Teaching Molecular Shapes with VSEPR

Molecules are three-dimensional, but almost everything students see of them is flat: a page, a whiteboard, a screen. Draw water as H–O–H in a straight line and many students will believe the molecule is straight, because that is what they were shown. Teaching VSEPR (valence shell electron pair repulsion) theory is therefore as much about getting students to think in three dimensions as about learning a table of shapes and angles. The good news is that the theory itself is short and logical, and it responds well to hands-on models.

This guide describes a three-lesson sequence for students of about 16 to 18 (it can be trimmed to two lessons for younger groups by stopping at four electron pairs). It includes timings, the practical activities that work best, the misconceptions to plan for, and a graded question bank.

Learning goals

Students should be able to:

1. State the core idea of VSEPR: regions of electron density around a central atom repel each other and arrange themselves as far apart as possible.
2. Count bonding pairs and lone pairs on a central atom from a Lewis structure, treating a multiple bond as one region.
3. Predict the electron-pair arrangement and the molecular shape for two to six regions.
4. Explain why lone pairs reduce bond angles (lone pair–lone pair > lone pair–bond pair > bond pair–bond pair repulsion).
5. Draw shapes using wedge and dash notation.

Prior knowledge

• Drawing dot-and-cross or Lewis structures, including lone pairs, double bonds and simple ions (see Lewis structures practice).
• Valence electrons from the group number.
• The idea that like charges repel.

If Lewis structures are weak, spend the first ten minutes of lesson 1 on them. VSEPR cannot start until students can find the lone pairs.

Lesson 1: why shapes exist (60 minutes)

Starter: is water straight? (5 minutes)

Draw H–O–H in a line and ask students to vote: is water linear or bent? Then ask what evidence could decide. (Water is polar, and a straight, symmetric molecule would not be; spectroscopy measures the angle directly.) Leave the question open.

Activity: balloon domains (20 minutes)

Give each group long modelling balloons (or ordinary round balloons) inflated to the same size. Twist or tie together two, then three, four, five and six at their knots. Each balloon represents one region of electron density around the central atom, and the balloons push each other into the arrangement that gives each the most room:

• 2 balloons → linear (180°)
• 3 → trigonal planar (120°)
• 4 → tetrahedral (109.5°)
• 5 → trigonal bipyramidal (90° and 120°)
• 6 → octahedral (90°)

The four-balloon cluster is the moment to watch for. Many students expect a flat cross with 90° angles and are surprised when the balloons settle into a tetrahedron. Ask them to explain why 109.5° beats 90°. Check for latex allergies before the lesson and use latex-free balloons if needed.

Main: from domains to shapes (25 minutes)

Now make one balloon a different colour and call it a lone pair. With four balloons, one white (lone pair) and three coloured (bonds), ask: “If you could only see the atoms, what shape would you see?” This separates the two ideas students most often merge:

• electron-pair geometry: the arrangement of all regions, including lone pairs;
• molecular shape: the arrangement of the atoms only.

Walk through CH₄, NH₃ and H₂O using the same four-balloon cluster. The electron-pair geometry is tetrahedral each time; the shape is tetrahedral, trigonal pyramidal and bent.

Plenary (10 minutes)

Return to the starter vote. Students now write a three-sentence explanation of why water is bent.

Lesson 2: angles, lone pairs and multiple bonds (60 minutes)

Starter: the angle puzzle (10 minutes)

Show the bond angles of CH₄ (109.5°), NH₃ (about 107°) and H₂O (about 104.5°). All three have four electron pairs. Ask pairs of students to suggest why the angles differ.

Guide them to the explanation: a lone pair is held by only one nucleus, so it sits closer to the central atom and spreads out more, pushing the bonding pairs together. Each lone pair squeezes the angle by roughly 2 to 2.5°. Rank the repulsions: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. Lone pairs and molecular shape and bond angles explained give the fuller treatment.

Model kits (25 minutes)

Students build, sketch in wedge–dash form and record: CO₂, BF₃, CH₄, NH₃, H₂O, SO₂, NH₄⁺, and the carbonate ion. Use kits with a separate piece (or a small blob of modelling clay) for lone pairs so they are not forgotten. Two teaching points emerge:

• Double bonds count as one region. CO₂ has two double bonds and is linear, not bent. SO₂ has two bonding regions and one lone pair, so it is bent at slightly less than 120°.
• Ions work the same way. NH₄⁺ has four bonding pairs and is tetrahedral, just like CH₄; add an electron for each negative charge and remove one for each positive charge when counting.

Wedge–dash drawing practice (15 minutes)

Model the convention once: a plain line lies in the page, a solid wedge comes towards the viewer, a hashed wedge goes away. Students draw CH₄ and NH₃, then swap drawings and build each other’s molecules from the sketch alone. If the partner can’t build it, the drawing is ambiguous.

