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

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Topic 4: Chemical Bonding II

Overview of Covalent Bond Representation

Chemists use several models to represent covalent bonds and predict 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). Each model offers unique insights and has specific strengths and limitations.

Lewis Structures: Show connectivity and electron pairs but do not predict 3D shape.
VSEPR Theory: Predicts 3D molecular geometry based on electron group repulsion.
Valence Bond Theory (VBT): Describes bonding via orbital overlap and hybridization.
Molecular Orbital Theory (MOT): Provides a quantum mechanical view of molecules, explaining electronic structure and properties.

VSEPR Theory (Valence Shell Electron-Pair Repulsion)

Principles of VSEPR Theory

VSEPR theory is used to predict the shapes 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, or 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.

Counting Electron Groups: Examples

H2O: 2 bonds + 2 lone pairs = 4 electron groups
PH3: 3 bonds + 1 lone pair = 4 electron groups
CCl2O: 3 bonds + 1 lone pair = 4 electron groups

Electron Group Geometry and Predicted Bond Angles

# 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 from ideal values.
Multiple bonds are more repulsive than single bonds.
Greater repulsion leads to distortion from usual 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 Their Representations

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

# 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 Species with 4, 5, and 6 Electron Groups

For 4 electron groups (tetrahedral), lone pairs are placed to minimize repulsion, affecting bond angles (e.g., NH3 and H2O).
For 5 electron groups (trigonal bipyramidal), lone pairs occupy equatorial positions first.
For 6 electron groups (octahedral), lone pairs occupy axial positions first.

Strategy for Drawing VSEPR Structures

Draw the Lewis structure.
Determine the electron group geometry.
Place lone pair electrons and determine VSEPR structure. Identify bond angles if required.

VSEPR Applied to Larger Molecules

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

Identifying the centers in the Lewis structure.
Determining the number of electron groups and lone pairs at each center.
Determining the molecular shape at each center.

Molecular Dipole and Overall Polarity

Bond Polarity

Atoms share electrons in covalent 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

A vector has both direction and magnitude.
Bond dipoles are added head-to-tail; the resultant vector represents 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	Polarity between two atoms in a bond	Overall polarity of the molecule
How is it determined?	Difference in electronegativity	Vector sum of all bond dipoles
How is it represented?	Arrow or $\delta^+$/$\delta^-$	Arrow showing net dipole direction

Isomers (Ch 20.5-20.7)

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. Subdivided into:
- Geometric Isomers: Different spatial arrangement (e.g., cis/trans).
- Optical Isomers: Non-superimposable mirror images (enantiomers).

Structural Isomers

Can be identified from Lewis structures.
Example: C3H8O (propanol vs. isopropanol).

Stereoisomers

Same formula and connectivity, different arrangement in space.
Includes geometric and optical isomers.

Geometric Isomers

Different spatial arrangement of atoms (e.g., cis/trans forms).
Example: SF2Cl2 (cis and trans forms).

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).

Valence Bond Theory (VBT)

Overview of VBT

Valence Bond Theory describes covalent bonding as the overlap of atomic orbitals, forming bonds between atoms. It bridges the gap between VSEPR and Molecular Orbital Theory.

Bonding occurs via overlap of half-filled atomic or hybrid orbitals.
Each atom donates one electron to the bond.
Types of bonds: sigma (σ) and pi (π) bonds.

Sigma and Pi Bonds

Sigma (σ) bonds: Cylindrical symmetry around the internuclear axis; formed by end-to-end overlap.
Pi (π) bonds: Electron density above and below the internuclear axis; formed by side-to-side overlap of p orbitals.
Single bonds are always sigma; double bonds are one sigma and one pi; triple bonds are one sigma and two pi.

Hybridization in VBT

Atomic orbitals mix to form hybrid orbitals (sp, sp2, sp3, etc.) to explain observed molecular shapes.
Example: CH4 uses sp3 hybridization for tetrahedral geometry.

Molecular Orbital Theory (MOT)

Overview of MOT

MOT provides a quantum mechanical description of molecules, explaining electronic structure, reactivity, color, and magnetism.

Electrons are described by molecular orbitals formed from atomic orbital combinations.
Bonding and antibonding orbitals are formed by constructive and destructive overlap.
Bond order is calculated as:

$\text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

Bond order indicates bond strength and stability; zero bond order means no bond.

Magnetism and Quantum Chemistry

Paramagnetic: Molecules with unpaired electrons; attracted to magnetic fields.
Diamagnetic: Molecules with all electrons paired; not attracted to magnetic fields.
Magnetism can be measured experimentally (e.g., Gouy balance).

Practice and Application

Draw VSEPR structures and identify bond angles for various molecules.
Determine molecular geometry and polarity using VSEPR and vector addition.
Identify and classify isomers (structural, geometric, optical).
Apply VBT and MOT to describe bonding and predict properties.

Additional info: Some tables and diagrams were inferred and summarized for clarity and completeness. For full details, refer to Table 10.1 in your textbook for molecular geometries and bond angles.