Atoms bond together to form molecules and compounds. The type of bond formed determines the properties of the substance. In this chapter you will explore covalent bonding, draw Lewis structures, predict molecular shapes using VSEPR theory, and explain why some molecules are polar and others are not.
6.1 Covalent Bonds and Lewis Structures
Definition
Covalent bond
A covalent bond is a form of chemical bond where pairs of electrons are shared between atoms to form molecules. Each atom contributes one electron to the shared pair (except in dative covalent bonds).
Definition
Dative covalent bond
A dative covalent bond (coordinate bond) is a type of covalent bond in which both electrons in the shared pair are donated by the same atom. This atom is called the donor; the atom receiving the pair is the acceptor. Example: NH₃ donates a lone pair to H⁺ to form NH₄⁺.
Atoms form covalent bonds to achieve a full outer electron shell (usually 8 electrons — the octet rule; 2 electrons for hydrogen). In a normal covalent bond, EACH atom contributes ONE electron to the shared pair. In a DATIVE (coordinate) covalent bond, ONE atom contributes BOTH electrons. Once formed, a dative covalent bond is indistinguishable from a normal covalent bond in terms of bond strength and length.
LEWIS STRUCTURES show all valence electrons. Shared pairs appear between atoms (bonding pairs); unshared pairs appear on individual atoms (lone pairs). Rules: 1. Count total valence electrons of all atoms. 2. Connect atoms with single bonds (using one pair each). 3. Complete the octets of outer atoms first. 4. Place remaining electrons on the central atom. 5. Form double or triple bonds if the central atom is short of an octet.
Watch Out
COMMON MISTAKE — Do NOT forget lone pairs when drawing Lewis structures. All valence electrons must be shown. For molecules like NH₃, the nitrogen has 1 lone pair in addition to 3 N–H bonding pairs. Omitting lone pairs leads to incorrect VSEPR shapes.
Worked Example
Draw the Lewis structure of carbon dioxide (CO₂). How many lone pairs and bonding pairs does each oxygen atom have?
Given
- C: 4 valence electrons
- Each O: 6 valence electrons
- Total valence electrons = 4 + 2(6) = 16
Find
Lewis structure; lone pairs and bonding pairs on each O
Solution
- 1Connect O–C–O with single bonds (uses 4 electrons): 12 remaining.
- 2Complete octets on both oxygens: add 3 lone pairs to each O (uses 12 electrons) → 0 remaining. But now C has only 4 electrons (from 2 single bonds) — short of octet.
- 3Convert one lone pair from each O into a bonding pair with C: this gives C=O=C (two double bonds).
- 4Final structure: O=C=O. Each O has 2 lone pairs and 2 bonding pairs (4 electron groups total).
Practice Question
Draw the Lewis structure for ammonia (NH₃). (a) How many bonding pairs does nitrogen form? (b) How many lone pairs does nitrogen have? (c) What is the total number of valence electrons in NH₃?
(6 marks)
6.2 VSEPR Theory and Molecular Shape
Definition
VSEPR
VSEPR (Valence Shell Electron Pair Repulsion) theory states that the shape of a molecule is determined by the repulsion between all electron pairs (bonding and lone pairs) in the valence shell of the central atom. Electron pairs arrange themselves to be as far apart as possible in order to minimise repulsion.
VSEPR THEORY states that electron pairs (both bonding and lone pairs) around the central atom repel each other and arrange themselves to be as far apart as possible. The SHAPE of the molecule is determined by the positions of the ATOMS, not the electron pairs. Lone pairs occupy more space than bonding pairs and cause greater repulsion, compressing bond angles.
SUMMARY OF VSEPR SHAPES: • 2 bonding pairs, 0 lone pairs → LINEAR (180°). Example: CO₂, BeCl₂. • 3 bonding pairs, 0 lone pairs → TRIGONAL PLANAR (120°). Example: BF₃, SO₃. • 4 bonding pairs, 0 lone pairs → TETRAHEDRAL (109,5°). Example: CH₄. • 3 bonding pairs, 1 lone pair → TRIGONAL PYRAMIDAL (<107°). Example: NH₃. • 2 bonding pairs, 2 lone pairs → BENT / V-SHAPE (<104,5°). Example: H₂O.
Exam Tip
EXAM TIP — Lone pairs repel more strongly than bonding pairs. The repulsion order is: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. This is why H₂O (2 lone pairs) has a smaller bond angle (104,5°) than NH₃ (1 lone pair, 107°), which is smaller than CH₄ (0 lone pairs, 109,5°).
