VSEPR Theory and Molecular Shapes

Prerequisites

Before diving into VSEPR theory, review:

• Covalent Bonding — Lewis structures and electron pairs
• Lewis Dot Structures — Drawing molecular structures
• Electronic Configuration — Understanding valence electrons

The Hook: Why Water Isn’t Linear

The Core Concept

What is VSEPR Theory?

VSEPR Theory predicts the 3D shapes of molecules based on one simple idea:

In simple terms: Imagine balloons tied together at a central point. The balloons naturally spread out as far as possible to minimize bumping into each other. That’s exactly what electron pairs do around an atom!

The Fundamental Principle

Electron pairs (bonding + lone pairs) around a central atom arrange themselves to be as far apart as possible to minimize repulsion.

graph TD
    A[Count electron pairs
around central atom] --> B{Total electron pairs}
    B -->|2 pairs| C[Linear
180°]
    B -->|3 pairs| D[Trigonal Planar
120°]
    B -->|4 pairs| E[Tetrahedral
109.5°]
    B -->|5 pairs| F[Trigonal Bipyramidal
90°, 120°]
    B -->|6 pairs| G[Octahedral
90°]

    style A fill:#e1f5fe
    style C fill:#fff9c4
    style D fill:#f0f4c3
    style E fill:#dcedc8
    style F fill:#c5e1a5
    style G fill:#aed581

VSEPR Methodology: The 4-Step Process

Step 1: Draw Lewis Structure

Identify central atom, count valence electrons, distribute as bonds and lone pairs.

Step 2: Count Electron Pairs

Steric Number (SN) = Number of bonding pairs + Number of lone pairs

Important: Count multiple bonds as ONE bonding region (double = 1, triple = 1)

Step 3: Determine Electron Geometry

Based on total electron pairs (SN), determine how electron pairs arrange in space.

Step 4: Determine Molecular Geometry

Consider only atoms (ignore lone pairs) to determine molecular shape.

Repulsion Order: Who Pushes Harder?

Not all electron pairs repel equally! Lone pairs occupy more space than bonding pairs.

Where:

• LP = Lone Pair
• BP = Bonding Pair

Why? Lone pairs are held by only one nucleus (the central atom), so they spread out more. Bonding pairs are held by two nuclei, so they’re more compressed.

Effect on Bond Angles

More lone pairs → Smaller bond angles

Example: CH₄ vs NH₃ vs H₂O

Memory trick: “More Lone, Less Angle!” → Each lone pair reduces angle by ~2.5°

The Complete VSEPR Guide

2 Electron Pairs (SN = 2)

Linear Geometry (180°)

Electron Geometry: Linear Molecular Geometries: 1 shape possible

Type: AX₂ (2 BP, 0 LP) — LINEAR

Example: BeCl₂, CO₂, HCN

Cl-Be-Cl    or    O=C=O
    180°           180°

Visual Description:

• Linear: Like a straight rod
• Both atoms on opposite sides of central atom
• Bond angle = 180°

3 Electron Pairs (SN = 3)

Electron Geometry: Trigonal Planar (120°)

Type 1: AX₃ (3 BP, 0 LP) — TRIGONAL PLANAR

Example: BF₃, BCl₃, CO₃²⁻, NO₃⁻

     F
     |
 F—B—F
   120°

Visual Description:

• Trigonal Planar: Like a Mercedes-Benz logo
• All atoms in the same plane
• Three arms at 120° angles
• Flat, symmetrical

Example: BF₃ (boron trifluoride)

• B at center, 3 F atoms around it
• All in one plane, 120° apart
• Looks like three spokes of a wheel

Type 2: AX₂E (2 BP, 1 LP) — BENT/V-SHAPED

Example: SO₂, O₃, NO₂⁻

   O
  //
 S
  \
   O
  ~120° (actually ~119°)

Visual Description:

• Bent: Like a boomerang
• Two atoms bonded, one lone pair (invisible)
• Lone pair pushes atoms closer
• Bond angle < 120° (usually ~119° for SO₂)

4 Electron Pairs (SN = 4)

Electron Geometry: Tetrahedral (109.5°)

Type 1: AX₄ (4 BP, 0 LP) — TETRAHEDRAL

Example: CH₄, CCl₄, NH₄⁺, SO₄²⁻

       H
       |
   H—C—H
       |
       H
   109.5°

Visual Description:

• Tetrahedral: Like a triangular pyramid
• 4 atoms arranged at corners of a tetrahedron
• Central atom at the center
• All bond angles = 109.5°
• 3D structure (NOT flat!)

Imagine: A tripod with one leg pointing up and three pointing down (but more symmetrical). Or a soccer ball’s stitching at a vertex.

