1. Why Do Atoms Form Bonds?
Atoms form chemical bonds through interactions involving their electrons and nuclei. Bond formation often lowers the total energy of a system, making the bonded arrangement more stable.
These are electrons in the outermost occupied shell. They play a major role in bonding and chemical reactions.
The octet rule is a useful introductory guideline: many main-group atoms tend to form bonds that give them eight electrons in their valence shell. It has important exceptions, including hydrogen, electron-deficient compounds, and expanded-valence descriptions.
2. Ionic Bonding
Ionic bonding involves electrostatic attraction between oppositely charged ions. Ionic compounds commonly form when electrons are transferred between atoms, producing cations and anions.
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl
Solid ionic compounds typically form crystal lattices rather than separate pairs of ions. They often have high melting points and conduct electricity when molten or dissolved in water, provided mobile ions are present.
3. Covalent Bonding
A covalent bond forms when atoms share one or more pairs of electrons. Covalent bonds are common between nonmetal atoms.
Single bond
One shared pair of electrons. Example: H—H.
Double bond
Two shared pairs of electrons. Example: O=O.
Triple bond
Three shared pairs of electrons. Example: N≡N.
Coordinate covalent bond
A shared pair is initially donated by one atom. Once formed, it behaves as a covalent bond.
Lewis structures represent valence electrons as dots and shared electron pairs as lines. They help us count electrons and predict likely bonding patterns.
4. Metallic Bonding
In the simplified metallic-bond model, positive metal centres are held together by attraction to delocalised electrons that can move throughout the structure.
Their mobile electrons can respond to an applied electric field, allowing electric charge to flow.
Metallic bonding also helps explain why many metals can be hammered into sheets or drawn into wires without immediately breaking apart.
5. Comparing Bond Types
| Bond type | Key feature | Example |
|---|---|---|
| Ionic | Attraction between oppositely charged ions | NaCl |
| Covalent | Shared electron pairs | H₂O |
| Metallic | Metal centres and delocalised electrons | Cu |
These are useful models, but bonding can have both ionic and covalent character. Real substances do not always fit neatly into a single category.
6. Molecular Shape
Electron pairs repel one another. The valence-shell electron-pair repulsion (VSEPR) model predicts that electron groups arrange themselves to reduce repulsion.
Two electron groups around the central carbon produce a linear molecular shape, with a bond angle of approximately 180°.
Three bonding groups and no lone pairs on the central atom give a planar shape, with angles of approximately 120°.
Four bonding groups around carbon produce a tetrahedral shape, with angles of approximately 109.5°.
Two bonding pairs and two lone pairs around oxygen give water a bent shape, with a bond angle of approximately 104.5°.
7. Bond Polarity and Molecular Polarity
When bonded atoms attract shared electrons unequally, the bond can have partial positive and partial negative ends. This is called a polar covalent bond.
Molecular polarity depends on both bond polarities and the three-dimensional arrangement of the bonds. For example, individual C=O bonds in CO₂ are polar, but their dipoles cancel because the molecule is linear.
8. Intermolecular Forces
Intermolecular forces are attractions between particles. They are distinct from the covalent or ionic bonds within a substance, and they influence properties such as boiling point, melting point, and viscosity.
Arise from temporary fluctuations in electron distributions and occur between all atoms and molecules.
Attractions between the permanent dipoles of polar molecules.
A particularly important attraction when hydrogen bonded to nitrogen, oxygen, or fluorine interacts with a suitable electron-rich site on another particle.
Key Ideas
Valence electrons are central to the way atoms combine.
Bond type, molecular shape, and intermolecular forces help explain the behaviour of substances.
Lewis structures and VSEPR are useful introductory tools, but they simplify the underlying quantum-mechanical picture.