Why atoms bond
Most atoms are more stable when their outer shell is full. Noble gases already have full shells, so they rarely react. Other atoms gain, lose or share electrons to achieve a noble-gas configuration. This tendency is called the octet rule: atoms tend to end up with eight valence electrons (hydrogen and helium are satisfied with two).
A chemical bond is the attraction that holds atoms together.
Types of bonds
| Ionic bond | Covalent bond | Metallic bond | |
|---|---|---|---|
| Between | Metal + non-metal | Non-metal + non-metal | Metal + metal |
| Electrons | Transferred | Shared | Pooled in a "sea" |
| Particles | Ions in a lattice | Molecules | Cations in electron sea |
| Melting point | High | Usually low | Usually high |
| Conducts electricity | When melted or dissolved | Usually not | Yes, as a solid |
| Example | NaCl | H₂O, CO₂ | Copper, iron |
Ionic bonding
Sodium has 1 valence electron; chlorine has 7. Sodium transfers its electron to chlorine:
- Na → Na⁺ (now has the electron arrangement of neon)
- Cl → Cl⁻ (now has the electron arrangement of argon)
The oppositely charged ions attract strongly, forming a crystal lattice. That’s why salt forms cubic crystals, has a high melting point (801 °C) and is brittle.
Writing ionic formulas
The total charge must equal zero.
| Ions | Formula |
|---|---|
| Na⁺ and Cl⁻ | NaCl |
| Mg²⁺ and Cl⁻ | MgCl₂ |
| Al³⁺ and O²⁻ | Al₂O₃ |
| Ca²⁺ and OH⁻ | Ca(OH)₂ |
Tip: the "criss-cross" method swaps the charge numbers to become subscripts, then reduces if possible.
Covalent bonding
Non-metals share electrons so each atom gets a full outer shell.
- Single bond: one shared pair (H–H, H–Cl)
- Double bond: two shared pairs (O=O, the C=O bonds in CO₂)
- Triple bond: three shared pairs (N≡N), very strong
Lewis structures
A Lewis structure shows valence electrons as dots or lines.
Steps:
- Count total valence electrons.
- Place the least electronegative atom in the centre (never hydrogen).
- Connect atoms with single bonds.
- Complete octets on outer atoms with lone pairs.
- Use multiple bonds if the central atom lacks an octet.
Water: O in the centre, two H atoms bonded, and two lone pairs on O.
Polar and non-polar bonds
When atoms with different electronegativities share electrons, the more electronegative atom pulls them closer, creating partial charges (δ+ and δ−). The bond is polar.
| Electronegativity difference | Bond type |
|---|---|
| 0 to about 0.4 | Non-polar covalent |
| about 0.4 to 1.7 | Polar covalent |
| above about 1.7 | Mostly ionic |
Molecular shape (VSEPR)
Electron pairs repel each other and spread out as far as possible.
| Molecule | Shape | Bond angle | Polar? |
|---|---|---|---|
| CO₂ | Linear | 180° | No (polar bonds cancel) |
| H₂O | Bent | 104.5° | Yes |
| NH₃ | Trigonal pyramidal | 107° | Yes |
| CH₄ | Tetrahedral | 109.5° | No |
Shape determines polarity, which determines many properties, like whether a substance dissolves in water ("like dissolves like").
Intermolecular forces
Forces between molecules are weaker than bonds but control boiling points and states.
- London dispersion forces: weakest; present in all molecules
- Dipole–dipole forces: between polar molecules
- Hydrogen bonds: strong attractions when H is bonded to N, O or F
Hydrogen bonding explains why water boils at 100 °C even though it’s a small molecule, why ice floats, and how DNA’s two strands stay together.
Metallic bonding
Metal atoms release their valence electrons into a shared "sea." This explains why metals:
- conduct electricity (electrons move freely)
- are malleable and ductile (layers slide without breaking bonds)
- are shiny (free electrons reflect light)
Common mistakes
- Thinking ionic compounds conduct electricity as solids. Ions must be free to move.
- Thinking a molecule with polar bonds must be polar. Symmetry can cancel polarity (CO₂).
- Confusing bonds with intermolecular forces. Boiling water breaks intermolecular forces, not the O–H bonds.
Try it yourself
- Predict the bond type: KBr, CH₄, Fe, NH₃.
- Write the formula for the compound formed by Ca²⁺ and N³⁻.
- Draw the Lewis structure for CO₂.
- Explain why oil and water don’t mix.
(Answers: ionic, covalent, metallic, covalent; Ca₃N₂; O=C=O with two lone pairs on each O; water is polar and oil is non-polar)
Quick recap
- Atoms bond to reach stable electron arrangements.
- Ionic bonds transfer electrons; covalent bonds share them; metallic bonds pool them.
- Lewis structures and VSEPR predict molecular shapes.
- Unequal sharing makes polar bonds; shape decides if a molecule is polar.
- Intermolecular forces such as hydrogen bonds control boiling points and solubility.