Chemical Bonding II: VSEPR, Valence Bond Theory, and Molecular Orbital Theory

Topic 4: Chemical Bonding II

Overview of Covalent Bond Representation

Chemical bonding in molecules can be represented using several models, each providing different insights into molecular structure and properties. The main theories include Lewis structures, Valence Shell Electron Pair Repulsion (VSEPR) theory, Valence Bond Theory (VBT), and Molecular Orbital Theory (MOT).

• Lewis Structures: Show connectivity and electron pairs.

• VSEPR Theory: Predicts 3D molecular shapes based on electron repulsion.

• Valence Bond Theory (VBT): Describes bonding via orbital overlap.

• Molecular Orbital Theory (MOT): Provides a quantum mechanical view of molecules.

VSEPR Theory (Valence Shell Electron-Pair Repulsion)

Principles of VSEPR Theory

VSEPR theory is used to predict the geometry of molecules based on the repulsion between electron groups around a central atom.

• Electron groups are regions around the central atom where electrons are concentrated (bonds or lone pairs).

• Electron groups are identified as single, double, or triple bonds, and lone pairs.

• The number of electron groups is determined from the Lewis structure.

• Electron groups repel each other and arrange themselves as far apart as possible.

Electron Group Geometry and Bond Angles

The geometry and bond angles depend on the number of electron groups around the central atom.

# of Electron Groups	Electron Group Geometry	Predicted Bond Angles
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

Drawing Molecules in 3-D

• Wedge bond: Comes out of the page.

• Hash bond: Goes into the page.

• Solid bond: Lies in the plane of the page.

• Maximize the number of bonds in the plane for clarity.

Effect of Lone Pairs and Multiple Bonds on Bond Angles

• Lone pairs are more repulsive than bonding pairs, causing bond angles to decrease.

• Multiple bonds are more repulsive than single bonds.

• Greater repulsion leads to distortion from ideal bond angles.

• AXE notation: Used to describe the number of atoms (A), bonded atoms (X), and lone pairs (E) around the central atom.

Molecular Geometry and Lone Pairs

Molecular geometry is determined by both the electron group geometry and the number of lone pairs.

# e- group	e- group geometry (no lone pairs)	1 lone pair	2 lone pairs	3 lone pairs	4 lone pairs
2	Linear	-	-	-	-
3	Trigonal planar	Bent	-	-	-
4	Tetrahedral	Trigonal pyramidal	Bent	-	-
5	Trigonal bipyramidal	Seesaw	T-shaped	Linear	-
6	Octahedral	Square pyramidal	Square planar	-	-

Placement of Lone Pairs in 5 and 6 Electron Group Geometries

• In trigonal bipyramidal geometry (5 EG), lone pairs are placed in equatorial positions.

• In octahedral geometry (6 EG), lone pairs are placed in axial positions first.

Strategy for Drawing VSEPR Structures

1. Draw the Lewis structure.

2. Determine the electron group geometry.

3. Place lone pair electrons and determine VSEPR structure. Identify bond angles if required.

VSEPR for Larger Molecules

For larger molecules, determine the geometry around each central atom by:

1. Identifying the centers in the Lewis structure.

2. Determining the number of electron groups and lone pairs.

3. Determining the molecular shape at each center.

Molecular Dipole and Overall Polarity

Bond Polarity

• Atoms share electrons in bonds, but sharing is often unequal due to differences in electronegativity.

• Unequal sharing creates a bond dipole.

• The greater the electronegativity difference ($\Delta EN$), the more polar the bond.

• Dipole moment is represented by $\delta^+$/$\delta^-$ or an arrow ($\rightarrow$).

Molecular Dipole Moment

• Sum the individual bond dipoles (vector addition) to determine the overall molecular dipole.

• Non-polar molecules have no net dipole moment.

Vector Addition in Dipole Moments

• Vectors have both direction and magnitude.

• Bond dipoles are added head-to-tail; the resultant vector gives the molecular dipole.

• Used to determine if a molecule is polar or non-polar.

Summary Table: Bond Dipole vs. Molecular Dipole

	Bond Dipole	Molecular Dipole
Definition	Dipole moment of a single bond due to electronegativity difference	Net dipole moment of the entire molecule
How is it determined?	By difference in electronegativity between two atoms	By vector addition of all bond dipoles
How is it represented?	Arrow or $\delta^+$/$\delta^-$	Arrow showing overall direction

Isomers

Types of Isomers

Isomers are molecules with the same chemical formula but different arrangements of atoms. They are classified as:

• Structural Isomers: Different connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

  • Geometric Isomers: Different spatial arrangement (e.g., cis/trans).

  • Optical Isomers: Non-superimposable mirror images (enantiomers).

Structural Isomers

• Identified by different connectivity in Lewis structures.

• Example: C3H8O (propanol vs. isopropanol).

Stereoisomers

• Same formula and connectivity, different arrangement in space.

• Subdivided into geometric and optical isomers.

Geometric Isomers

• Different spatial arrangement, such as cis/trans forms.

• Example: SF2Cl2 (cis/trans).

Optical Isomers

• Non-superimposable mirror images, called enantiomers.

• Chiral molecules have at least one carbon with four different groups attached.

• Enantiomers can have different biological activities (e.g., drug effectiveness, smell).

Advanced Bonding Theories

Valence Bond Theory (VBT)

VBT describes chemical bonding as the overlap of atomic orbitals, forming bonds between atoms.

• Sigma (σ) bonds: Cylindrical symmetry about the internuclear axis; formed by end-to-end overlap.

• Pi (π) bonds: Side-by-side overlap; electron density above and below the axis.

• Hybridization explains observed molecular shapes by mixing atomic orbitals (e.g., sp, sp2, sp3).

Example: In methane (CH4), carbon undergoes sp3 hybridization to form four equivalent bonds.

Molecular Orbital Theory (MOT)

MOT provides a quantum mechanical description of molecules, where electrons occupy molecular orbitals formed from atomic orbitals.

• Describes bonding, electronic structure, and properties such as magnetism and color.

• Bond order is calculated as:

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

• Paramagnetic molecules have unpaired electrons; diamagnetic molecules have all electrons paired.

Practice and Application

• Draw VSEPR structures and determine bond angles for various molecules (e.g., AlCl3, CH4).

• Identify molecular geometry and bond angles for each atom in larger molecules (e.g., ethanoic acid).

• Classify isomers and determine relationships between molecules.

• Apply VBT and MOT to describe bonding and predict properties.

Additional info: The notes include textbook-style tables and diagrams for VSEPR geometries and isomer classification, and provide strategies for drawing and analyzing molecular structures using multiple bonding theories.