Advanced Theories of Chemical bonding

CHEMISTRY - ADVANCED THEORIES OF COVALENT BONDING

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CHAPTER: ADVANCED THEORIES OF COVALENT BONDING

Major Theories Covered

  • Valence Bond Theory
  • Hybrid Atomic Orbitals
  • Multiple Bonds
  • Molecular Orbital Theory

BONDING THEORIES

Overview of Key Concepts

Lewis Theory and VSEPR

  • Dots arrayed about a central atom represent bonds and lone pairs of electrons.
  • Arrangement of atoms in three-dimensions helps predict molecular geometry.
  • Molecular polarity can be predicted from molecular structure.

Valence Bond Theory

  • Electrons reside within orbitals on each atom.
  • Orbitals may be hybrids, with overlapping orbitals in space forming bonds.

Molecular Orbital Theory

  • Electrons exist in orbitals that are spread over the entire molecule, not attached to individual atoms.
  • This model accounts for experimental observations of magnetism that simpler theories cannot.

VALENCE BOND THEORY

Weaknesses of VSEPR Theory

  • VSEPR is a model that predicts three-dimensional molecular shapes.
  • Bond angles in VSEPR models, like 109° and 120°, do not correspond to geometries of s, p, and d atomic orbitals.

Overview of Valence Bond Theory

  • Covalent bonds are formed via the overlap of atomic orbitals containing single electrons, creating electron pairs.
  • Orbitals overlap when portions of atomic orbitals occupy the same region of space:
    • Overlap of two s-orbitals forms a single bond.
    • Overlap of one s-orbital and one p-orbital forms a single bond.
    • Overlap of two p-orbitals forms a single bond.

ENERGY AND ORBITAL OVERLAP

Orbital Overlap Process

  • The energy of a system is influenced by how orbitals overlap.
  • No overlap at large distances results in zero energy.
  • As orbitals approach, the attraction of one nucleus to its electrons increases while repulsions also start to act.
  • Initially, attractive forces exceed repulsive forces, causing energy to decrease as a bond starts forming.
  • As overlap increases, both attractions and repulsions escalate until energy reaches a minimum at an optimal distance, stabilizing the bond.

Bond Energy and Bond Length

  • Bond energy, ΔE, is defined as the difference between the minimum energy (at bond length) and fully separated atoms.
  • It represents energy required to completely separate atoms or the energy released upon bond formation.
  • Example: 7.24 x 10⁻¹⁹ J required to break one H-H bond.

Distance and Orientation in Bond Formation

  • Formation of a bond is dependent on both the distance between atomic nuclei and the orientation of the overlapping orbitals.
  • Ideal orbital orientation is along the line connecting two nuclei, maximizing overlap.
  • Overlap between two s-orbitals is simple due to spherical symmetry. For p-orbitals, proper orientation is critical for effective overlap.

SIGMA (σ) AND PI (π) BONDS

Bonding Characteristics

  • Sigma (σ) bonds result from overlaps including:
    • Two s-orbitals
    • One s-orbital and one p-orbital
    • Two p-orbitals
  • Pi (π) bonds arise from side-by-side overlaps of two p-orbitals.

Bonds in Molecular Structures

  • Single bonds are always σ-bonds.
  • Multiple bonds consist of σ-bonds and π-bonds:
    • Double bond = one σ-bond and one π-bond.
    • Triple bond = one σ-bond and two π-bonds.
  • Example: Butadiene, C₆H₆, is utilized in synthetic rubber and included as a case study.

HYBRID ATOMIC ORBITALS

Orbital Overlap in Water

  • Water (H₂O) features a central oxygen atom joined to two hydrogen atoms.
  • Electron configuration of oxygen: 1s² 2s² 2p⁴
  • A simple overlap model predicts 90° bond angle, conflicting with experimental evidence showing an actual angle of 104.9°.

Implications from VSEPR Theory

  • Four regions of electron density around oxygen dictate a tetrahedral classification.
  • Despite two lone pairs of electrons, water adopts a bent molecular geometry.
  • Incorporating hybridization into valence bond theory provides an accurate framework for bonding.

Characteristics of Hybrid Orbitals

  • Hybrid orbitals only exist when atoms bond, not in isolation.
  • These orbitals derive from the combination of atomic orbitals, maintaining equivalency in shape and energy within a hybridized set.
  • The number of hybrid orbitals created is equivalent to the number of atomic orbitals mixed.
  • Orientation and shape depend on electron-pair geometries and VSEPR theory, with σ-bonds formed from hybrid orbitals and π-bonds from unhybridized ones.

Types of Hybridization

sp Hybridization

  • One valence s-orbital combines with one valence p-orbital to form two equivalent sp hybrid orbitals.

sp² Hybridization

  • A central atom with three regions of electron density shapes a set of three sp² hybrid orbitals, forming a trigonal planar configuration at 120°.

sp³ Hybridization

  • For four regions of electron density, four sp³ hybrid orbitals create a tetrahedral configuration.
  • Example: Ethane (C₂H₆) features sp³ hybridized carbons, with rotation about the σ-bond readily occurring.

Higher Hybridizations

  • sp³d hybridization involves five bonding orbitals: one s, three p, and one d orbital, forming a trigonal bipyramidal shape.
  • sp³d² hybridization involves six bonding orbitals: one s, three p, and two d orbitals, leading to an octahedral shape.

MULTIPLE BONDS IN MOLECULES

Example Structures

Ethene (C₂H₄)

  • Each carbon surrounded by two hydrogen atoms and another carbon, requiring sp² hybridization.
  • Forms one σ-bond and one π-bond between carbon atoms, totaling five σ-bonds and one π-bond in C₂H₄.

Acetylene (C₂H₂)

  • Each carbon atom coupled with one hydrogen and another carbon, necessitating sp hybridization.
  • Results in one σ-bond and one π-bond between the carbon atoms, along with σ-bonds connecting to the single hydrogen atoms. Total: three σ-bonds and two π-bonds

MOLECULAR ORBITALS

Orbital Diagrams and Mixing

  • Comparison of molecular orbitals for B₂, C₂, N₂ (significant s-p mixing) versus O₂, F₂, Ne₂ (less s-p mixing).
  • Bonding and antibonding interactions illustrated in diagrams.

Paramagnetism in O₂

  • Experimental evidence shows O₂ is attracted to magnets, whereas N₂ is not.
  • Contrasting Lewis dot structures for N₂ and O₂ with respect to unpaired electrons leads to differences in magnetic properties.

Bond Order and Stabilization

  • Electron pairing is preferable for stability but unpaired