Molecular Geometry Shapes

Why This Matters

Molecular geometry isn't just about memorizing shapes—it's the foundation for understanding polarity, intermolecular forces, and reactivity throughout AP Chemistry. When you're asked why water has a high boiling point or why $\text{CO}_2$ is nonpolar despite having polar bonds, you're really being tested on how electron pair repulsion, lone pair effects, and three-dimensional arrangement determine molecular behavior. These concepts connect directly to Unit 2's focus on compound structure and Unit 3's emphasis on how molecular shape influences physical properties like solubility and boiling point.

The key principle here is VSEPR theory (Valence Shell Electron Pair Repulsion): electron pairs around a central atom arrange themselves to minimize repulsion, and lone pairs repel more strongly than bonding pairs. This creates predictable distortions from ideal geometries. Don't just memorize that water is bent at 104.5°—understand why lone pairs compress bond angles and how that bent shape creates a dipole moment that explains water's remarkable properties.

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Geometries with No Lone Pairs: Ideal Shapes

When a central atom has only bonding pairs, electron groups spread out symmetrically to maximize distance from each other. These "ideal" geometries have predictable bond angles because there's no lone pair repulsion to distort them.

Linear

• Two electron groups at 180°—the simplest geometry where atoms arrange in a straight line
• Nonpolar when substituents are identical (e.g., $\text{CO}_2$), as bond dipoles cancel perfectly
• Common in molecules with double/triple bonds to the central atom, where multiple bonds count as one electron group

Trigonal Planar

• Three electron groups at 120°—all atoms lie in the same plane with the central atom
• Symmetrical arrangement creates nonpolar molecules when all three substituents are identical (e.g., $\text{BF}_3$)
• Appears in carbonate ions and aldehydes, making it essential for understanding organic functional groups

Tetrahedral

• Four electron groups at 109.5°—the most common geometry in organic chemistry
• Three-dimensional arrangement maximizes separation of electron pairs in all directions (e.g., $\text{CH}_4$)
• Foundation for understanding chirality and the geometry of carbon in biological molecules

Trigonal Bipyramidal

• Five electron groups with two distinct positions—axial (90° from equatorial) and equatorial (120° apart)
• Axial positions experience more repulsion than equatorial positions, which matters for predicting where lone pairs go
• Example: $\text{PCl}_5$ demonstrates how expanded octets occur in Period 3+ elements with available d orbitals

Octahedral

• Six electron groups at 90°—all positions are equivalent in the ideal geometry
• Highly symmetrical structure leads to nonpolar molecules when all substituents are identical (e.g., $\text{SF}_6$)
• Central atom uses $sp^3d^2$ hybridization, requiring elements from Period 3 or below

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Geometries Derived from Tetrahedral: One to Two Lone Pairs

When lone pairs replace bonding pairs in a tetrahedral electron geometry, the molecular shape changes while bond angles compress. Lone pairs occupy more space than bonding pairs, pushing bonded atoms closer together.

Trigonal Pyramidal

• Three bonding pairs + one lone pair gives a pyramidal shape with bond angles ~107° (e.g., $\text{NH}_3$)
• Lone pair compression reduces the ideal 109.5° tetrahedral angle by about 2.5°
• Creates a polar molecule even with identical substituents, since the geometry is asymmetrical

Bent (from Tetrahedral)

• Two bonding pairs + two lone pairs creates a bent shape with bond angles ~104.5° (e.g., $\text{H}_2\text{O}$)
• Two lone pairs cause greater compression than one, explaining why water's angle is smaller than ammonia's
• Responsible for water's polarity and its ability to form hydrogen bonds—a critical concept for intermolecular forces

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Geometries Derived from Trigonal Bipyramidal: Lone Pairs in Equatorial Positions

When lone pairs appear in a trigonal bipyramidal electron geometry, they always occupy equatorial positions first. Equatorial positions have only two 90° interactions (with axial groups), while axial positions have three 90° interactions—so lone pairs minimize repulsion by going equatorial.

Seesaw

• Four bonding pairs + one lone pair with the lone pair in an equatorial position (e.g., $\text{SF}_4$)
• Bond angles of ~90° and ~120° are compressed slightly from ideal due to lone pair repulsion
• Asymmetrical shape creates a polar molecule, important for predicting physical properties

T-Shaped

• Three bonding pairs + two lone pairs with both lone pairs occupying equatorial positions (e.g., $\text{ClF}_3$)
• Bond angles of approximately 87.5°—compressed from 90° by the two equatorial lone pairs
• Demonstrates the equatorial preference rule for lone pair placement in bipyramidal systems

Linear (from Trigonal Bipyramidal)

• Two bonding pairs + three lone pairs places all three lone pairs in equatorial positions (e.g., $\text{XeF}_2$)
• Results in 180° bond angle with fluorines in axial positions
• Shows how electron geometry (trigonal bipyramidal) differs from molecular geometry (linear)

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Geometries Derived from Octahedral: Symmetric Lone Pair Placement

In octahedral electron geometry, all six positions are initially equivalent. When lone pairs are added, they position themselves to maximize distance from each other—going 180° apart when possible.

Square Pyramidal

• Five bonding pairs + one lone pair creates a square base with one atom above (e.g., $\text{BrF}_5$)
• Bond angles of approximately 85°—compressed from 90° by the single lone pair
• First step in the octahedral → square planar progression as lone pairs are added

Square Planar

• Four bonding pairs + two lone pairs with lone pairs positioned 180° apart (e.g., $\text{XeF}_4$)
• Bond angles of exactly 90° between adjacent fluorines in the square plane
• Lone pairs occupy axial positions (above and below the plane) to minimize repulsion with each other

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Quick Reference Table

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Self-Check Questions

1. Both $\text{CO}_2$ and $\text{H}_2\text{O}$ have three atoms, yet one is linear and one is bent. Explain how lone pairs account for this difference and predict which molecule is polar.

2. Compare the bond angles in $\text{CH}_4$, $\text{NH}_3$, and $\text{H}_2\text{O}$. What trend do you observe, and what principle explains it?

3. Why do lone pairs in a trigonal bipyramidal electron geometry preferentially occupy equatorial rather than axial positions? Use $\text{SF}_4$ as your example.

4. Both $\text{XeF}_4$ (square planar) and $\text{CH}_4$ (tetrahedral) have four atoms bonded to a central atom. Explain why their geometries and bond angles differ.

5. An FRQ asks you to explain why $\text{BF}_3$ is nonpolar but $\text{NF}_3$ is polar, even though both have three fluorines bonded to a central atom. How would you structure your response using molecular geometry concepts?