Ch.6 - Chemical Bonding II

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR Theory is used to predict the geometry of molecules based on the repulsion between electron groups around a central atom. The arrangement of these groups determines the molecular shape.

• Electron Groups: Defined as the sum of bonding pairs (shared electrons in bonds) and lone pairs (non-bonding electrons) around the central atom.

• Lone Pair Electrons: Occupy more space than bonding pairs, increasing electron repulsion and affecting molecular geometry.

• Application: By counting the number of electron groups (bonding and lone pairs), the geometry can be predicted (e.g., linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral).

• Example: Ammonia (NH3): Nitrogen has 4 electron groups (3 bonding pairs, 1 lone pair), resulting in a tetrahedral electron geometry but a trigonal pyramidal molecular shape.

• Example: Selenium tetrafluoride (SeF4): Selenium has 5 electron groups (4 bonding pairs, 1 lone pair), resulting in a seesaw molecular shape.

Hybridization

Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals that can form stronger bonds and increase molecular stability.

• Concept: Atomic orbitals (s, p, d) combine to form hybrid orbitals (sp, sp2, sp3, sp3d, sp3d2).

• Relationship to Electron Groups: The number of electron groups around a central atom determines the type of hybridization:

• Unhybridized Orbitals: Orbitals not involved in hybridization remain as pure atomic orbitals (e.g., d orbitals in sp3 hybridization).

• Example: In SBr4, sulfur has 5 electron groups, so its hybridization is sp3d.

Molecular Orbital (MO) Theory

Molecular Orbital Theory explains chemical bonding as the combination of atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

• Electron Orbital Diagrams: Electrons fill atomic orbitals according to the Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule.

• Bonding and Anti-bonding Orbitals: When atomic orbitals combine, they form bonding (lower energy, increased stability) and anti-bonding (higher energy, decreased stability) molecular orbitals.

• Bonding Molecular Orbital (σ, π): Region of increased electron density between nuclei, promoting bond formation.

• Anti-bonding Molecular Orbital (σ*, π*): Region with a node (zero electron density) between nuclei, preventing bond formation.

• Filling Order: Electrons fill molecular orbitals from lowest to highest energy, following the same principles as atomic orbitals.

• Example: For N2 (Z = 7), the electron configuration in molecular orbitals is determined by combining the atomic orbitals of two nitrogen atoms.

MO Theory: Homonuclear Diatomic Molecules

Homonuclear diatomic molecules consist of two identical atoms. The order of molecular orbitals (σ2p and π2p) can change depending on the element due to differences in electronegativity and atomic size.

• For B2, C2, N2: The π2p orbitals are lower in energy than the σ2p orbital.

• For O2, F2, Ne2: The σ2p orbital is lower in energy than the π2p orbitals.

• Example: The electron configuration for F2– is written by filling the molecular orbitals with the appropriate number of electrons.

MO Theory: Heteronuclear Diatomic Molecules

Heteronuclear diatomic molecules are composed of two different elements. The more electronegative atom’s atomic orbitals are lower in energy and dominate the molecular orbital diagram.

• Electronegativity: Increases up and to the right on the periodic table; more electronegative atoms have lower energy orbitals.

• MO Diagram Selection: The diagram is based on the more electronegative element.

• Example: For CO, the molecular orbital diagram is constructed using the atomic orbitals of carbon and oxygen, with oxygen’s orbitals lower in energy.

MO Theory: Bond Order

Bond order is a measure of the number of chemical bonds between a pair of atoms. It is calculated using the molecular orbital diagram.

• Bond Order Formula:

• Interpretation: A bond order greater than zero indicates a stable molecule; a bond order of zero means the molecule is unstable and does not exist.

• Bond Order and Bond Type:

  • Single bond: Bond order = 1

  • Double bond: Bond order = 2

  • Triple bond: Bond order = 3

• Example: For NO–, calculate the bond order using the MO diagram and the formula above.

Practice Problems and Applications

• Identify the number of electron groups, lone pairs, and bonding groups in various molecules using Lewis structures and VSEPR theory.

• Determine the hybridization and unhybridized orbitals for central atoms in molecules such as KrBr4 and BeCl2.

• Construct and fill molecular orbital diagrams for homonuclear and heteronuclear diatomic molecules (e.g., O22+, P2, CF+, BN).

• Calculate bond order and predict molecular stability and bond strength for ions and molecules (e.g., SH–, O2, O2–, O2+, CN, CN–).