Type 2: AX₃E (3 BP, 1 LP) — TRIGONAL PYRAMIDAL

Example: NH₃, PCl₃, H₃O⁺

      N
     /|\
    H H H
    ~107°

Visual Description:

• Trigonal Pyramidal: Like a tripod or a stool
• 3 atoms form a triangular base
• Central atom above the base (not in the plane!)
• Lone pair at the top (invisible)
• Bond angle ≈ 107° (less than 109.5° due to LP repulsion)

Imagine: A three-legged camera tripod. The camera (central atom) sits above the three legs (bonded atoms). There’s an invisible fourth leg (lone pair) pointing up.

Type 3: AX₂E₂ (2 BP, 2 LP) — BENT/V-SHAPED

Example: H₂O, H₂S, SCl₂

    O:  (2 lone pairs)
   / \
  H   H
 104.5°

Visual Description:

• Bent: Like a boomerang or an open book
• 2 atoms bonded at an angle
• 2 lone pairs (invisible) push from opposite sides
• Bond angle ≈ 104.5° (H₂O), smaller than 109.5°

Imagine: Take a tetrahedral structure, remove two opposite corners (atoms), replace with lone pairs. The remaining two atoms form a bent shape.

5 Electron Pairs (SN = 5)

Electron Geometry: Trigonal Bipyramidal (90° and 120°)

Key Feature: TWO types of positions:

• Axial: 2 positions above and below (90° to equatorial)
• Equatorial: 3 positions in the middle plane (120° to each other)

Important: Lone pairs ALWAYS occupy equatorial positions first (less repulsion!)

Type 1: AX₅ (5 BP, 0 LP) — TRIGONAL BIPYRAMIDAL

Example: PCl₅, PF₅

      Cl (axial)
       |
Cl—P—Cl  (equatorial, 120° apart)
       |
      Cl (axial)

Visual Description:

• Trigonal Bipyramidal: Like two triangular pyramids joined at the base
• 3 atoms in the middle (equatorial) at 120°
• 2 atoms above and below (axial) at 90° to equatorial
• NOT all bond angles equal!

Imagine: A bow tie or two Egyptian pyramids stuck together base-to-base.

Type 2: AX₄E (4 BP, 1 LP) — SEE-SAW/DISTORTED TETRAHEDRAL

Example: SF₄, XeO₂F₂

      F (axial)
       |
    F—S—F (equatorial)
       |
      F (axial)
  (LP at equatorial)

Visual Description:

• See-Saw: Like a playground see-saw or teeter-totter
• Lone pair occupies one equatorial position
• 2 axial atoms (above/below)
• 2 equatorial atoms (one position empty due to LP)
• Asymmetrical, distorted

Bond angles: ~90° (axial-equatorial), ~120° (equatorial-equatorial), but distorted by LP

Imagine: A see-saw at a playground — two kids on the ends (axial), one kid on one side (equatorial), empty seat on other side (lone pair).

Type 3: AX₃E₂ (3 BP, 2 LP) — T-SHAPED

Example: ClF₃, BrF₃

      F
      |
    F—Cl—F
   (2 LP at equatorial)

Visual Description:

• T-Shaped: Like the letter “T”
• 2 lone pairs occupy equatorial positions
• 2 atoms at axial positions (up and down)
• 1 atom at remaining equatorial position
• Forms a T

Bond angles: ~90° (axial to equatorial), slightly less than 90° due to LP repulsion

Imagine: A capital letter T. The vertical part has atoms at top and bottom (axial), the horizontal part has one atom extending to the side (equatorial).

Type 4: AX₂E₃ (2 BP, 3 LP) — LINEAR

Example: XeF₂, I₃⁻

F—Xe—F
(3 LP at equatorial)

Visual Description:

• Linear: Straight line (like AX₂)
• 3 lone pairs occupy ALL equatorial positions
• 2 atoms at axial positions (least repulsion)
• Bond angle = 180°

Imagine: A barbell again, but with invisible balloons (lone pairs) in the middle plane pushing the atoms to opposite sides.

6 Electron Pairs (SN = 6)

Electron Geometry: Octahedral (90°)

Key Feature: All 6 positions are equivalent (unlike TBP)

Type 1: AX₆ (6 BP, 0 LP) — OCTAHEDRAL

Example: SF₆, PF₆⁻, [Fe(CN)₆]³⁻

      F
      |
  F—S—F
      |
      F
(+ 1 F front, 1 F back)

Visual Description:

• Octahedral: Like two square pyramids joined at the base
• 6 atoms at corners of an octahedron
• All bond angles = 90°
• Highly symmetrical

Imagine: A die (6-sided cube) with atoms at each face’s center. Or two Egyptian pyramids (square base) stuck together.

