Intermolecular Forces flashcards

Intermolecular Forces, Valence Bond Theory & Molecular Orbital Theory

Overview of Bonding Theories

• These theories explain carbon's ability to bond with four different atoms despite having fewer than four unpaired electrons.

Valence Bond Theory (VB Theory)

• Hybridization: Orbitals hybridize or blend to create new orbitals so that four electrons become unpaired.

  • Types of Hybridization:

    • sp hybridization: 2 electron domains around the atom.

    • sp2 hybridization: 3 electron domains around the atom.

    • sp3 hybridization: 4 electron domains around the atom.

• Lone Pairs: Lone pairs count as a "connection" in hybridization.

• Bond Types:

  • Single bonds are sigma bonds (σ).

  • Double bonds consist of 1 sigma and 1 pi (π) bond.

  • Triple bonds consist of 1 sigma and 2 pi bonds.

Molecular Orbital (MO) Theory

• States that electrons delocalize and spread across the entire molecule rather than just hybridizing to form new orbitals.

• Significance: Explains electrical conductivity in some covalent compounds and the stability of benzene.

Solutions and Dissolving Salts

Three Steps of Dissolution

1. Break the lattice crystal (endothermic - requires energy).

2. Break Intermolecular Forces (IMFs) in water (endothermic).

3. Form new IMFs between water and ions (exothermic).

Hydration Spheres

• Form when ions interact with water molecules.

• Overall heat of hydration (ΔHhyd) can be positive or negative:

  • Positive ΔHhyd: Ion-dipole IMFs weaker than ionic bonds.

  • Negative ΔHhyd: Ion-dipole IMFs stronger than ionic bonds.

Bond Energy Principles

• Breaking Bonds: Endothermic process (requires energy).

• Forming Bonds: Exothermic process (releases energy).

Main Types of Intermolecular Bonds

1. Ion-Dipole Forces: Strongest; occur between ionic and polar covalent substances.

2. Hydrogen Bonds: Form between molecules containing O-H, N-H, or F-H bonds; strong dipole-dipole forces.

3. Dipole-Dipole Forces: Occur between polar covalent substances.

4. London Forces (Van der Waals): Weakest; present in all substances, significant between nonpolar covalent or polar and nonpolar substances.

5. Strength Comparison: Can only compare strengths if molecules have similar masses.

Properties Influenced by Intermolecular Forces (IMFs)

1. Boiling Points: Higher with stronger IMFs.

2. Melting Points: Higher with stronger IMFs.

3. Viscosity: Higher with stronger IMFs.

4. Hydration Energy: Negative with stronger IMFs; indicates stability.

5. Vapor Pressure: Lower with stronger IMFs.

Isomers

• Definition: Same chemical formula but different structures leading to different properties.

• Double Bonds: Limit rotation, leading to isomer variations classified as:

  • Cis Isomers: Same side of the double bond.

  • Trans Isomers: Opposite sides of the double bond.

Effect of Mass on Intermolecular Forces

1. Heavier polar compounds have stronger IMFs than lighter polar compounds.

2. Heavier nonpolar compounds have stronger IMFs than lighter nonpolar compounds.

3. Polarizability: Heavier molecules are more polarizable; can affect IMFs.

4. Comparison of IMFs between nonpolar molecules:

• A longer chain nonpolar molecule may exhibit stronger IMFs than a shorter chain nonpolar molecule.

5. Exceptions: A long chain nonpolar substance could potentially have stronger IMFs compared to a polar substance of lower mass.

Techniques for Separation of Mixtures

1. Chromatography: Separates chemicals by polarity and molar mass; substance most similar to solvent travels furthest.

2. Distillation: Liquid with weakest IMFs boils off first and is collected.

3. Filtration: Liquids and dissolved solids pass through filter paper, while insoluble solids are collected and washed.

4. Evaporation: Liquids evaporate, leaving dissolved solids behind.

State Changes and Phase Changes

• Melting: Endothermic process (energy absorbed).

• Freezing: Exothermic reaction (energy released).

• Understanding these principles helps explain the strength of IMFs that hold substances in solid, liquid, or gas states.