VSEPR Theory, Valence Bond Theory, and Hybridization: Molecular Shapes and Bonding

VSEPR Theory, Valence Bond Theory, and Hybridization

VSEPR Theory – Predicting Molecular Shape

The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the three-dimensional geometry of molecules based on the repulsion between electron pairs (both bonding and lone pairs) around a central atom. Electron pairs arrange themselves to minimize repulsion, resulting in specific molecular shapes.

• Key Principle: Electron pairs repel each other and adopt positions as far apart as possible.

• Limitations: VSEPR does not explain why bonds form, nor does it account for bond strengths or bond lengths.

• Example: In H2S, the bond angle is 92°, less than the ideal 109.5° for four electron pairs. This deviation is explained by the use of pure p-orbitals and repulsion effects.

Valence Bond (VB) Theory – Why Bonds Form

Valence Bond Theory describes a covalent bond as the overlap of atomic orbitals from two atoms, with two electrons of opposite spin occupying the overlap region. This overlap lowers the system's potential energy, stabilizing the molecule.

• Bond Formation: Occurs when singly occupied atomic orbitals on two atoms overlap.

• Electron Spins: The two electrons in the bond must have opposite spins.

• Energy Consideration: Overlap brings nuclei closer, lowering potential energy.

• Examples:

  • HF: 1s orbital of H overlaps with 2p orbital of F.

  • F2: Two 2p orbitals overlap end-to-end.

• Limitation: Pure atomic orbitals cannot explain certain molecular shapes (e.g., linear BeCl2 or 180° bond angle in CO2).

Hybridization – Blending Orbitals to Match Geometry

Hybridization is the process of mixing atomic orbitals (s, p, and sometimes d) to form new, equivalent hybrid orbitals that match the observed geometry of molecules. This concept explains how atoms form the correct number of bonds and achieve the correct spatial arrangement.

Why Hybridization is Needed

• Atoms often do not have enough unpaired electrons in their ground state to form the required number of bonds.

• Electrons are promoted and orbitals are hybridized to create the necessary number of equivalent bonding orbitals.

Determining Hybridization: Step-by-Step

1. Draw the Lewis dot structure of the molecule.

2. Use VSEPR to determine the number of electron domains (bonding and lone pairs) around the central atom.

3. Match the number of electron domains to the hybridization scheme:

Electron Domains	Hybridization	Electron Geometry
2	sp	Linear
3	sp2	Trigonal planar
4	sp3	Tetrahedral
5	sp3d	Trigonal bipyramidal
6	sp3d2	Octahedral

The sp, sp2, sp3 Hybrids – How They Form

sp Hybridization (Linear, 180°)

• Example: BeH2

• Be ground state: [He]2s2 (no unpaired electrons)

• Promote one 2s electron to 2p: 2s12p1

• Mix one 2s and one 2p orbital → two equivalent sp hybrid orbitals

• sp orbitals point 180° apart (linear geometry)

• Each sp orbital overlaps with H 1s to form a sigma bond

sp2 Hybridization (Trigonal Planar, 120°)

• Example: BF3

• B ground state: 2s22p1 (only one unpaired electron)

• Promote one 2s electron to 2p: 2s12p2

• Mix one s and two p orbitals → three sp2 hybrid orbitals in a plane, 120° apart

• One p orbital remains unhybridized (perpendicular to the plane)

• Each sp2 overlaps with F 2p orbital (sigma bonds)

sp3 Hybridization (Tetrahedral, 109.5°)

• Example: CH4

• C ground state: 2s22p2 (two unpaired electrons)

• Promote one 2s electron to 2p: 2s12p3 (four unpaired electrons)

• Mix one s and three p orbitals → four sp3 hybrid orbitals

• sp3 orbitals point to corners of a tetrahedron (109.5° apart)

• Each overlaps with H 1s to form sigma bonds

• Other sp3 examples:

  • NH3: Three sp3 orbitals form N–H bonds, one holds a lone pair. Bond angle ~107.3°.

  • H2O: Two sp3 orbitals form O–H bonds, two hold lone pairs. Bond angle ~105°.

Hybridization Involving d Orbitals (3rd Period and Beyond)

When the central atom forms more than four bonds, d orbitals participate in hybridization.

• sp3d: Five hybrid orbitals, trigonal bipyramidal geometry

  • Example: PBr5

• sp3d2: Six hybrid orbitals, octahedral geometry

  • Example: SF6

Multiple Bonds – Sigma (σ) and Pi (π) Bonds

After forming the sigma bond framework with hybrid orbitals, any remaining unhybridized p orbitals can form pi bonds. The combination of sigma and pi bonds determines the bond order and restricts rotation in double and triple bonds.

• Single bond: 1 sigma (σ) bond (end-to-end overlap)

• Double bond: 1 σ + 1 pi (π) bond (sideways overlap of p orbitals)

• Triple bond: 1 σ + 2 π bonds

Molecule	Hybridization of Central Atom	Bonding Description
C2H4 (ethylene)	C is sp2	σ framework from sp2 orbitals; leftover p electrons form one π bond between carbons. Restricts rotation (cis/trans isomers).
C2H2 (acetylene)	C is sp	σ bonds from sp hybrids; two unhybridized p orbitals on each C form two perpendicular π bonds (triple bond).
CH2O (formaldehyde)	C is sp2	C forms σ bonds to two H and one O (via sp2-sp2 overlap). Unhybridized p on C and O form a π bond. O also has two lone pairs in sp2 orbitals.

Summary – Connecting VSEPR and VB/Hybridization

1. VSEPR determines the number of electron domains and their spatial arrangement.

2. The electron domain number indicates the hybridization type (sp, sp2, sp3, sp3d, sp3d2).

3. Valence Bond Theory explains sigma bond formation from hybrid orbitals and the resulting geometry.

4. Remaining p orbitals form π bonds, determining bond order and restricting rotation in multiple bonds.

5. Observed bond angles are often slightly less than ideal when lone pairs are present, due to greater lone pair–bonding pair repulsion compared to bonding pair–bonding pair repulsion.

Additional info: The concepts of VSEPR, hybridization, and sigma/pi bonding are foundational for understanding molecular geometry, bond strength, and reactivity in general chemistry. Mastery of these topics is essential for predicting molecular properties and behavior.