5.1 VSEPR theory and molecular shapes

Molecular Geometry Predictions

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts molecular shapes by looking at how electron groups arrange themselves around a central atom. The core idea: electron domains repel each other, so they spread out as far apart as possible to minimize that repulsion. This gives us predictable 3D geometries that directly influence a molecule's chemical behavior, polarity, and reactivity.

VSEPR Theory and Electron Domain Geometry

An electron domain is any region of electron density around a central atom. That includes bonding pairs (shared electrons forming a bond) and lone pairs (non-bonding valence electrons sitting on the central atom). Double and triple bonds each count as a single electron domain, since all the electrons in that bond occupy roughly the same region of space.

The total number of electron domains determines the electron domain geometry:

To find the number of electron domains, draw the Lewis structure first, then count the groups of electrons (bonds + lone pairs) around the central atom.

Predicting Molecular Geometry

Here's how to predict geometry step by step:

1. Draw the Lewis structure for the molecule or ion.
2. Count the total electron domains (bonding pairs + lone pairs) around the central atom. Remember: multiple bonds count as one domain.
3. Identify the electron domain geometry from the table above.
4. Determine how many of those domains are lone pairs vs. bonding pairs.
5. Name the molecular geometry based on where the atoms are (ignoring lone pairs).

When there are zero lone pairs, the molecular geometry matches the electron domain geometry:

• $BeF_2$: 2 bonding pairs, 0 lone pairs → linear
• $BCl_3$: 3 bonding pairs, 0 lone pairs → trigonal planar
• $CH_4$: 4 bonding pairs, 0 lone pairs → tetrahedral
• $PCl_5$: 5 bonding pairs, 0 lone pairs → trigonal bipyramidal
• $SF_6$: 6 bonding pairs, 0 lone pairs → octahedral

Molecular Shape and Bond Angles

Determining Molecular Shape

This is a distinction that trips people up: electron domain geometry describes the arrangement of all electron domains, while molecular geometry describes only the positions of the bonded atoms. You can't see lone pairs in the physical shape of a molecule, but they're still there influencing everything.

When all electron domains are bonding pairs, the two geometries are identical. The molecular geometries listed above (linear, trigonal planar, tetrahedral, etc.) apply directly. Things get more interesting when lone pairs enter the picture.

Bond Angles in Molecular Geometries

Bond angles are set by the electron domain geometry but modified by lone pairs. The ideal angles for each geometry are:

• Linear: 180°
• Trigonal planar: 120°
• Tetrahedral: 109.5°
• Trigonal bipyramidal: 90° (axial-equatorial) and 120° (equatorial-equatorial)
• Octahedral: 90°

Some concrete examples with no lone pairs on the central atom:

• $CO_2$: linear, 180° bond angles
• $SO_3$: trigonal planar, 120° bond angles
• $SiH_4$: tetrahedral, 109.5° bond angles

These ideal angles only hold perfectly when all the domains are equivalent. Once lone pairs show up, the angles shift.

Electron Domain vs. Molecular Geometry

Differences Between Electron Domain and Molecular Geometry

The key distinction worth repeating: electron domain geometry accounts for every electron group around the central atom, while molecular geometry only reflects where the bonded atoms sit in space.

If a molecule has no lone pairs on the central atom, the two geometries are the same. But as soon as lone pairs are present, the molecular geometry becomes a "subset" of the electron domain geometry. For example, water ($H_2O$) has four electron domains (tetrahedral electron domain geometry), but only two of those are bonding pairs. The molecular geometry you'd report is bent, not tetrahedral.

Effect of Lone Pairs on Molecular Geometry

Lone pairs repel more strongly than bonding pairs. This happens because lone pair electrons are held closer to the central atom and spread over a larger angular region, so they push harder on neighboring electron domains.

The repulsion hierarchy is:

This ordering explains why bond angles compress when lone pairs are present, and why lone pairs in trigonal bipyramidal geometries always occupy equatorial positions (where they have more room).

Lone Pairs and Molecular Geometry

Lone Pair Effects on Bond Angles

Because lone pairs take up more angular space than bonding pairs, they squeeze the bonding pairs closer together. The more lone pairs on the central atom, the greater the compression of bond angles away from the ideal values.

For the tetrahedral electron domain family, this trend is clear:

• $CH_4$: 0 lone pairs → 109.5° (ideal tetrahedral)
• $NH_3$: 1 lone pair → approximately 107° (compressed from 109.5°)
• $H_2O$: 2 lone pairs → approximately 104.5° (compressed even further)

Each additional lone pair pushes the bond angles down by a few degrees.

Examples of Lone Pair Effects

Starting from tetrahedral electron domain geometry (4 domains):

• 1 lone pair → trigonal pyramidal (e.g., $NH_3$, bond angles ~107°)
• 2 lone pairs → bent (e.g., $H_2O$, bond angles ~104.5°)

Starting from trigonal bipyramidal electron domain geometry (5 domains):

Lone pairs preferentially occupy equatorial positions to minimize 90° repulsive interactions.

• 1 lone pair → seesaw (e.g., $SF_4$)
• 2 lone pairs → T-shaped (e.g., $ClF_3$)
• 3 lone pairs → linear (e.g., $I_3^-$)

Starting from octahedral electron domain geometry (6 domains):

• 1 lone pair → square pyramidal (e.g., $BrF_5$)
• 2 lone pairs → square planar (e.g., $XeF_4$); the two lone pairs sit opposite each other to minimize lone pair–lone pair repulsion

In all these cases, the bond angles deviate from the ideal values of the parent electron domain geometry due to the extra repulsion from lone pairs.