Chapter 8 Student Notes

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Bonding Theories and Molecular Structure

Introduction to Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule. Understanding molecular shapes is essential for predicting reactivity, polarity, and physical properties.

Electron density determines molecular geometry.
Models such as Lewis Structures, VSEPR Theory, Valence Bond Theory, and Molecular Orbital Theory are used to predict and explain molecular shapes.
Hybridization helps explain observed bond angles and molecular shapes.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the geometry of molecules based on the repulsion between electron groups (bonding pairs and lone pairs) around a central atom.

Electron groups (bonding pairs and lone pairs) arrange themselves to minimize repulsion.
The number of electron groups determines the electron geometry.
Molecular geometry is determined by the positions of atoms, considering lone pairs.

Steps for Determining VSEPR Geometry:

Draw the Lewis structure of the molecule.
Count the number of electron groups around the central atom.
Assign the electron geometry based on the number of electron groups.
Determine the molecular geometry by considering the positions of atoms and lone pairs.

VSEPR Geometries Table

Electron Groups	Electron Geometry	Lone Pairs	Molecular Geometry	Bond Angles	Examples
2	Linear	0	Linear	180°	CO2
3	Trigonal planar	0	Trigonal planar	120°	BF3
3	Trigonal planar	1	Bent	~120°	SO2
4	Tetrahedral	0	Tetrahedral	109.5°	CH4
4	Tetrahedral	1	Trigonal pyramidal	~107°	NH3
4	Tetrahedral	2	Bent	~104.5°	H2O
5	Trigonal bipyramidal	0	Trigonal bipyramidal	90°, 120°	PCl5
5	Trigonal bipyramidal	1	Seesaw	~90°, ~120°	SF4
5	Trigonal bipyramidal	2	T-shaped	~90°	ClF3
5	Trigonal bipyramidal	3	Linear	180°	I3-
6	Octahedral	0	Octahedral	90°	SF6
6	Octahedral	1	Square pyramidal	~90°	BrF5
6	Octahedral	2	Square planar	90°	XeF4

Hybridization and Valence Bond Theory

Valence Bond Theory explains bonding by the overlap of atomic orbitals. Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals.

sp hybridization: Linear geometry, 180° bond angle.
sp2 hybridization: Trigonal planar geometry, 120° bond angle.
sp3 hybridization: Tetrahedral geometry, 109.5° bond angle.
sp3d and sp3d2 hybridizations: Trigonal bipyramidal and octahedral geometries, respectively.

Sigma (σ) bonds are formed by the head-on overlap of orbitals, while pi (π) bonds are formed by the side-to-side overlap of p orbitals.

Single bonds are always sigma bonds.
Double bonds consist of one sigma and one pi bond.
Triple bonds consist of one sigma and two pi bonds.

Bond order is the number of chemical bonds between a pair of atoms. Higher bond order generally means greater bond strength and shorter bond length.

Molecular Polarity

Molecular polarity depends on both the polarity of individual bonds and the overall geometry of the molecule.

A molecule is polar if it has a net dipole moment.
Symmetrical molecules (e.g., CO2, CCl4) are often nonpolar even if they contain polar bonds.
Asymmetrical molecules (e.g., H2O, NH3) are usually polar.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces (IMFs) are forces of attraction between molecules, affecting boiling points, melting points, and solubility.

London Dispersion Forces (LDF): Present in all molecules; strength increases with molecular size and polarizability.
Dipole-Dipole Interactions: Occur between polar molecules.
Hydrogen Bonding: A strong dipole-dipole interaction occurring when H is bonded to N, O, or F.
Ion-Dipole Forces: Occur between ions and polar molecules; important in solutions.

Polarizability is the ease with which an electron cloud can be distorted. Larger atoms/molecules are more polarizable.

Effects of Intermolecular Forces

Stronger IMFs lead to higher boiling and melting points.
Hydrogen bonding is responsible for high boiling points in H2O, NH3, and HF.
LDFs are the only IMFs present in nonpolar molecules.

Molecular Orbital Theory

Introduction to Molecular Orbitals (MO)

Molecular Orbital Theory describes electrons in molecules as occupying molecular orbitals formed from the combination of atomic orbitals.

Bonding orbitals are lower in energy and stabilize the molecule.
Antibonding orbitals are higher in energy and destabilize the molecule.
Bond order can be calculated using the MO diagram.
Paramagnetic molecules have unpaired electrons; diamagnetic molecules have all electrons paired.

MO Diagrams and Examples

MO diagrams show the energy levels of molecular orbitals for diatomic molecules (e.g., H2, O2, N2).
O2 is paramagnetic due to two unpaired electrons in its π* orbitals.
Bond order and magnetic properties can be predicted from MO diagrams.

Practice and Application

Sample Problems and Applications

Draw Lewis structures and predict molecular geometry using VSEPR theory.
Determine hybridization and identify sigma and pi bonds in molecules.
Classify molecules as polar or nonpolar based on geometry and bond polarity.
Identify types of intermolecular forces present in various compounds.
Use MO diagrams to determine bond order and magnetic character.

Summary Table: Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion	All molecules	Weakest	CH4
Dipole-Dipole	Polar molecules	Intermediate	HCl
Hydrogen Bonding	H bonded to N, O, or F	Strongest (among dipole-dipole)	H2O
Ion-Dipole	Ions and polar molecules	Very strong	Na+ in H2O

Additional info:

Electron geometry and molecular geometry are identical when there are no lone pairs on the central atom.
Practice drawing Lewis structures and MO diagrams for exam preparation.
Polarizability increases with molecular size, affecting the strength of London dispersion forces.