Molecular Geometry and VSEPR

students, imagine trying to predict the shape of a tent just from where the poles are attached ⛺. In chemistry, molecules also have shapes, and those shapes help explain why substances behave the way they do. In this lesson, you will learn how the Valence Shell Electron Pair Repulsion model, or VSEPR, helps predict molecular geometry. You will also see how molecular shape connects to bonding, intermolecular forces, polarity, and the properties of materials.

What VSEPR Tries to Explain

Atoms in a molecule are not arranged randomly. The electrons around a central atom repel each other, so they spread out as much as possible. The VSEPR model uses this idea to predict the electron pair geometry and the molecular geometry of a molecule.

The main idea is simple: electron pairs around a central atom repel one another because they are all negatively charged. They arrange themselves to stay as far apart as possible. This gives a molecule a specific 3D shape.

There are two important terms here:

• Electron pair geometry: the arrangement of all electron pairs, including bonding pairs and lone pairs, around the central atom.
• Molecular geometry: the arrangement of atoms only, ignoring lone pairs.

For example, in water $\mathrm{H_2O}$, the oxygen atom has four electron pairs around it, but only two are bonding pairs. So the electron pair geometry is tetrahedral, while the molecular geometry is bent.

VSEPR is very useful because shape affects many properties such as melting point, boiling point, solubility, and whether a molecule is polar.

Counting Electron Domains

To use VSEPR, first count the electron domains around the central atom. An electron domain is any region of electron density. These include:

• a single bond
• a double bond
• a triple bond
• a lone pair

A key point is that multiple bonds count as one electron domain because they occupy one region of space around the atom.

Let’s look at a few examples.

Example 1: Carbon dioxide, $\mathrm{CO_2}$

The central carbon has two double bonds, one to each oxygen. That gives two electron domains. Two domains repel each other into a straight line, so the electron pair geometry is linear and the molecular geometry is also linear.

The bond angle is $180^\circ$.

Example 2: Methane, $\mathrm{CH_4}$

Carbon has four single bonds and no lone pairs. That gives four electron domains. They spread out into a tetrahedral arrangement.

The bond angle is about $109.5^\circ$.

Example 3: Ammonia, $\mathrm{NH_3}$

Nitrogen has three bonding pairs and one lone pair, so there are four electron domains total. The electron pair geometry is tetrahedral, but because one space is occupied by a lone pair, the molecular geometry is trigonal pyramidal.

The bond angle is slightly less than $109.5^\circ$, about $107^\circ$, because lone pairs repel more strongly than bonding pairs.

Shapes and Why Lone Pairs Matter

Lone pairs take up more space than bonding pairs. This is because lone pairs are attracted to only one nucleus, while bonding pairs are shared between two nuclei. As a result, lone pairs push bonding pairs closer together.

The repulsion strength usually follows this order:

$$\text{lone pair-lone pair} > \text{lone pair-bond pair} > \text{bond pair-bond pair}$$

This explains why bond angles often decrease when lone pairs are present.

Let’s compare two molecules with four electron domains:

• $\mathrm{CH_4}$ has four bonding pairs and no lone pairs, so the shape is tetrahedral.
• $\mathrm{NH_3}$ has three bonding pairs and one lone pair, so the shape is trigonal pyramidal.
• $\mathrm{H_2O}$ has two bonding pairs and two lone pairs, so the shape is bent.

Water has a smaller bond angle of about $104.5^\circ$ because two lone pairs compress the bond angle even more.

These differences matter because shape influences the overall distribution of charge in a molecule.

Common Molecular Geometries You Should Know

Here are some important shapes in IB Chemistry HL.

Linear

A linear molecule has atoms arranged in a straight line. The bond angle is $180^\circ$.

Examples:

• $\mathrm{CO_2}$
• $\mathrm{BeCl_2}$

Trigonal planar

Three electron domains arrange themselves in one plane with bond angles of about $120^\circ$.

Example:

• $\mathrm{BF_3}$

If one of the three domains is a lone pair, the molecular geometry becomes bent.

Tetrahedral

Four electron domains arrange in 3D space with bond angles of about $109.5^\circ$.

Examples:

• $\mathrm{CH_4}$
• $\mathrm{CCl_4}$

Trigonal pyramidal

Four electron domains with one lone pair give this shape.

Example:

• $\mathrm{NH_3}$

Bent

This shape occurs when lone pairs change a tetrahedral or trigonal planar arrangement.

Examples:

• $\mathrm{H_2O}$
• $\mathrm{SO_2}$

Trigonal bipyramidal and octahedral

These occur when the central atom has five or six electron domains.

