Chemical Bonding
Why Atoms Form Bonds
Atoms bond because a full outer shell of electrons is a stable, low-energy arrangement — the same arrangement the noble gases (Group 0) already have. Almost every other atom has an incomplete outer shell, so it will react and bond to reach that stable state.
There are three main ways atoms achieve a full outer shell:
- Ionic bonding — electrons are transferred
- Covalent bonding — electrons are shared
- Metallic bonding — electrons are delocalised (pooled)
The type of bond depends on which elements are involved.
Ionic Bonding
An ionic bond forms between a metal and a non-metal. The metal atom loses one or more outer electrons, and the non-metal atom gains them.
- Metals lose electrons to form positive ions (cations)
- Non-metals gain electrons to form negative ions (anions)
- The oppositely charged ions attract each other strongly — this attraction is the ionic bond
Example — sodium chloride (NaCl):
- Sodium (Na) has 1 outer electron. It loses it to become Na+.
- Chlorine (Cl) has 7 outer electrons. It gains one to become Cl−.
- Na+ and Cl− attract each other and build up into a giant ionic lattice.
Properties of Ionic Compounds
| Property | Explanation |
|---|---|
| High melting and boiling points | Many strong bonds hold the lattice together — a lot of energy is needed to break them |
| Conduct electricity when molten or dissolved | Ions become free to move and carry charge |
| Do not conduct as solids | Ions are locked in fixed positions |
| Often soluble in water | Water molecules can pull ions out of the lattice |
| Hard but brittle | Layers can shift, bringing like charges together, which cracks the crystal |
Covalent Bonding
A covalent bond forms between non-metal atoms that share pairs of electrons. Each shared pair counts towards a full outer shell for both atoms.
- A single bond shares one pair of electrons (e.g., H–H in hydrogen)
- A double bond shares two pairs (e.g., O=O in oxygen)
- A triple bond shares three pairs (e.g., N≡N in nitrogen)
Examples:
- Water (H₂O) — each hydrogen shares a pair with oxygen
- Carbon dioxide (CO₂) — two double bonds between carbon and oxygen
- Methane (CH₄) — carbon shares a pair with each of four hydrogens
Simple Molecules vs Giant Covalent Structures
Most covalent substances form small molecules with weak forces between the molecules:
- Low melting and boiling points (little energy needed to separate molecules)
- Usually do not conduct electricity (no free charges)
- Examples: water, oxygen, carbon dioxide
A few form giant covalent structures where huge numbers of atoms are bonded in a continuous network:
- Very high melting points
- Examples: diamond and graphite (both pure carbon), silicon dioxide (sand)
| Structure | Example | Melting point | Conducts electricity? |
|---|---|---|---|
| Simple molecule | Water | Low | No |
| Giant covalent | Diamond | Very high | No |
| Giant covalent | Graphite | Very high | Yes (free electrons) |
Metallic Bonding
In a metal, the outer electrons leave the individual atoms and become delocalised — they are free to move throughout the whole structure. This leaves a regular arrangement of positive metal ions surrounded by a "sea" of shared electrons.
The attraction between the positive ions and the sea of electrons is the metallic bond.
This structure explains why metals:
- Conduct electricity and heat — delocalised electrons carry charge and energy
- Are malleable and ductile — layers of ions slide over each other without breaking the bond, so metals bend and stretch instead of shattering
- Have high melting points — metallic bonds are strong
Comparing the Three Bond Types
| Feature | Ionic | Covalent | Metallic |
|---|---|---|---|
| Elements involved | Metal + non-metal | Non-metal + non-metal | Metal (+ metal) |
| What happens to electrons | Transferred | Shared | Delocalised |
| Particles formed | Ions | Molecules or networks | Ions + electron sea |
| Conducts electricity | When molten/dissolved | Usually not | Yes |
| Melting point | High | Low (simple) / high (giant) | Usually high |
Key Terms
- Ion — a charged atom that has lost or gained electrons
- Cation — a positive ion (has lost electrons)
- Anion — a negative ion (has gained electrons)
- Ionic bond — attraction between oppositely charged ions
- Covalent bond — a shared pair of electrons between atoms
- Metallic bond — attraction between positive metal ions and delocalised electrons
- Delocalised electron — an electron free to move through a whole structure
- Lattice — a regular, repeating arrangement of particles
Common Mistakes
- Saying ionic compounds conduct electricity as solids — they only conduct when molten or dissolved, once the ions are free to move
- Confusing "sharing" (covalent) with "transferring" (ionic) electrons
- Thinking a double bond shares two electrons — it shares two pairs (four electrons)
- Describing metals as having covalent bonds — metals use metallic bonding with delocalised electrons
Tips and Tricks
- Metal + non-metal → ionic; non-metal + non-metal → covalent; metal only → metallic
- "Ionic = I owe you" (electrons are given away); "Covalent = Co-operate" (electrons are shared)
- If it conducts when molten but not solid, it is almost certainly ionic
- High melting point + conducts as a solid → think metal