Plenary (10 minutes)

Mini whiteboards: call out a formula; students hold up the shape and angle.

Lesson 3: expanded octets and the method (60 minutes)

Five and six regions (20 minutes)

Use the five-balloon cluster to show that trigonal bipyramidal positions are not equivalent: three equatorial (120° apart) and two axial (90° to the equator). Lone pairs go equatorial, where they have fewer 90° neighbours. Build or show:

• SF₄: 4 bonding + 1 lone → seesaw
• ClF₃: 3 bonding + 2 lone → T-shaped
• XeF₂: 2 bonding + 3 lone → linear
• SF₆: 6 bonding → octahedral
• XeF₄: 4 bonding + 2 lone → square planar (the lone pairs are opposite each other)

Consolidating the method (20 minutes)

Give students a five-step routine and have them apply it without help to new molecules: (1) find the central atom; (2) count its valence electrons, adjust for charge and bonded atoms; (3) find bonding and lone pairs; (4) choose the electron-pair geometry; (5) name the shape from the atoms and adjust the angle. A summary list of shapes is in molecular shapes list.

Linking shape to polarity (15 minutes)

Close the loop from lesson 1: water is polar because it is bent. Compare CO₂ with SO₂, and CCl₄ with CHCl₃. Symmetric shapes cancel bond dipoles; asymmetric shapes do not. This prepares students for intermolecular forces (see teaching intermolecular forces).

Exit ticket (5 minutes)

Two molecules not seen before, for example PCl₃ and ICl₄⁻.

Misconceptions to expect

Misconception	What to do
“Shapes are flat, as drawn on paper.”	Balloons and kits before any drawings; teach wedge–dash early; have students build from each other’s sketches
“Lone pairs don’t count, because they’re not bonds.”	The coloured balloon: lone pairs take up space and determine the arrangement even though they’re invisible in the shape name
“The shape name includes the lone pairs.”	Keep two columns in every table: electron-pair geometry and molecular shape
“A double bond counts as two regions.”	Build CO₂; ask whether a double bond can point in two directions at once
“Water’s angle is 90°, because the H atoms are drawn at right angles.”	Four regions give a tetrahedral base; 90° would be worse for repulsion
“Bonds repel because the atoms repel.”	Make clear that it is the electron pairs that repel
“VSEPR always gives exact angles.”	Show PH₃ (about 93.5°) against NH₃ (about 107°); VSEPR predicts the direction of change, not every number
“Square planar and tetrahedral are the same for four atoms.”	Compare CH₄ with XeF₄: same number of atoms, different number of electron regions

Differentiation

Support

• A printed flowchart and a partly filled shape table (regions, bonding, lone, shape, angle), completed as the lessons progress.
• Keep kits on the desk throughout, not just during the kit activity. Many students need the physical model for every question at first.
• Limit early practice to central atoms from period 2 with four or fewer regions.

Stretch

• Explain why lone pairs take equatorial positions in trigonal bipyramidal structures, counting the 90° interactions each position has.
• Discuss why PH₃ and H₂S have angles close to 90°, and what that says about the limits of the model.
• Apply VSEPR to the carbon atoms in ethene and ethyne, and to each central atom in a larger molecule such as ethanoic acid.

Question bank

Recall

1. What does VSEPR stand for? (Valence shell electron pair repulsion.)
2. State the bond angle in a tetrahedral molecule. (109.5°.)

Application 3. Predict the shape of PCl₃ and give its approximate bond angle. (Three bonding pairs, one lone pair: trigonal pyramidal, a little under 109.5°. Accept about 107° for a VSEPR answer.) 4. Predict the shape of BeCl₂ in the gas phase. (Two bonding pairs: linear, 180°.) 5. What shape is the ICl₄⁻ ion? (Four bonding pairs, two lone pairs: square planar, 90°.)

Explanation 6. Explain why the bond angle in NH₃ is smaller than in CH₄. (Both have four pairs; NH₃ has one lone pair, which repels more strongly than bonding pairs and pushes the N–H bonds closer together.) 7. CO₂ and SO₂ both have two oxygen atoms bonded to a central atom. Explain why their shapes differ. (CO₂: two double-bond regions, no lone pairs, linear. SO₂: two bonding regions plus one lone pair, bent.)

Challenge 8. Explain why XeF₄ is square planar rather than tetrahedral, and why it is non-polar.

Key takeaways

• Start with three-dimensional models (balloons, then kits) before any drawings.
• Separate electron-pair geometry from molecular shape from the first lesson, and keep both in every table.
• Treat lone pairs as full regions that take up more room than bonds; treat multiple bonds as one region.
• Teach wedge–dash drawing explicitly and check it by having students build from each other’s sketches.
• Use shape to explain polarity, which leads straight into intermolecular forces.