Chapter 11

Introduction to Covalent Bonding

• Learning changes everything.

• Textbook: Chemistry: The Molecular Nature of Matter and Change, Tenth Edition by Martin S. Silberberg and Patricia G. Amateis.

• Chapter 11: Theories of Covalent Bonding

  • 11.1: Valence Bond (VB) Theory and Orbital Hybridization

  • 11.2: Modes of Orbital Overlap and the Types of Covalent Bonds

  • 11.3: Molecular Orbital (MO) Theory and Electron Delocalization

Valence Bond (VB) Theory

• Basic Principle of VB Theory:

  • A covalent bond is formed when the orbitals of two atoms overlap, and a pair of electrons occupy the overlap region.

  • The overlapping space can accommodate a maximum of two electrons, which must have opposite (paired) spins.

  • Stronger covalent bonds result from greater orbital overlap; the closer the nuclei are to the electrons, the stronger the bond.

  • The extent of orbital overlap depends on the shape and direction of the orbitals.

  • Hybridization of atomic orbitals is possible in some covalent compounds.

Orbital Orientation and Maximum Overlap

• Overlapping orbitals must align spatially to maximize interaction.

• Figures Illustrated:

  • H₂, HF, and F₂ showing different overlapping patterns of 1s and 2p orbitals.

VB Theory and Orbital Hybridization

• In certain molecules, valence orbitals from isolated atoms can mix to form new hybrid orbitals in the molecule, a process known as hybridization.

• New orbitals formed (hybrid orbitals) have spatial orientations that correspond with observed molecular shapes.

• Without hybridization, some observed shapes cannot be accounted for.

Features of Hybrid Orbitals

• The number of formed hybrid orbitals is equal to the number of atomic orbitals that are mixed.

• The types of hybrid orbitals vary according to the atomic orbitals involved in the mixing.

• Hybrid orbitals are oriented to maximize overlap with orbitals from the bonding atom.

Examples of Hybrid Orbitals

sp Hybrid Orbitals in BeCl₂

• Illustration:

  • Simple depiction of sp hybridization with linear orientation (-180°).

  • Be atom undergoes hybridization of one 2s and one 2p orbital.

sp² Hybrid Orbitals in BF₃

• Illustration:

  • Depicts mixing of two 2p orbitals and one 2s to yield three sp² hybrid orbitals for boron.

sp³ Hybrid Orbitals in CH₄, NH₃, H₂O

• CH₄ (Methane):

  • Carbon atom hybridizes to form four sp³ orbitals leading to tetrahedral molecular geometry.

• NH₃ (Ammonia):

  • Nitrogen has one lone pair and three bonding pairs, forming a trigonal pyramidal shape.

• H₂O (Water):

  • Involves hybridization in oxygen with two lone pairs leading to a bent molecular geometry.

sp³d and sp³d² Hybrid Orbitals

• PCl₅:

  • Involves mixing of d orbitals with p and s orbitals, leading to trigonal bipyramidal geometry.

• SF₆:

  • Involves hybridization to form six sp³d² orbitals resulting in octahedral shape.

Composition and Orientation of Hybrid Orbitals

• Summary table of hybridization types:

  • Information Layout:

  • Atomic orbitals mixed

  • Hybrid orbitals formed

  • Unhybridized orbitals remaining

  • Orientation of hybrid orbitals (e.g., linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral).

From Molecular Formula to Hybrid Orbitals

• Create hybrid orbital diagrams from molecular formulas, through steps involving Lewis structures and determining electron group arrangements.

Sample Problem 11.1: Hybrid Orbitals in Compounds

Problem Statement

• Use partial orbital diagrams to describe hybridization in methanol (CH₃OH) and sulfur tetrafluoride (SF₄).

Solution Strategy

1. Draw Lewis structure to identify electron groups.

2. Use electron-group arrangements to determine hybrid orbitals.

Methanol (CH₃OH)

• Lewis Structure: Tetrahedral arrangement around C and O atoms; both have sp³ hybridized orbitals.

• Carbon forms four bonds (four half-filled sp³ orbitals); oxygen uses two sp³ orbitals for bonding, and two are filled with lone pairs.

Sulfur Tetrafluoride (SF₄)

• Lewis Structure: Sulfur has four bonds and one lone pair leading to a trigonal bipyramidal arrangement.

• The hybridization is sp³d, with one hybrid orbital occupied by a lone pair and four half-filled.

Limitations of the Hybridization Model

• Hybridization does not consistently match observed molecular shapes, especially in larger nonmetal hydrides (e.g., H₂S showing a 92° bond angle).

• In expanded valence shells, d-orbital contributions may not effectively hybridize.

Types of Covalent Bonds

• Sigma (σ) Bonds:

  • Formed by end-to-end overlap of orbitals; characterized by highest electron density along bond axis.

  • Every single bond is a σ bond.

• Pi (π) Bonds:

  • Formed by sideways overlap of orbitals; generally weaker than σ bonds, resulting in two regions of electron density above and below the bond axis.

  • A double bond consists of one σ bond and one π bond, and a triple bond consists of one σ and two π bonds.

Molecular Orbital (MO) Theory

• MO theory views a molecule as comprising molecular orbitals (MOs) that extend over the entire molecule occupied by delocalized electrons.

• Combination of wave functions from atomic orbitals forms molecular wave functions (MOs).

• Bonding MOs: Increased electron density between nuclei; Antibonding MOs: Nodes present leading to zero electron density between nuclei.

Molecular Orbital Diagrams

• MOs are filled similar to AOs in order of energy.

• Bond Order Calculation:
  ext{Bond Order} = rac{1}{2} ( ext{Number of bonding electrons} - ext{Number of antibonding electrons})

Sample Problems: Using MO Theory

• Predict stability and bond order of various species (e.g., N₂, O₂) using MO diagrams to determine bonding characteristics.

Conclusions

• Understanding hybridization, types of bonds, and molecular orbital theory is crucial for explaining molecular structure and behavior in chemical bonding.

• These concepts underlie much of organic and inorganic chemistry and are essential for advanced studies in chemical interactions.