What this chapter is about
This chapter explains why and how atoms combine to form molecules and compounds. You have already studied atomic structure and the arrangement of electrons in shells and subshells. Now you learn the driving force behind bond formation: atoms tend to achieve a stable, low-energy configuration, often resembling the electron arrangement of noble gases.
The chapter covers three main types of chemical bonds—ionic, covalent and coordinate—and the theories that describe them: the octet rule, Lewis structures, Valence Bond Theory (VBT) and the concept of hybridisation. You also meet the Valence Shell Electron Pair Repulsion (VSEPR) model, which predicts molecular shapes, and the basics of Molecular Orbital Theory (MOT), which explains properties like bond order and magnetism.
After studying this chapter you should be able to draw Lewis dot structures, predict molecular geometry, describe hybridisation in common molecules, calculate bond order using MOT, and relate structure to physical properties such as polarity and bond strength.
Key ideas
- Octet rule: Atoms tend to gain, lose or share electrons until their valence shell has eight electrons, achieving noble-gas stability. Hydrogen is an exception; it needs only two electrons (duplet).
- Ionic bond: Formed when one atom transfers electrons to another, creating cations and anions held together by electrostatic attraction. Example: Na loses one electron to Cl, forming Na⁺Cl⁻.
- Covalent bond: Formed when two atoms share one or more pairs of electrons. A single bond involves one shared pair, a double bond two pairs, and a triple bond three pairs.
- Coordinate (dative) bond: A covalent bond where both shared electrons come from one atom. In NH₄⁺, the nitrogen of NH₃ donates its lone pair to H⁺.
- VSEPR theory: Electron pairs around a central atom repel each other and arrange themselves to minimise repulsion. Lone pairs repel more than bonding pairs, distorting bond angles.
- Hybridisation: Atomic orbitals of similar energy mix to form new, equivalent hybrid orbitals. sp gives linear geometry (180°), sp² gives trigonal planar (120°), sp³ gives tetrahedral (109.5°).
- Molecular Orbital Theory (MOT): Atomic orbitals combine to form bonding and antibonding molecular orbitals. Electrons fill these according to Aufbau, Pauli and Hund's rules. Bond order = (bonding electrons − antibonding electrons) / 2.
- Polarity: A bond is polar if the two atoms differ in electronegativity; a molecule is polar if the vector sum of bond dipoles is non-zero.
Formulas and facts to remember
1. Bond order (MOT) = (Nb − Na) / 2, where Nb = electrons in bonding MOs, Na = electrons in antibonding MOs. Higher bond order means stronger, shorter bond.
2. Formal charge = Valence electrons of free atom − [Non-bonding electrons + ½(Bonding electrons)].
3. Dipole moment μ = q × d, where q is charge magnitude and d is separation distance. SI unit is coulomb-metre; often given in debye (1 D ≈ 3.336 × 10⁻³⁰ C·m).
4. Hybridisation and geometry:
- sp → linear, 180° (BeCl₂, CO₂)
- sp² → trigonal planar, 120° (BF₃, C₂H₄)
- sp³ → tetrahedral, 109.5° (CH₄, NH₄⁺)
- sp³d → trigonal bipyramidal, 90° and 120° (PCl₅)
- sp³d² → octahedral, 90° (SF₆)
5. VSEPR repulsion order: Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
6. Sigma (σ) bond: Formed by head-on overlap of orbitals along the internuclear axis.
7. Pi (π) bond: Formed by sideways overlap of p orbitals above and below the internuclear axis.
8. Resonance: When a single Lewis structure cannot represent actual electron distribution, we draw multiple structures; the real molecule is a hybrid.
Worked examples
### Example 1: Drawing the Lewis structure of CO₂
Problem: Draw the Lewis structure of carbon dioxide and state its geometry.
Solution: 1. Count valence electrons: C has 4, each O has 6. Total = 4 + 6 + 6 = 16 electrons. 2. Place C in the centre (less electronegative). Attach O atoms with single bonds: O–C–O. This uses 4 electrons; 12 remain. 3. Distribute remaining electrons to give each atom an octet. With single bonds, C has only 4 electrons. Form double bonds: O=C=O. 4. Check: Each O has 4 non-bonding electrons plus 4 bonding electrons = 8. C has 8 bonding electrons = 8. Octets satisfied. 5. No lone pairs on C; two bonding domains. VSEPR predicts linear geometry, bond angle 180°.
Answer: O=C=O; linear shape.
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### Example 2: Determining hybridisation and shape of H₂O
Problem: Predict the hybridisation and shape of a water molecule.
Solution: 1. Oxygen has 6 valence electrons. Two are used in bonds with two H atoms; four remain as two lone pairs. 2. Total electron domains around O = 2 bond pairs + 2 lone pairs = 4. 3. Four domains require sp³ hybridisation (tetrahedral arrangement of orbitals). 4. Because two positions are lone pairs, the molecular shape is bent (angular), not tetrahedral. 5. Lone pair–bond pair repulsion compresses the H–O–H angle below 109.5° to about 104.5°.
Answer: sp³ hybridised; bent shape with bond angle ≈ 104.5°.
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### Example 3: Bond order of O₂ using MOT
Problem: Calculate the bond order of O₂ and comment on its magnetic behaviour.
Solution: 1. Each O atom has 8 electrons; O₂ has 16 electrons. 2. Fill molecular orbitals in order (for homonuclear diatomics of second period from O onwards): σ1s² < σ*1s² < σ2s² < σ*2s² < σ2p² < π2p⁴ < π*2p². 3. Bonding electrons (in σ1s, σ2s, σ2p, π2p): 2 + 2 + 2 + 4 = 10. 4. Antibonding electrons (in σ*1s, σ*2s, π*2p): 2 + 2 + 2 = 6. 5. Bond order = (10 − 6) / 2 = 2. 6. Two electrons in π*2p orbitals are unpaired (one in each degenerate orbital by Hund's rule), so O₂ is paramagnetic.
Answer: Bond order = 2 (a double bond); O₂ is paramagnetic.
Common mistakes
- Thinking all sp³ molecules are tetrahedral → Shape depends on how many positions are lone pairs; NH₃ is sp³ but pyramidal, H₂O is sp³ but bent.
- Ignoring lone pairs when predicting geometry → Always count lone pairs; they occupy space and affect angles.
- Confusing molecular polarity with bond polarity → CO₂ has polar C=O bonds but is non-polar overall because dipoles cancel.
- Assuming higher bond order always means more bonds shown in Lewis structure → MOT bond order can be fractional (e.g., 1.5 in O₂⁺) even when Lewis structure shows integer bonds.
- Writing bond order as a negative number → If antibonding electrons exceed bonding electrons, bond order is zero or molecule does not exist; never report a negative bond order.
Quick revision
1. Octet rule: atoms share, gain or lose electrons to reach eight valence electrons.
2. VSEPR: electron pairs repel; shape is decided by bonding pairs; angles shrink when lone pairs are present.
3. Hybridisation = mixing of orbitals; sp³ (109.5°), sp² (120°), sp (180°).
4. Bond order (MOT) = (Nb − Na) / 2; higher bond order means stronger and shorter bond.
5. Polar molecule: net dipole ≠ 0; symmetric molecules often have zero net dipole.
6. Sigma bonds form by head-on overlap; pi bonds form by sideways overlap and exist in double/triple bonds.