• Five domains: trigonal bipyramidal, with bond angles of $90^\circ$, $120^\circ$, and $180^\circ$
• Six domains: octahedral, with bond angles of $90^\circ$ and $180^\circ$

Examples:

• $\mathrm{PCl_5}$ is trigonal bipyramidal
• $\mathrm{SF_6}$ is octahedral

How to Apply VSEPR in IB Chemistry HL

When solving exam questions, use a clear step-by-step method:

1. Draw the Lewis structure.
2. Count electron domains around the central atom.
3. Determine the electron pair geometry.
4. Remove lone pairs from the shape description to find the molecular geometry.
5. Estimate bond angles.
6. Decide whether the molecule is polar or nonpolar if asked.

Let’s try a full example with sulfur dioxide, $\mathrm{SO_2}$.

First, draw the Lewis structure. Sulfur is the central atom, and the molecule has two bonding regions and one lone pair on sulfur. That gives three electron domains total.

Three domains means trigonal planar electron pair geometry. But because there is one lone pair, the molecular geometry is bent. The bond angle is a little less than $120^\circ$.

This shape helps explain why $\mathrm{SO_2}$ is a polar molecule. The bent geometry means the bond dipoles do not cancel.

Now compare that with carbon dioxide, $\mathrm{CO_2}$. Even though each $\mathrm{C=O}$ bond is polar, the molecule is linear, so the two bond dipoles cancel out. That makes the molecule nonpolar.

This is a great example of why shape matters more than just bond polarity.

Shape, Polarity, and Intermolecular Forces

Molecular geometry affects whether a molecule has a net dipole moment. A molecule is polar if its bond dipoles do not cancel.

Polarity matters because it affects intermolecular forces, which are attractions between molecules. Stronger intermolecular forces usually mean:

• higher boiling point
• higher melting point
• greater viscosity
• different solubility behavior

For example, water has a bent shape and is polar. This allows hydrogen bonding, a strong intermolecular force. That is why water has unusually high boiling point compared with similar-sized molecules.

In contrast, $\mathrm{CO_2}$ is linear and nonpolar, so it only has London dispersion forces between molecules. As a result, it has a much lower boiling point.

Another example is ammonia, $\mathrm{NH_3}$. Its trigonal pyramidal shape makes it polar, and it can form hydrogen bonds because nitrogen is bonded to hydrogen. This helps explain its physical properties.

So, students, when you see a molecule in an exam, think beyond the formula. Ask: What shape does it have? Is it symmetric? Do dipoles cancel? What intermolecular forces are possible? 🧠

Connection to Structure and Materials

Molecular geometry is part of the bigger topic of bonding and structure because shape helps explain why substances have different properties.

• Simple molecular substances often have low melting and boiling points because the intermolecular forces between molecules are weak compared with covalent bonds inside molecules.
• Polar molecules may dissolve better in polar solvents like water.
• Nonpolar molecules often dissolve better in nonpolar solvents.

The model also helps in understanding materials. For example, the shape of molecules can influence how they pack together in a solid, which affects density and melting point. In biology, molecular shape is crucial too, because enzyme-substrate interactions depend on shape matching.

VSEPR is a model, not a perfect picture. It works very well for many main-group compounds, but it does not explain everything in full detail. Still, it is one of the most important tools for predicting molecular shape in IB Chemistry HL.

Conclusion

Molecular geometry and VSEPR help you predict the 3D shapes of molecules by using electron pair repulsion. By counting electron domains and recognizing the effect of lone pairs, you can determine shapes such as linear, trigonal planar, tetrahedral, trigonal pyramidal, bent, trigonal bipyramidal, and octahedral.

This topic is important because shape affects polarity, intermolecular forces, and the physical properties of substances. It also links directly to the broader IB Chemistry HL theme of Structure 2 — Models of Bonding and Structure. In exams, a strong answer often includes the Lewis structure, electron domain count, molecular geometry, bond angle, and an explanation of polarity or intermolecular forces.

Study Notes

• VSEPR stands for Valence Shell Electron Pair Repulsion.
• Electron pairs around a central atom repel and arrange themselves as far apart as possible.
• An electron domain is a bond or a lone pair.
• Multiple bonds count as one electron domain.
• Electron pair geometry includes lone pairs and bonding pairs.
• Molecular geometry includes only the positions of atoms.
• Lone pairs repel more strongly than bonding pairs, so they reduce bond angles.
• Common shapes include linear $180^\circ$, trigonal planar $120^\circ$, tetrahedral $109.5^\circ$, trigonal pyramidal, bent, trigonal bipyramidal, and octahedral.
• Shape affects polarity by determining whether bond dipoles cancel.
• Polarity affects intermolecular forces, which affect melting point, boiling point, viscosity, and solubility.
• VSEPR is a key model for understanding structure-property relationships in IB Chemistry HL.