AP Chemistry: Molecular Geometry Prediction

Why does a water molecule bend while carbon dioxide stays straight? The answer determines how substances interact, from the shape of DNA to the design of pharmaceuticals and materials. Predicting molecular geometry is not just about drawing shapes; it’s about understanding the fundamental forces that dictate chemical behavior, biological function, and the properties of everything around you. This guide will equip you with the systematic framework of Valence Shell Electron Pair Repulsion (VSEPR) theory to visualize and name the three-dimensional structures of molecules.

The Foundation: VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory is the central model for predicting molecular shapes. Its core postulate is simple yet powerful: electron groups—whether they are bonding pairs or lone pairs—arrange themselves in three-dimensional space to be as far apart as possible. This minimization of repulsive forces dictates the initial geometric template of a molecule. An electron group is defined as any region of electron density: a single bond, a double bond, a triple bond, or a lone pair. Each of these counts as one "group" for the initial geometry prediction because the electrons within them occupy a region of space that repels other regions.

The first critical step is determining the steric number of the central atom. This is the sum of the number of atoms bonded to the central atom and the number of lone pairs on the central atom. For example, in ammonia ($\text{ce}NH3$), nitrogen is bonded to three hydrogen atoms and has one lone pair, giving a steric number of 4. This number is the key that unlocks the initial electron domain geometry.

From Electron Domains to Molecular Geometry

The steric number determines the electron domain geometry—the shape described by the positions of all electron groups. The molecular geometry is what you get when you ignore the lone pairs and look only at the positions of the atoms. This distinction is crucial, as lone pairs cause repulsive distortions.

• Steric Number 2: Two electron groups maximize separation at 180°, resulting in a linear electron domain geometry. If both groups are bonding pairs, the molecular geometry is also linear (e.g., $\text{ce}CO2$).
• Steric Number 3: Three groups arrange in a trigonal planar geometry with ideal 120° angles. Examples include boron trifluoride ($\text{ce}BF3$). If one of those groups is a lone pair, the molecular geometry becomes bent or angular (e.g., $\text{ce}SO2$), with a bond angle slightly less than 120° due to greater lone pair repulsion.
• Steric Number 4: Four electron groups adopt a tetrahedral geometry with approximately 109.5° angles. Methane ($\text{ce}CH4$) is the perfect example. Key distortions occur here:
• One lone pair yields a trigonal pyramidal molecular geometry (e.g., $\text{ce}NH3$), with bond angles compressed to about 107°.
• Two lone pairs yield a bent molecular geometry (e.g., $\text{ce}H2O$), with bond angles compressed further to about 104.5°.

Expanding to Five and Six Electron Groups

For central atoms with expanded octets (typically period 3 or higher), we encounter five and six electron groups.

Steric Number 5: The five groups arrange in a trigonal bipyramidal electron domain geometry. This shape has two distinct positions: three equatorial positions (in a plane, 120° apart) and two axial positions (above and below the plane). Lone pairs always occupy the more spacious equatorial positions first to minimize repulsion.

• 0 lone pairs: Trigonal bipyramidal molecular geometry (e.g., $\text{ce}PCl5$).
• 1 lone pair: The molecular geometry is seesaw (or distorted tetrahedron), as in sulfur tetrafluoride ($\text{ce}SF4$).
• 2 lone pairs: T-shaped molecular geometry, as in chlorine trifluoride ($\text{ce}ClF3$).
• 3 lone pairs: Linear molecular geometry (e.g., $\text{ce}XeF2$).

Steric Number 6: Six electron groups achieve maximum separation in an octahedral geometry with 90° angles. All positions are equivalent.

• 0 lone pairs: Octahedral molecular geometry (e.g., $\text{ce}SF6$).
• 1 lone pair: Square pyramidal molecular geometry (e.g., $\text{ce}BrF5$).
• 2 lone pairs: With two lone pairs, they will be opposite each other to minimize repulsion, resulting in a square planar molecular geometry (e.g., $\text{ce}XeF4$).

Predicting and Understanding Bond Angles

The ideal bond angles stem from perfect geometric shapes: 180° (linear), 120° (trigonal planar), 109.5° (tetrahedral), 90°/120° (trigonal bipyramidal), and 90° (octahedral). VSEPR theory explains deviations through a hierarchy of repulsive strength: Lone Pair-Lone Pair > Lone Pair-Bonding Pair > Bonding Pair-Bonding Pair. A lone pair, being held closer to the nucleus, exerts a stronger repulsive force than bonding pairs, which are pulled away by the bonded atom. This is why the bond angles in $\text{ce}NH3$ and $\text{ce}H2O$ are successively smaller than the ideal tetrahedral angle.

Furthermore, in structures like the trigonal bipyramid, the presence of a lone pair distorts angles because repulsion is not uniform. Double and triple bonds also count as one electron group but have greater electron density than single bonds, leading to slight repulsive differences, such as the bond angles in molecules like formaldehyde ($\text{ce}H2CO$) being slightly distorted from ideal trigonal planar.

Common Pitfalls

1. Confusing Electron Domain Geometry with Molecular Geometry: This is the most frequent error. Always determine the steric number and electron domain geometry first. Then derive the molecular geometry by considering only the positions of the atoms. For $\text{ce}H2O$, the electron domain geometry is tetrahedral, but the molecular geometry is bent.
2. Misidentifying the Central Atom or Lone Pairs: In complex molecules or polyatomic ions, incorrectly identifying the central atom will derail everything. Remember, the central atom is usually less electronegative and bonded to multiple atoms. Also, always account for charge when placing lone pairs; a negative charge on a central atom often means an extra lone pair.
3. Incorrect Lone Pair Placement in Trigonal Bipyramidal Systems: Placing a lone pair in an axial position creates destabilizing 90° interactions with three equatorial groups. Placing it in an equatorial position creates only two 90° interactions. Lone pairs always go equatorial first in a trigonal bipyramidal framework.
4. Assuming All Bond Angles are Ideal: Memorizing shapes is not enough. You must be able to compare bond angles. For instance, in a seesaw molecule ($\text{ce}SF4$), the axial bond angles are less than 180°, and the equatorial bond angles are less than 120° due to the lone pair's strong repulsion. Always ask: where are the lone pairs, and how will they compress the angles?

Summary

• VSEPR theory predicts shape based on the repulsion between electron groups (bonds and lone pairs), which arrange to maximize separation.
• The steric number (number of atoms bonded + number of lone pairs on the central atom) determines the initial electron domain geometry, which is then used to find the molecular geometry.
• Lone pairs exert stronger repulsion than bonding pairs, distorting bond angles away from ideal values (e.g., 109.5° → ~107° in $\text{ce}NH3$ → ~104.5° in $\text{ce}H2O$).
• In trigonal bipyramidal systems, lone pairs occupy equatorial positions to minimize repulsion, leading to derived shapes like seesaw and T-shaped.
• Mastering the progression from steric number to electron domain geometry to molecular geometry, while accounting for lone pair distortions, allows you to accurately predict and visualize the 3D structure of any molecule governed by VSEPR principles.