Intermolecular Forces: Attraction between Particles (Chapter 6 Study Notes)

Intermolecular Forces: Attraction between Particles

Lecture Outline

• 6.1 Intramolecular forces versus intermolecular forces

• 6.2 London Dispersion Forces

• 6.3 Interactions Involving Polar Molecules

• 6.4 Trends in Solubility

• 6.5 Phase Diagrams: Intermolecular Forces at Work

• 6.6 Some Remarkable Properties of Water

6.1 Intramolecular Forces versus Intermolecular Forces

Definitions and Differences

Understanding the distinction between intramolecular and intermolecular forces is essential for explaining the physical properties of substances. These forces determine the state of matter and many observable phenomena.

• Intramolecular forces: The forces that hold atoms together within a molecule (e.g., covalent, ionic, metallic bonds).

• Intermolecular forces: The forces of attraction or repulsion between molecules (e.g., hydrogen bonding, dipole-dipole, London dispersion).

• Physical properties such as boiling point, melting point, and solubility depend on the relative strength of these forces compared to the kinetic energy of the particles.

States of Matter:

• As kinetic energy decreases (from gas to liquid to solid), attractive intermolecular forces become more significant.

Example: Water in space forms a blob due to intermolecular forces acting between water molecules.

6.2 London Dispersion Forces

Nature and Trends

London dispersion forces are the weakest type of intermolecular force but are present in all molecules, especially nonpolar ones. They arise from temporary fluctuations in electron density, creating instantaneous dipoles.

• Definition: London dispersion forces are attractions resulting from temporary dipoles induced in atoms or molecules due to electron movement.

• Strength increases with:

  • Molecular weight (more electrons = greater polarizability)

  • Surface area (greater contact between molecules)

• These forces explain why larger, heavier molecules have higher boiling points.

Example: Pentane (C5H12) and neopentane (C5H12) have the same molecular weight but different boiling points due to differences in surface area.

6.3 Interactions Involving Polar Molecules

Dipole-Dipole Forces

Dipole-dipole forces occur between polar molecules, where the positive end of one molecule is attracted to the negative end of another.

• Definition: Dipole-dipole forces are attractions between the positive and negative ends of polar molecules.

• These forces are stronger than London dispersion forces but weaker than hydrogen bonds.

• They influence boiling points, melting points, and solubility.

Example: Acetone (CH3COCH3) has a higher boiling point than butane (C4H10) due to dipole-dipole interactions.

Hydrogen Bonding

Hydrogen bonding is a special, strong type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms (O, N, or F).

• Definition: Hydrogen bonding is the attraction between a hydrogen atom bonded to O, N, or F and a lone pair on another O, N, or F atom.

• Hydrogen bonds are responsible for many unique properties of water and biological molecules.

Example: Water (H2O) molecules can form up to four hydrogen bonds, leading to high boiling and melting points.

Applications: Hydrogen bonding is crucial for protein and DNA structure.

Ion-Dipole Forces

Ion-dipole forces occur when ionic compounds are mixed with polar solvents, such as water. These are important in aqueous solutions.

• Definition: Ion-dipole forces are attractions between ions and the partial charges of polar molecules.

• These forces are stronger than dipole-dipole and London dispersion forces.

Example: NaCl dissolving in water involves ion-dipole interactions between Na+/Cl- ions and water molecules.

Identifying Intermolecular Forces

• Nonpolar molecules (e.g., CH4): Only London dispersion forces.

• Polar molecules (e.g., HCl): Dipole-dipole and London dispersion forces.

• Molecules with O-H, N-H, or F-H bonds (e.g., CH3COOH): Hydrogen bonding, dipole-dipole, and London dispersion forces.

6.4 Trends in Solubility

Solubility Principles

Solubility describes how much of a substance (solute) can dissolve in a solvent at a given temperature. The nature of intermolecular forces determines solubility.

• Solution: Homogeneous mixture of solute and solvent.

• Solubility (g/mL): Maximum amount of solute that dissolves in a given amount of solvent at a specified temperature.

• Miscible: Substances that mix in all proportions without separating.

• Immiscible: Substances that do not mix (e.g., oil and water).

• "Like dissolves like": Polar solvents dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes.

Example: Water dissolves NaCl (ionic, polar) but not grease (nonpolar).

6.5 Phase Diagrams: Intermolecular Forces at Work

Phase Changes and Diagrams

Phase diagrams show the states of matter (solid, liquid, gas) of a substance as a function of temperature and pressure. Intermolecular forces influence phase transitions.

• Vaporization: Molecules gain enough kinetic energy to overcome intermolecular forces and enter the gas phase.

• Pressure ($P$): Defined as force per unit area ($P = \frac{F}{A}$).

• Atmospheric pressure: Pressure exerted by the weight of air above Earth's surface.

• Triple point: The unique set of conditions where all three phases coexist.

• Critical point: The temperature and pressure above which a substance cannot exist as a liquid.

Example: Water's phase diagram features a positive slope for the solid-liquid transition, indicating that liquid water is denser than ice.

6.6 Some Remarkable Properties of Water

Unique Physical Properties

Water exhibits several remarkable properties due to hydrogen bonding, making it essential for life and environmental processes.

• High specific heat: Water absorbs large amounts of heat with little temperature change.

• High heat of vaporization: Large amounts of heat are required for evaporation, aiding temperature regulation.

• Density anomaly: Liquid water is denser than solid ice, causing ice to float.

• Surface tension: Water molecules at the surface experience stronger attraction, minimizing surface area.

• Capillary action: Water can flow against gravity in narrow tubes due to cohesive and adhesive forces.

Example: Water covers most of Earth's surface and constitutes about 66% of the human body.

Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion	All molecules/atoms	Weakest	CH4, noble gases
Dipole-Dipole	Polar molecules	Intermediate	HCl, acetone
Hydrogen Bonding	H bonded to O, N, or F	Strong	H2O, NH3
Ion-Dipole	Ions and polar molecules	Strongest	NaCl in water