Do lone pairs repel more than bonded pairs?

Yes! Bonded electrons are stretched out between two nuclei, which keeps them somewhat contained. Lone pairs are held by just one nucleus, so they puff out closer to the central atom, taking up more space and squeezing the bond angles together slightly.

Why do double bonds count as only one group?

Even though a double bond contains four electrons, those electrons are all forced to occupy the same general region of space between the two bonded atoms. Since they can't split up and move to opposite sides of the central atom, VSEPR treats them as a single repulsive domain.

What happens if a molecule does not have a central atom?

For larger, more complex molecules without a single central atom, you apply VSEPR rules to each interior atom individually. You end up describing the local shape around each specific atom rather than a single shape for the entire large molecule.

How does VSEPR predict molecular shape?

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the 3D shape of a molecule by looking at the electrons around its central atom. Because electrons are negatively charged, they naturally repel each other. To stay as stable as possible, these electron groups arrange themselves in space so they are as far apart as they can get. Think of tying a bunch of balloons together at the knots. They naturally push against each other and fan out into a specific symmetrical shape. VSEPR applies this exact same logic to electron pairs pushing away from one another around a central nucleus, creating predictable geometric structures.

Counting electron groups

To use VSEPR, you first need to know how many 'groups' of electrons surround the central atom. An electron group can be a single bond, a double bond, a triple bond, or a lone pair of non-bonding electrons. No matter how many electrons are shared in a bond, a single, double, or triple bond all count as just one electron group because they point in the same direction.

Electron geometry vs. Molecular shape

There is a subtle but important difference between where the electrons are and what the molecule actually looks like. The 'electron geometry' looks at all electron groups, including lone pairs. The 'molecular shape' only looks at the atoms. Lone pairs are invisible when we describe the final molecular shape, but they still act like invisible balloons, pushing the bonded atoms into specific angles.

Common shapes you will see

If you have two electron groups, they push 180 degrees apart, making a 'linear' shape. Three groups form a flat triangle called 'trigonal planar' with 120-degree angles. Four groups pop into 3D space to form a 'tetrahedral' shape with angles of 109.5 degrees. When you replace bonds with lone pairs in these geometries, you get variations like 'bent' or 'trigonal pyramidal'.

Where students slip up

The most common mistake is forgetting to draw the Lewis structure first. If you just look at the chemical formula, you might miss the lone pairs on the central atom. For example, $CO_2$ and $H_2O$ both have three atoms, but $CO_2$ is linear while $H_2O$ is bent because oxygen has two hidden lone pairs pushing the hydrogen atoms down.

Worked through

Predict the molecular shape of ammonia ( $NH_3$ ).

First, draw the Lewis structure for $NH_3$ . Nitrogen is in the center, bonded to three hydrogen atoms, and has one lone pair of electrons left over to complete its octet. Next, count the electron groups around the central nitrogen. There are 3 single bonds and 1 lone pair, giving a total of 4 electron groups. Four electron groups mean the electron geometry is tetrahedral. Finally, determine the molecular shape. Since one of those four groups is an invisible lone pair, the three hydrogen atoms are pushed down to form a pyramid base. This specific shape is called trigonal pyramidal.

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Where this comes from: OpenStax Chemistry 2e, Chapter 7: Chemical Bonding and Molecular Geometry · Khan Academy: Chemical bonds and molecular geometry unit