Valence Bond Theory, Hybridization, and Bond Types (Chapter 7 Study Notes)

Valence Bond Theory

Introduction to Valence Bond Theory

Valence Bond Theory explains how atomic orbitals of atoms overlap to form chemical bonds in molecules. It provides a quantum mechanical description of bonding, focusing on the interactions between electrons in atomic orbitals.

• Atomic orbitals combine to form bonds when atoms approach each other.

• The overlap region between orbitals is where the bond forms, maximizing electron density between nuclei.

• Bond strength is related to the extent of orbital overlap.

• Valence Bond Theory is foundational for understanding molecular geometry and reactivity.

• Example: The H2 molecule forms when two hydrogen 1s orbitals overlap.

Hybridization

Concept of Hybridization

Hybridization is the process by which atomic orbitals mix to form new, equivalent hybrid orbitals that are oriented to maximize bonding in molecules. This concept explains the observed shapes of molecules that cannot be described by simple atomic orbitals.

• Hybrid orbitals are formed by combining s, p, and sometimes d orbitals.

• Hybridization accounts for molecular geometries such as tetrahedral, trigonal bipyramidal, and octahedral.

• Example: Carbon in methane (CH4) forms four sp3 hybrid orbitals, resulting in a tetrahedral geometry.

Hybridization in Specific Elements

• Beryllium (Be): In its ground state, Be has filled 2s orbitals and empty 2p orbitals. To form two bonds, Be promotes an electron from 2s to 2p, then hybridizes to form two sp orbitals.

• Boron (B): Boron promotes an electron and hybridizes one 2s and two 2p orbitals to form three sp2 hybrid orbitals.

• Carbon (C): Carbon promotes an electron and hybridizes one 2s and three 2p orbitals to form four sp3 hybrid orbitals.

• Expanded Octets: Elements in period 3 and beyond can hybridize d orbitals, forming sp3d (five orbitals, trigonal bipyramidal) or sp3d2 (six orbitals, octahedral) hybrid orbitals.

Relationship Between Geometry and Hybridization

The electron-domain geometry of a molecule determines the hybridization state of the central atom.

Electron-Domain Geometry	Hybridization	Example
Linear	sp	BeCl2
Trigonal planar	sp2	BCl3
Tetrahedral	sp3	CH4
Trigonal bipyramidal	sp3d	PCl5
Octahedral	sp3d2	SF6

Practice: Determining Hybridization

• (a) NH2-: sp3 hybridization (tetrahedral electron domain geometry)

• (b) SF4: sp3d hybridization (trigonal bipyramidal geometry)

• (c) SO32-: sp3 hybridization (tetrahedral electron domain geometry)

• (d) SF6: sp3d2 hybridization (octahedral geometry)

Bond Types

Sigma (σ) Bonds

Sigma bonds are the strongest type of covalent bond, formed by head-to-head overlap of orbitals. They have cylindrical symmetry about the internuclear axis.

• Formed by: Overlap of s-s, s-p, or p-p orbitals along the axis connecting two nuclei.

• All single bonds are sigma bonds.

• Example: H–H bond in H2 is a sigma bond.

Pi (π) Bonds

Pi bonds are formed by side-to-side overlap of p orbitals, with electron density above and below the internuclear axis.

• Formed by: Overlap of unhybridized p orbitals.

• Present in: Double and triple bonds (in addition to a sigma bond).

• Example: The C=C bond in ethylene (C2H4) contains one sigma and one pi bond.

Single and Multiple Bonds

• Single bonds: Always sigma bonds due to greater overlap and bond strength.

• Double bonds: Consist of one sigma and one pi bond.

• Triple bonds: Consist of one sigma and two pi bonds.

• Example:

  • H–H: one sigma bond

  • H2C=CH2: one sigma, one pi bond

  • N≡N: one sigma, two pi bonds

Delocalized Electrons and Resonance

Resonance Structures

Resonance occurs when electrons are delocalized over two or more atoms, resulting in multiple valid Lewis structures. The true electronic structure is a hybrid of these forms.

• Delocalized electrons are not confined to a single bond or atom.

• Resonance structures are depicted with double-headed arrows between them.

• Example: Nitrate ion (NO3-) has three resonance structures with delocalized pi electrons.

Resonance in Benzene

Benzene (C6H6) is a classic example of resonance, with six pi electrons delocalized over six carbon atoms, forming a ring of electron density.

• Structure: Alternating single and double bonds in a hexagonal ring.

• Delocalization: All C–C bonds are equivalent due to electron delocalization.

Practice: Describing Bonds in Molecules

• Formaldehyde (H2CO): The C–H bonds are sigma bonds formed by overlap of sp2 hybrid orbitals of carbon with 1s orbitals of hydrogen. The C=O bond consists of a sigma bond (sp2 of carbon with sp2 of oxygen) and a pi bond (unhybridized p orbitals).

• Acetonitrile (CH3CN):

  • Bond angles: 109.5° around CH3 carbon (sp3), 180° around CN carbon (sp).

  • Hybridization: CH3 carbon is sp3, CN carbon is sp.

  • Total bonds: 5 sigma, 2 pi bonds.

• Nitrate ion (NO3-): Has delocalized pi bonds due to resonance.

• Delocalized bonding: SO3, SO32-, H2CO, O3 exhibit delocalized bonding; NH4+ does not.

Summary Table: Hybridization and Geometry

Atomic Orbital Set	Hybrid Orbital Set	Geometry	Examples
1 s, 1 p	2 sp	Linear	BeCl2, CO2
1 s, 2 p	3 sp2	Trigonal planar	BCl3, SO3
1 s, 3 p	4 sp3	Tetrahedral	CH4, NH3, H2O
1 s, 3 p, 1 d	5 sp3d	Trigonal bipyramidal	PCl5, SF4
1 s, 3 p, 2 d	6 sp3d2	Octahedral	SF6, PF6-

Additional info: These notes expand on the original slides by providing definitions, examples, and context for hybridization, bond types, and resonance, ensuring a self-contained study guide for General Chemistry students.