Molecular Shapes — VSEPR Summary
| Property | Electron pairs around central atom | Shape (name and bond angle) |
|---|---|---|
| 2 BP, 0 LP | CO₂, BeCl₂ | Linear — 180° |
| 3 BP, 0 LP | BF₃, SO₃ | Trigonal planar — 120° |
| 4 BP, 0 LP | CH₄, CCl₄ | Tetrahedral — 109,5° |
| 3 BP, 1 LP | NH₃, PCl₃ | Trigonal pyramidal — ~107° |
| 2 BP, 2 LP | H₂O, H₂S | Bent / V-shape — ~104,5° |
Worked Example
Predict the shape and bond angle of the methane molecule (CH₄) using VSEPR theory.
Given
- Central atom: C (4 valence electrons)
- 4 H atoms each contribute 1 electron
- Lewis structure: C surrounded by 4 single bonds, no lone pairs
Find
Shape and bond angle of CH₄
Solution
- 1Count electron pairs around C: 4 bonding pairs, 0 lone pairs.
- 2By VSEPR, 4 electron pairs arrange to be as far apart as possible → tetrahedral arrangement.
- 3All 4 electron pairs are bonding pairs → molecular shape = TETRAHEDRAL.
- 4Bond angle = 109,5°.
Practice Question
Use VSEPR theory to predict the shape and approximate bond angle of (a) H₂O and (b) BF₃. For each molecule, state the number of bonding pairs and lone pairs on the central atom.
(8 marks)
6.3 Electronegativity, Bond Polarity and Bond Energy
Definition
Electronegativity
Electronegativity is a chemical property which describes the power of an atom in a molecule to attract the shared electron pair towards itself. The greater the electronegativity difference between two bonded atoms, the more polar the bond. Fluorine is the most electronegative element (χ = 4.0).
ELECTRONEGATIVITY TRENDS in the periodic table: • Increases across a period from left to right (more protons → stronger pull on electrons). • Decreases down a group (electrons are further from the nucleus — shielded by more electron shells). • Fluorine (F) is the most electronegative element (χ = 4,0 on the Pauling scale). • Noble gases are excluded as they do not typically form bonds.
BOND POLARITY: If two bonded atoms have DIFFERENT electronegativities, the shared electrons are pulled toward the more electronegative atom. The bond has a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative atom. This is a POLAR COVALENT BOND. Rule of thumb: • Δχ < 0,4 → non-polar covalent bond • 0,4 ≤ Δχ < 1,7 → polar covalent bond • Δχ ≥ 1,7 → ionic bond
Definition
Polar molecules
A polar molecule is one in which there is an uneven distribution of electron density so that one part of the molecule has a partial negative charge (δ−) and another part has a partial positive charge (δ+). Polar molecules have a net dipole moment. Example: H₂O, HCl, NH₃.
Definition
Non-polar molecules
A non-polar molecule is one in which there is an even distribution of electron density so that no permanent dipole moment exists. This occurs when: (1) all bonds are non-polar (e.g., H₂, Cl₂), or (2) the molecule is symmetrical so that bond dipoles cancel (e.g., CO₂ linear, CCl₄ tetrahedral).
MOLECULAR POLARITY vs BOND POLARITY: A molecule can have polar bonds yet still be NON-POLAR overall, if the molecule is SYMMETRICAL and the bond dipoles cancel. Example: CO₂ has two polar C=O bonds, but the molecule is linear and the two dipoles point in opposite directions — they cancel → non-polar molecule. H₂O has two polar O–H bonds, but the molecule is BENT → the dipoles do NOT cancel → polar molecule.
Definition
Bond energy
The amount of energy that must be added to break the bond that has formed.
BOND LENGTH AND BOND ORDER: A single bond is the longest and weakest covalent bond. A double bond is shorter and stronger. A triple bond is the shortest and strongest. This is because more shared electron pairs draw the two nuclei closer together and hold them more tightly. Example for carbon-carbon bonds: • C–C (single): bond length ≈ 154 pm, bond energy ≈ 347 kJ·mol⁻¹ • C=C (double): bond length ≈ 134 pm, bond energy ≈ 614 kJ·mol⁻¹ • C≡C (triple): bond length ≈ 120 pm, bond energy ≈ 839 kJ·mol⁻¹
Exam Tip
EXAM TIP — To determine if a molecule is polar: (1) Check if individual bonds are polar (Δχ ≠ 0). (2) Use VSEPR to determine the shape. (3) Check if bond dipoles are symmetrically arranged and cancel. If they cancel → non-polar. If they don't cancel → polar.
Polar vs Non-polar Molecules
| Property | Polar molecule | Non-polar molecule |
|---|---|---|
| Bond dipoles | Do NOT cancel — asymmetric arrangement | Cancel — symmetric arrangement |
| Net dipole moment | Non-zero | Zero |
| Examples | H₂O, NH₃, HCl, SO₂ | CO₂, CCl₄, BF₃, H₂ |
| Solubility in water | Generally soluble (like dissolves like) | Generally insoluble in water |
Practice Question
Explain why CO₂ is a non-polar molecule despite having polar C=O bonds.
(5 marks)