- Ionic = electron transfer (metal→non-metal), giant lattice, conducts when molten/dissolved.
- Covalent = shared electrons (non-metals); simple molecules have low mp (weak intermolecular forces); giant covalent have very high mp.
- Metallic = positive ions + sea of delocalised electrons → conductive and malleable.