Type 2: AX₅E (5 BP, 1 LP) — SQUARE PYRAMIDAL

Example: BrF₅, IF₅, [NiCN₅]³⁻

      F
      |
  F—Br—F
   | |
    F—F
(LP opposite to apex)

Visual Description:

• Square Pyramidal: Like the Egyptian pyramids
• 4 atoms form a square base
• 1 atom at the apex (top)
• Lone pair at the bottom (opposite to apex)

Bond angles: ~90° (mostly), slightly distorted by LP

Imagine: The Great Pyramid of Giza — square base, apex at top, lone pair hidden underground.

Type 3: AX₄E₂ (4 BP, 2 LP) — SQUARE PLANAR

Example: XeF₄, ICl₄⁻, [PtCl₄]²⁻

    F
    |
F—Xe—F
    |
    F
(2 LP above and below plane)

Visual Description:

• Square Planar: Like a flat square
• 4 atoms at corners of a square
• 2 lone pairs above and below the plane
• All atoms in one plane
• Bond angles = 90°

Imagine: A perfectly flat piece of paper with atoms at the four corners. Lone pairs are ghosts floating above and below the paper.

Summary Table: All VSEPR Shapes

Interactive Demo: Explore VSEPR Molecular Geometries

Use this interactive 3D visualization to explore all VSEPR geometries. Select preset molecules like CH4, NH3, H2O, or customize electron pair counts. Drag to rotate the model and observe how lone pairs affect bond angles. Watch the difference between electron geometry and molecular geometry in real-time!

Memory Tricks & Patterns

The Lone Pair Effect

“Each lonely pair steals 2.5° from the party!”

For tetrahedral base (109.5°):

• 0 LP: CH₄ = 109.5°
• 1 LP: NH₃ = 107° (−2.5°)
• 2 LP: H₂O = 104.5° (−5°)

Shape Family Tree

2 EP: Linear only
3 EP: Planar (flat) shapes only
4 EP: 3D shapes start
5 EP: Complex shapes, LP prefer equatorial
6 EP: 3D symmetrical shapes

Visual Mnemonics

BeCl₂: Becomes Linear (180°) BF₃: Boron Forms Trigonal planar (120°) CH₄: Carbon Has Tetrahedral (109.5°) NH₃: Nitrogen’s Pyramid (107°) H₂O: Highly Bent (104.5°) PCl₅: Phosphor BP shape (Bipyramidal) SF₆: Sulfur Octahedral (6 faces)

Special Cases and Exceptions

Multiple Bonds Count as ONE Region

Expanded Octets (Period 3 and Beyond)

Elements in Period 3 and higher can accommodate more than 8 electrons using vacant d-orbitals.

Examples:

• PCl₅: P has 10 electrons (5 bonds)
• SF₆: S has 12 electrons (6 bonds)
• IF₇: I has 14 electrons (7 bonds)

Why? They have accessible 3d, 4d, 5d orbitals for electron accommodation.

Resonance Doesn’t Change Shape

Example: CO₃²⁻

• Has 3 resonance structures (double bond rotates)
• But VSEPR shape remains trigonal planar
• Why? Average bond order doesn’t change electron regions

Common Mistakes to Avoid

Practice Problems

Level 1: Foundation (NCERT Style)

Level 2: JEE Main Type

Level 3: JEE Advanced Type

Quick Revision Box

Most Tested Molecules in JEE

Real-World Applications

Teacher’s Summary

Related Topics

Within Chemical Bonding

• Covalent Bonding - Lewis structures, sigma and pi bonds
• Hybridization - Why these shapes exist (VBT explanation)
• Molecular Orbital Theory - Advanced bonding perspective
• Ionic Bonding - Comparison with ionic structures
• Hydrogen Bonding - Shape and polarity effects

Atomic Structure Foundation

• Orbital Shapes - s, p, d orbital geometry
• Electronic Configuration - Valence electrons

Organic Chemistry Applications

• Organic Principles - Hybridization - Carbon geometry in organic molecules
• Reaction Intermediates - Carbocation (sp2) and carbanion shapes
• Aldehydes and Ketones - Carbonyl geometry

Coordination Chemistry

• Coordination Compounds - Complex ion geometry
• Crystal Field Theory - d-orbital splitting

Inorganic Applications

• p-Block Elements - Oxoacid structures
• Group 15 - NH3, PCl5 shapes
• Group 16 - H2O, H2S shapes

Physics Connection

• Electrostatics - Charge repulsion basics
• Electric Dipole - Molecular polarity

Mathematics

• Vector Algebra - Dipole moment vector addition