Electron geometry is the three-dimensional arrangement of all electron groups around a central atom, including both bonding pairs and lone pairs. It describes where electrons sit in space, not where the surrounding atoms end up. There are five fundamental electron geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. Which one applies depends entirely on how many electron groups surround the central atom.
Why Electrons Arrange Themselves This Way
Electrons carry a negative charge, so groups of electrons repel each other. The most stable arrangement is the one where electron groups are as far apart from each other as possible, minimizing that electrostatic repulsion. This principle is the foundation of VSEPR theory (valence shell electron pair repulsion), the model chemists use to predict electron geometry.
Think of it like tying balloons to a central point. Two balloons naturally point in opposite directions. Three spread out into a flat triangle. Four push into a three-dimensional shape. The electrons do the same thing, settling into whichever arrangement puts the most distance between each group.
The Five Electron Geometries
Each electron geometry corresponds to a specific number of electron groups, sometimes called the steric number. An “electron group” can be a single bond, a double bond, a triple bond, or a lone pair. It doesn’t matter whether a bond is single or double for counting purposes; each one counts as one group.
- Linear (2 electron groups): The two groups point in opposite directions, forming a straight line with a 180° angle. Carbon dioxide (CO₂) is a classic example.
- Trigonal planar (3 electron groups): Three groups spread evenly across a flat plane, 120° apart. Boron trifluoride (BF₃) takes this shape.
- Tetrahedral (4 electron groups): Four groups push into three dimensions, pointing toward the corners of a tetrahedron at 109.5° from each other. Methane (CH₄) is the textbook case.
- Trigonal bipyramidal (5 electron groups): Five groups arrange with three in a central plane at 120° and two pointing straight up and down at 90° to the plane. Phosphorus pentafluoride (PF₅) follows this geometry.
- Octahedral (6 electron groups): Six groups point toward the corners of an octahedron, with 90° angles between adjacent groups and 180° between opposite groups. Sulfur hexafluoride (SF₆) is the standard example.
Electron Geometry vs. Molecular Geometry
This distinction trips up more students than almost anything else in introductory chemistry. Electron geometry accounts for every electron group around the central atom. Molecular geometry only describes where the actual atoms are. When every electron group is a bond (no lone pairs), the two geometries are identical. The moment lone pairs enter the picture, they diverge.
Water is the perfect illustration. Oxygen in H₂O has four electron groups: two bonds to hydrogen atoms and two lone pairs. That gives it a tetrahedral electron geometry. But since you can’t “see” the lone pairs as physical atoms, the molecular shape is bent, with just the two hydrogen atoms and the oxygen forming a V shape. The electron geometry is tetrahedral; the molecular geometry is bent.
Ammonia (NH₃) works the same way. Nitrogen has four electron groups: three bonds and one lone pair. The electron geometry is tetrahedral, but the molecular shape is trigonal pyramidal because you’re only looking at where the three hydrogen atoms sit relative to nitrogen.
How Lone Pairs Change Bond Angles
Lone pairs don’t just change the name of the molecular shape. They physically squeeze the bond angles smaller. A lone pair spreads out more than a bonding pair because it isn’t being pulled toward a second nucleus. It’s held close to just one atom, so it occupies more space.
The repulsion hierarchy works like this: lone pair vs. lone pair is the strongest repulsion, lone pair vs. bonding pair is moderate, and bonding pair vs. bonding pair is the weakest. In ammonia, this means the lone pair pushes the three N-H bonds slightly closer together, compressing the bond angle below the ideal 109.5° of a perfect tetrahedron. Water, with two lone pairs, gets squeezed even further.
How to Determine Electron Geometry
You can figure out the electron geometry of any molecule in a few steps. Start by drawing the Lewis structure, which shows you where all the bonding pairs and lone pairs are. Then count the total number of electron groups around the central atom. Remember, a double or triple bond still counts as one group.
Once you have that count, match it to the geometry:
- 2 groups: linear
- 3 groups: trigonal planar
- 4 groups: tetrahedral
- 5 groups: trigonal bipyramidal
- 6 groups: octahedral
That’s it for electron geometry. If you also need the molecular geometry, take the additional step of noting how many of those groups are lone pairs and adjusting the shape name accordingly. But the electron geometry itself depends only on the total count.
The Connection to Hybridization
If you’re studying organic chemistry, you’ll notice that electron geometry lines up directly with orbital hybridization. A linear arrangement (2 groups) corresponds to sp hybridization. Trigonal planar (3 groups) matches sp² hybridization. Tetrahedral (4 groups) matches sp³. For five and six groups, the hybridizations are sp³d and sp³d², respectively.
This isn’t a coincidence. Hybridization is the mathematical description of the same physical reality: the central atom blends its orbitals to point electron density in the directions that minimize repulsion. So knowing the electron geometry immediately tells you the hybridization, and vice versa. If a carbon atom has a double bond and two single bonds (three electron groups), you know it’s trigonal planar and sp² hybridized without any additional calculation.
Why Electron Geometry Matters
Electron geometry is the starting point for understanding nearly everything about a molecule’s physical behavior. It determines the molecular shape, which in turn controls whether a molecule is polar or nonpolar. Polarity affects boiling points, solubility, and how molecules interact with each other. Carbon dioxide is linear and nonpolar, which is why it’s a gas at room temperature. Water is bent and highly polar, which is why it’s a liquid that dissolves salts and sugars so effectively.
Getting the electron geometry right is the first domino. Every prediction about a molecule’s shape, polarity, and reactivity follows from correctly counting electron groups and identifying where they sit in three-dimensional space.