| Term | Meaning |
|---|---|
| Electron Domain | Any region of electron density: single bond, double bond, triple bond, or lone pair |
| Steric Number (SN) | Total electron domains = bonding domains + lone pairs |
| Electron Geometry | Arrangement of all electron domains (including lone pairs) |
| Molecular Geometry | Arrangement of atoms only (lone pairs excluded from name) |
Step 1 — Draw the Lewis structure.
Step 2 — Count electron domains around the central atom.
Step 3 — Find the Steric Number (SN) → gives Electron Geometry.
Step 4 — Count lone pairs → gives Molecular Geometry.
Step 5 — Predict bond angles (lone pairs compress them).
X ─── A ─── X Bond Angle: 180°
X
/
X ─── A Bond Angle: 120°
\
X| Lone Pairs | Molecular Geometry | Bond Angle | Example |
|---|---|---|---|
| 0 | Trigonal Planar | 120° | BF₃, SO₃ |
| 1 | Bent (V-shaped) | <120° | SO₂, O₃ |
X
/
X ─── A ─── X Bond Angle: 109.5°
\
X| Lone Pairs | Molecular Geometry | Bond Angle | Example |
|---|---|---|---|
| 0 | Tetrahedral | 109.5° | CH₄, CCl₄ |
| 1 | Trigonal Pyramidal | ~107° | NH₃, PCl₃ |
| 2 | Bent (V-shaped) | ~104.5° | H₂O, H₂S |
Why angles compress: Each lone pair pushes bonding pairs closer.
NH₃: one LP → 109.5° → 107° | H₂O: two LP → 109.5° → 104.5°
X (axial)
|
X ─────── A ───────X Axial angle: 90°
/ \ Equatorial angle: 120°
X X
(axial)Two types of positions:
Rule: Lone pairs always occupy equatorial positions (fewer 90° close contacts = less repulsion).
| Lone Pairs | Molecular Geometry | Example |
|---|---|---|
| 0 | Trigonal Bipyramidal | PCl₅ |
| 1 | Seesaw | SF₄ |
| 2 | T-shaped | ClF₃ |
| 3 | Linear | XeF₂, I₃⁻ |
X (axial)
|
X ─── X ─── A ─── X Bond Angle: 90°
|
X (axial)All 6 positions are equivalent.
Two lone pairs prefer trans (180° apart) positions to maximise separation.
| Lone Pairs | Molecular Geometry | Example |
|---|---|---|
| 0 | Octahedral | SF₆ |
| 1 | Square Pyramidal | BrF₅, IF₅ |
| 2 | Square Planar | XeF₄, ICl₄⁻ |
| SN | Electron Geometry | Lone Pairs | Molecular Geometry | Bond Angle | Example |
|---|---|---|---|---|---|
| 2 | Linear | 0 | Linear | 180° | CO₂ |
| 3 | Trigonal Planar | 0 | Trigonal Planar | 120° | BF₃ |
| 3 | Trigonal Planar | 1 | Bent | <120° | SO₂ |
| 4 | Tetrahedral | 0 | Tetrahedral | 109.5° | CH₄ |
| 4 | Tetrahedral | 1 | Trigonal Pyramidal | ~107° | NH₃ |
| 4 | Tetrahedral | 2 | Bent | ~104.5° | H₂O |
| 5 | Trig. Bipyramidal | 0 | Trig. Bipyramidal | 90°/120° | PCl₅ |
| 5 | Trig. Bipyramidal | 1 | Seesaw | ~90°/~120° | SF₄ |
| 5 | Trig. Bipyramidal | 2 | T-shaped | ~90° | ClF₃ |
| 5 | Trig. Bipyramidal | 3 | Linear | 180° | XeF₂ |
| 6 | Octahedral | 0 | Octahedral | 90° | SF₆ |
| 6 | Octahedral | 1 | Square Pyramidal | ~90° | BrF₅ |
| 6 | Octahedral | 2 | Square Planar | 90° | XeF₄ |
Lone pairs are not shown in the molecular geometry name, but they profoundly determine the shape.
| Molecule | Lone Pairs | Bond Angle | Change |
|---|---|---|---|
| CH₄ | 0 | 109.5° | Baseline |
| NH₃ | 1 | 107° | −2.5° |
| H₂O | 2 | 104.5° | −5° |
Each lone pair compresses angles by ~2–2.5°.
Multiple bonds act as one domain but carry more electron density → repel more.
Example: Formaldehyde (H₂C=O)
General rule:
More electronegative groups pull bonding electrons away from the central atom → those pairs occupy less space → repel less → angles compress.
Example:
In NF₃, the N–F bonding pairs are drawn toward F, reducing their repulsion. The lone pair on N dominates, compressing the angle further.
This distinction is frequently tested:
| Electron Geometry | Molecular Geometry | |
|---|---|---|
| Based on | All domains (LP + BP) | Atoms only |
| Lone pairs included? | Yes | No |
| Determined by | Steric Number | SN minus lone pairs |
| H₂O example | Tetrahedral (SN=4) | Bent (2 LP excluded) |
Geometry determines whether bond dipoles cancel:
| Molecule | Geometry | Polarity |
|---|---|---|
| CO₂ | Linear | Nonpolar |
| H₂O | Bent | Polar |
| BF₃ | Trigonal Planar | Nonpolar |
| NH₃ | Trigonal Pyramidal | Polar |
| CH₄ | Tetrahedral | Nonpolar |
| XeF₄ | Square Planar | Nonpolar |
VSEPR theory delivers a powerful prediction from a single idea: electron pairs repel and spread apart. The steric number sets the electron geometry; lone pairs then sculpt the molecular geometry by compressing angles. While qualitative, it is an indispensable first step from Lewis structures toward a full quantum mechanical picture of bonding.
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VSEPR Theory: Molecular Shapes, Postulates & Geometry Notes PDF is a fundamental concept in inorganic chemistry. Understanding the mechanisms, reaction conditions, and stereo-chemical outcomes is crucial for mastering organic chemistry. Our curated resources provide step-by-step visualizations to help you excel.
SELF TEST
According to VSEPR theory, what determines the shape of a molecule?
LEARNING SUPPORT
Lone pairs are attracted to only one nucleus (the central atom), so they occupy more space and spread out more than bonding pairs, which are pulled between two nuclei. This extra space results in stronger electrostatic repulsion against neighboring pairs.
Electron Geometry considers the arrangement of all electron domains (both bonding and lone pairs). Molecular Geometry only describes the arrangement of the actual atoms, though its shape is determined by the positions of the lone pairs.
VSEPR theory provides very good estimates (like 109.5° for tetrahedral), but exact angles depend on the specific atoms involved and the difference in repulsion between lone pairs and bonding pairs (e.g., water is 104.5° instead of 109.5°).