Shapes of Molecules and Bond Polarity

Shapes of Molecules and Bond Polarity

VSEPR theory (Valence Shell Electron Pair Repulsion) states that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsion. This determines the shape and bond angles of molecules and ions.

Electron Pair Types

There are two types of electron pair:

• Bonding pairs (bp) — shared between two atoms in a covalent bond
• Lone pairs (lp) — non-bonding pairs that remain on the central atom

Lone pairs repel more strongly than bonding pairs because they are held closer to the central atom and spread over a larger volume. The order of repulsion strength is:

lp–lp > lp–bp > bp–bp

This means lone pairs compress bond angles below the ideal values.

Common Molecular Shapes

Electron pairs	Bonding pairs	Lone pairs	Shape	Bond angle	Example
2	2	0	Linear	180°	BeCl₂, CO₂
3	3	0	Trigonal planar	120°	BF₃, AlCl₃
3	2	1	Bent (V-shaped)	~117°	SO₂
4	4	0	Tetrahedral	109.5°	CH₄, NH₄⁺
4	3	1	Trigonal pyramidal	~107°	NH₃, PCl₃
4	2	2	Bent (V-shaped)	~104.5°	H₂O
5	5	0	Trigonal bipyramidal	90° & 120°	PCl₅
6	6	0	Octahedral	90°	SF₆
6	4	2	Square planar	90°	XeF₄

Predicting Shape — Step by Step

1. Draw the dot-and-cross diagram for the molecule

2. Count the total number of electron pairs (bonding + lone) around the central atom (treat a double or triple bond as one bonding region)

3. Identify how many are bonding regions and how many are lone pairs

4. Determine the shape using VSEPR principles

5. Adjust the bond angle for each lone pair (reduce by approximately 2–2.5° per lone pair)

Worked Example: ClF₃

Chlorine has 7 outer electrons. Three are used in bonds to F atoms, leaving two lone pairs.

• Total electron regions around Cl: 5 (3 bp + 2 lp)
• Base geometry: trigonal bipyramidal
• Lone pairs occupy equatorial positions (where they experience less repulsion)
• Result: T-shaped, bond angles approximately 87.5° (less than 90° due to lp compression)

Bond Polarity

A polar bond forms when two atoms with different electronegativities share electrons unequally. The more electronegative atom acquires a partial negative charge (δ⁻) and the less electronegative atom acquires a partial positive charge (δ⁺).

Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond. It increases across a period and decreases down a group (fluorine is the most electronegative element, value 4.0 on the Pauling scale).

Polar and Non-Polar Molecules

A molecule is polar (has a net dipole moment) only if:

1. It contains polar bonds, AND

2. The polar bonds are arranged asymmetrically so dipoles do not cancel

Examples of non-polar molecules with polar bonds:

• CO₂ — linear, two C=O dipoles point in opposite directions and cancel
• BF₃ — trigonal planar, three B–F dipoles cancel by symmetry
• CCl₄ — tetrahedral, four C–Cl dipoles cancel
• SF₆ — octahedral, all dipoles cancel

Examples of polar molecules:

• H₂O — bent shape, two O–H dipoles do not cancel (net dipole points towards oxygen)
• NH₃ — trigonal pyramidal, lone pair creates asymmetry
• CHCl₃ — three C–Cl dipoles do not cancel with the single C–H dipole
• HCl — only one bond, so the dipole is the molecular dipole

Intermolecular Forces

The polarity of molecules determines the types of intermolecular forces they exhibit:

• London dispersion forces (all molecules) — arise from temporary induced dipoles; strength increases with number of electrons / surface area
• Permanent dipole–dipole forces — between polar molecules; the δ⁺ end of one molecule attracts the δ⁻ end of another
• Hydrogen bonding — a special, strong type of dipole–dipole interaction occurring when H is bonded to N, O, or F (highly electronegative atoms with a lone pair)

Dative (Coordinate) Bonding and Shape

In a dative bond, both electrons in the bonding pair come from one atom. Once formed, a dative bond is identical to an ordinary covalent bond and counts as one bonding pair for VSEPR.

Example: NH₄⁺

• NH₃ donates its lone pair to H⁺
• Result: 4 bonding pairs, 0 lone pairs → tetrahedral, 109.5°
• The shape changes from trigonal pyramidal (in NH₃) to tetrahedral

Example: Al₂Cl₆

• Each Al has 3 ordinary bonds + 1 dative bond from a Cl lone pair
• Each Al is surrounded by 4 bonding pairs → tetrahedral around each Al

Shapes of Ions

Ions follow the same VSEPR rules. Count electrons carefully — subtract one for each positive charge, add one for each negative charge.

• NO₃⁻: N has 5 outer e⁻ + 1 (charge) = 6; three bonds to O atoms, no lone pairs → trigonal planar, 120°
• SO₄²⁻: S has 6 outer e⁻ + 2 (charge) = 8; four bonds to O atoms, no lone pairs → tetrahedral, 109.5°

Exam Tips

• Always count electron regions (a double bond = 1 region) rather than individual bonds
• Draw 3D representations using wedge and dashed bonds where appropriate
• When asked about molecular polarity, first identify polar bonds, then consider whether the shape causes them to cancel
• State the direction of the overall dipole when describing a polar molecule