Molecular Polarity, Valence Bond Theory, and Hybridization: Study Notes

Resonance and Molecular Structure

Resonance in Molecules

Resonance occurs when it is possible to draw more than one valid Lewis structure for a molecule. The actual electronic structure is a weighted average of these possibilities, known as a resonance hybrid.

• Resonance hybrid: The true structure may resemble one resonance form more than others.

• Criteria for resonance structure importance:

  • Each atom should satisfy the octet rule.

  • Like charges on adjacent atoms are unfavorable.

  • Smaller formal charges (positive or negative) are preferable.

  • A more negative formal charge should reside on a more electronegative atom.

• Exceptions to the octet rule are not uncommon.

Example: The nitrate ion, $\text{NO}_3^-$, has multiple resonance structures.

Molecular Shape and VSEPR Theory

Valence-Shell Electron-Pair Repulsion (VSEPR) Model

The VSEPR model predicts the shape of molecules by considering the repulsion between electron charge clouds around a central atom.

• Electron clouds arrange themselves to minimize repulsion.

• Molecular geometry is determined by the number of charge clouds (bonding and lone pairs).

Example: Water ($\text{H}_2\text{O}$) has a bent shape due to two bonding pairs and two lone pairs on oxygen.

Molecular Polarity

Importance of Molecular Polarity

Molecular polarity significantly affects both chemical and physical properties of compounds.

• Chemical reactivity is often sensitive to molecular polarity.

• Physical properties such as solubility, melting and boiling points, and lattice arrangement in solids are strongly affected by polarity.

Molecular Polarity of Diatomic Molecules

The polarity of a diatomic molecule is determined by the difference in electronegativity between the two atoms.

• Polar covalent bond: Unequal sharing of electrons, resulting in a dipole moment ($\mu \neq 0$).

• Nonpolar covalent bond: Equal sharing of electrons, resulting in no dipole moment ($\mu = 0$).

Example:

• Carbon monoxide ($\text{CO}$): polar molecule, $\mu \neq 0$.

• Chlorine ($\text{Cl}_2$): nonpolar molecule, $\mu = 0$.

Molecular Polarity of Larger Molecules

For larger molecules, both bond polarities and molecular shape determine overall polarity.

• The molecular dipole moment ($\mu$) is the vector sum of individual bond dipoles.

• If bond dipoles cancel due to molecular geometry, the molecule is nonpolar.

Example:

• Carbon dioxide ($\text{CO}_2$): linear, bond dipoles cancel, $\mu = 0$ (nonpolar).

• Water ($\text{H}_2\text{O}$): bent, bond dipoles do not cancel, $\mu \neq 0$ (polar).

Molecular Polarity and Physical Properties

Isomers with different molecular polarities can have different physical properties, such as boiling points.

Isomer	Structure	Boiling Point (°C)	Dipole Moment ($\mu$)
cis-Dichloroethylene	HClC=CHCl	60.3	Nonzero
trans-Dichloroethylene	ClHC=CHCl	47.5	Zero

Example: The cis isomer has a higher boiling point due to its nonzero dipole moment.

Theories of Covalent Bonding

Overview

The octet rule and Lewis structures model valence electron distribution, while VSEPR predicts molecular shape. Quantum mechanics introduces atomic orbitals. Two main theories combine these ideas:

• Valence Bond (VB) Theory

• Molecular Orbital (MO) Theory

Valence Bond Theory

Valence bond theory explains covalent bond formation as the overlap of two half-filled atomic orbitals, each containing one electron of opposite spin.

• Bond strength depends on the amount of orbital overlap.

• Bonds formed by orbitals other than s have directionality.

• Destructive interference occurs if overlapping orbitals are not in the same phase.

Examples:

• Hydrogen molecule ($\text{H}_2$): overlap of two 1s orbitals.

• Fluorine molecule ($\text{F}_2$): overlap of two 2p orbitals.

• Hydrogen chloride ($\text{HCl}$): overlap of H 1s and Cl 3p orbitals.

Valence Bond Theory Practice

To describe bonding in molecules like BrCl or H2O:

• Draw the Lewis structure.

• Write the condensed orbital-filling diagrams for each atom.

• Identify which half-filled orbitals overlap to form bonds.

• Ensure the octet rule is satisfied for each atom.

Valence Bond Theory Summary

• Covalent bonds form through overlap of atomic orbitals, each containing one electron of opposite spin.

• Bonded atoms maintain their own atomic orbitals; the electron pair in the overlapping region is shared.

• Greater orbital overlap leads to stronger, more directional bonds.

Hybridization

Concept of Hybridization

Linus Pauling proposed that atomic orbitals in molecules can be rearranged and shaped differently from those in isolated atoms. Hybridization is the process of averaging atomic orbitals to form new, equivalent hybrid orbitals.

• Hybridization explains molecular shapes and bond angles observed experimentally.

sp3 Hybridization

sp3 hybridization involves the mixing of one s and three p atomic orbitals to form four degenerate sp3 hybrid orbitals.

• These orbitals arrange tetrahedrally with bond angles of 109.5°.

Example: Methane ($\text{CH}_4$) has four equivalent sp3 hybrid orbitals overlapping with hydrogen 1s orbitals.

sp3 Hybridization in Other Molecules

• Ammonia ($\text{NH}_3$): Three sp3 hybrid orbitals of nitrogen overlap with hydrogen 1s orbitals, forming a trigonal pyramidal shape with bond angles of 109.5°.

• Water ($\text{H}_2\text{O}$): Two sp3 hybrid orbitals of oxygen overlap with hydrogen 1s orbitals, resulting in a bent shape.

Hybridization and Molecular Shape

Hybridization applies to any central atom with four charge clouds. Other types of hybridization (sp, sp2, etc.) apply to atoms with different numbers of charge clouds.

Summary Table: Key Concepts

Concept	Description	Example
Resonance	Weighted average of multiple valid Lewis structures	$\text{NO}_3^-$
VSEPR Model	Predicts molecular shape by minimizing electron repulsion	$\text{H}_2\text{O}$ (bent)
Molecular Polarity	Depends on bond polarity and molecular geometry	$\text{CO}_2$ (nonpolar), $\text{H}_2\text{O}$ (polar)
Valence Bond Theory	Bonds form by overlap of half-filled atomic orbitals	$\text{H}_2$, $\text{F}_2$, $\text{HCl}$
Hybridization	Atomic orbitals mix to form equivalent hybrid orbitals	sp3 in $\text{CH}_4$, $\text{NH}_3$, $\text{H}_2\text{O}$

Key Equations

• Dipole Moment: $\mu = q \times r$

• Formal Charge: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

Additional info: These notes expand on brief points and images to provide full academic context, including definitions, examples, and equations relevant to General Chemistry students.