Chemical Bonding Class 11: VSEPR Theory Made Simple

VSEPR theory is the part of Chemical Bonding and Molecular Structure in Class 11 that lets you predict the shape of a molecule with nothing but a pencil. In this post you will learn the core idea, a four-step method, worked examples from linear BeCl₂ to T-shaped ClF₃, common mistakes, and quick questions with answers.

The big idea: electron pairs push apart

VSEPR stands for Valence Shell Electron Pair Repulsion. Its main postulate is simple:

• Electron pairs around a central atom, both bonding pairs and lone pairs, arrange themselves as far apart as possible to keep repulsion as low as possible.
• Repulsion is strongest between two lone pairs and weakest between two bond pairs:

lone pair-lone pair (lp-lp) > lone pair-bond pair (lp-bp) > bond pair-bond pair (bp-bp)

Think of balloons tied together at their knots. However many you tie, they spread out to keep as far from each other as they can. A lone pair is like a fatter balloon: it takes more room and squeezes the others.

Shapes with no lone pairs

If the central atom A has only bonding pairs to atoms B (type ABₙ), the shape is the ideal one:

The chapter's examples match these shapes: BeCl₂ is linear, BCl₃ is trigonal planar, CH₄ is tetrahedral, PCl₅ is trigonal bipyramidal and SF₆ is octahedral. (The 180°, 120° and 109.5° angles are the ones the chapter gives for sp, sp² and sp³ hybridisation.)

Shapes with lone pairs

When the central atom has lone pairs (written E), the lone pairs take up more space and push the bonding pairs closer, so the ideal angles shrink. The shape is named only by where the atoms are, not the lone pairs.

The four-step method

1. Find the central atom and count the atoms bonded to it.
2. Count the lone pairs on the central atom. For molecules where each outer atom forms one single bond with the central atom: lone pairs = (valence electrons of the central atom − number of bonds) ÷ 2.
3. Add bonding pairs and lone pairs to get the total electron pairs. This gives the basic arrangement (4 pairs means tetrahedral arrangement, and so on).
4. Name the shape from the atoms only, and remember that lone pairs shrink the angles.

Worked examples

Example 1: BeCl₂. Be has 2 valence electrons and forms 2 bonds. Lone pairs = (2 − 2) ÷ 2 = 0. Two bonding pairs, no lone pairs: type AB₂, linear, 180°.

Example 2: BCl₃. B has 3 valence electrons and forms 3 bonds. Lone pairs = (3 − 3) ÷ 2 = 0. Type AB₃: trigonal planar, 120°.

Example 3: CH₄, NH₃ and H₂O side by side. All three have four electron pairs around the central atom, but different numbers of lone pairs:

• CH₄: C has 4 valence electrons, 4 bonds. Lone pairs = (4 − 4) ÷ 2 = 0. Perfect tetrahedron, 109.5°.
• NH₃: N has 5 valence electrons, 3 bonds. Lone pairs = (5 − 3) ÷ 2 = 1. Type AB₃E, trigonal pyramidal, 107°.
• H₂O: O has 6 valence electrons, 2 bonds. Lone pairs = (6 − 2) ÷ 2 = 2. Type AB₂E₂, bent, 104.5°.

The pattern is the key lesson: as lone pairs increase from 0 to 1 to 2, the angle shrinks from 109.5° to 107° to 104.5°, because lp-bp and lp-lp repulsions are stronger than bp-bp repulsion.

Example 4: PCl₅. Five bonding pairs, no lone pairs: trigonal bipyramidal. Its two axial bonds are slightly longer and weaker than the three equatorial bonds, because they feel more repulsion from the equatorial bond pairs.

Example 5: shape and polarity. Shape decides whether bond dipoles cancel. CO₂ is linear, so its two equal C=O bond dipoles point in opposite directions and cancel: the dipole moment is zero. In symmetrical BF₃ and CCl₄ the bond dipoles also cancel, giving zero net dipole moment.

Common mistakes to avoid

• Naming the shape from all electron pairs. NH₃ has four electron pairs, but its shape is trigonal pyramidal, not tetrahedral. The shape describes atoms only.
• Forgetting lone pairs. Always count lone pairs on the central atom before deciding the shape.
• Mixing up the repulsion order. It is lp-lp > lp-bp > bp-bp, not the other way round.
• Assuming every four-pair molecule has 109.5°. Only CH₄ does here; NH₃ is 107° and H₂O is 104.5°.
• Using the counting shortcut where it does not fit. The (valence − bonds) ÷ 2 rule in Step 2 assumes single bonds to monovalent atoms. For molecules like SO₂, use the table above.

Quick check questions

1. What is the shape of an AB₃ molecule with no lone pairs? Answer: Trigonal planar, as in BCl₃.

2. Why is the bond angle in H₂O smaller than in NH₃? Answer: H₂O has two lone pairs and NH₃ has one. Two lone pairs push the bonding pairs together more strongly, so the angle is smaller (104.5° against 107°).

3. How many lone pairs does the central atom in ClF₃ have, and what is its shape? Answer: Two lone pairs (type AB₃E₂). Its shape is T-shaped.

4. Which repulsion is strongest: lp-lp, lp-bp or bp-bp? Answer: lp-lp.

5. Why is the dipole moment of CO₂ zero? Answer: CO₂ is linear, so the two equal C=O bond dipoles are opposite and cancel.

6. NH₃ has four electron pairs. Why is it not called tetrahedral? Answer: One of the four pairs is a lone pair. The shape is named from the three N-H bonds only, which form a trigonal pyramid.

Once VSEPR clicks, hybridisation and bond angles in the rest of the chapter become much easier. To practise more shapes step by step, study this chapter with Joy.