Hybridization, Molecular Geometry, and Molecular Orbital Theory: Study Notes

Hybridization and Molecular Geometry

Electron Group Geometry and Hybrid Orbitals

Understanding the geometry and hybridization of molecules is essential for predicting molecular shapes and bonding properties. The electron group geometry is determined by the number of electron domains (bonding and lone pairs) around the central atom, while hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Electron Group Geometry: The arrangement of electron groups (bonds and lone pairs) around a central atom.

• Hybridization: The process of combining atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds.

Common Hybridizations:

• sp: Linear geometry (180° bond angle)

• sp2: Trigonal planar geometry (120° bond angle)

• sp3: Tetrahedral geometry (109.5° bond angle)

• sp3d: Trigonal bipyramidal geometry (90°, 120° bond angles)

• sp3d2: Octahedral geometry (90° bond angles)

Example: In methane (CH4), the central carbon atom has four electron groups and is sp3 hybridized, resulting in a tetrahedral geometry.

Determining Hybridization and Geometry in Molecules

Stepwise Approach

1. Draw the Lewis structure of the molecule.

2. Count the number of electron groups (bonds and lone pairs) around the central atom.

3. Assign the electron group geometry based on the number of groups.

4. Determine the hybridization that matches the geometry.

Example: For SF6 (sulfur hexafluoride):

• Six bonding pairs around S → octahedral geometry

• Hybridization: sp3d2

Observed Bonding and Hybridization in Organic Molecules

Assigning Hybridization to Specific Atoms

In organic molecules, each atom's hybridization can be determined by the number of regions of electron density (bonds and lone pairs) around it.

• Single bonds (σ bonds): Count as one region.

• Double/triple bonds: Each counts as one region (π bonds do not affect hybridization count).

• Lone pairs: Each counts as one region.

Example: In ethanol (CH3CH2OH), the carbon atoms are sp3 hybridized, while the oxygen is also sp3 hybridized due to two lone pairs and two bonds.

Molecular Orbital Theory and Bond Order

Molecular Orbital Diagrams

Molecular orbital (MO) theory describes the distribution of electrons in molecules using orbitals that extend over the entire molecule. Electrons fill molecular orbitals in order of increasing energy, following the Pauli exclusion principle and Hund's rule.

• Bonding orbitals: Lower in energy, stabilize the molecule.

• Antibonding orbitals: Higher in energy, destabilize the molecule.

Bond Order Formula:

$\text{Bond Order} = \frac{1}{2} (\text{Number of electrons in bonding MOs} - \text{Number of electrons in antibonding MOs})$

Example: For O2 (oxygen):

• Bonding electrons: 10

• Antibonding electrons: 6

• Bond order: (10 - 6)/2 = 2

Bond Energy and Paramagnetism

Bond Energy

Bond energy is related to bond order; higher bond order generally means stronger, shorter bonds and higher bond energy.

• As bond order increases, bond energy increases.

• As bond order decreases, bond energy decreases.

Example: N2 (triple bond, bond order 3) has a higher bond energy than O2 (double bond, bond order 2).

Paramagnetism

A molecule is paramagnetic if it has one or more unpaired electrons in its molecular orbitals. Paramagnetic substances are attracted to magnetic fields.

• O2: Paramagnetic (has two unpaired electrons in π* orbitals)

• N2: Diamagnetic (all electrons paired)

Summary Table: Hybridization and Geometry

Electron Groups	Geometry	Bond Angle	Hybridization
2	Linear	180°	sp
3	Trigonal planar	120°	sp2
4	Tetrahedral	109.5°	sp3
5	Trigonal bipyramidal	90°, 120°	sp3d
6	Octahedral	90°	sp3d2

Practice Problems (from Worksheet)

1. Name the electron group geometry and the hybrid orbital used by the inner atoms in various molecules (e.g., SF6, CH2Cl2, PH3, etc.).

2. Indicate the hybridization of specified atoms and observed bonding around those atoms in organic molecules.

3. Complete molecular orbital diagrams for diatomic molecules and calculate bond order.

4. Compare bond energies using molecular orbital diagrams.

5. Identify paramagnetic diatomic molecules.

Additional info: These problems reinforce concepts from Chapter 10 (Molecular Shapes & Valence Bond Theory) and Chapter 9 (Chemical Bonding I: The Lewis Model), as well as introduce basic molecular orbital theory (Chapter 10/11).