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GENERAL CHEMISTRY II

Chapter 1: The Kinetic Molecular Model and Intermolecular Forces of Attraction in Matter

Section 1.1: Kinetic Molecular Theory of Solids and Liquids

  • Essential Question (EQ): Why do solids and liquids behave differently?

Kinetic Molecular Theory

  • Explains properties of solids and liquids via intermolecular forces of attraction and the kinetic energy of particles.

Key Points

  1. Matter Composition: All matter is made of tiny particles.

  2. Constant Motion: Particles are in constant motion.

  3. Temperature Variation: Speed of particles is proportional to temperature; higher temperature leads to greater speed.

  4. States of Matter Differences: Solids, liquids, and gases differ in:

    • Distances between particles

    • Freedom of motion

    • Interaction extent among particles.

States of Matter

  • Solid: Defined shape and volume

  • Liquid: Definite volume but takes shape of the container

  • Gas: Assumes shape and volume of the container.

Activity 1: Comparison of States

  • Rank phases in increasing distance between particles: Solid < Liquid < Gas

  • Characterize movement of particles:

    • Solid: Vibration in place

    • Liquid: Medium speed, limited distances

    • Gas: Random, fast, covers large distances

  • Arrangement of molecules:

    • Gas: far apart

    • Liquid: close together, free to move

    • Solid: close together, fixed

  • Increasing volume of empty space: Solid < Liquid < Gas

Properties of Matter

Property

Gas

Liquid

Solid

Volume/Shape

Assumes volume and shape of container

Fixed volume; takes shape of container

Fixed volume; fixed shape

Density

Low

High

High

Compressibility

Easy to compress

Cannot be appreciably compressed

Cannot be appreciably compressed

Molecular Motion

Random, fast, cover large distances

Random, medium speed, limited distances

Vibration in place

Intermolecular Forces of Attraction

Section 1.2: Definition

  • Intermolecular Forces: Attractive forces between molecules or particles in solid or liquid states.

  • Weaker than intramolecular forces (within molecules).

Types of Intermolecular Forces

  1. Dipole-dipole

  2. Hydrogen bonding

  3. Ion-dipole

  4. London dispersion

  5. Dipole-induced dipole

Dipole-Dipole Forces

  • Exist between polar molecules.

  • Attraction between oppositely charged ends of dipoles.

Hydrogen Bonding

  • Strong dipole-dipole force between hydrogen and highly electronegative atoms (N, O, F).

  • Found in polar molecules containing hydrogen bonded to these elements.

Ion-Dipole Forces

  • Act between an ion and a polar molecule.

  • Crucial for explaining ionic compounds' solubility in polar solvents (e.g., water).

London Dispersion Forces

  • Weakest intermolecular force.

  • Occurs when non-polar molecules form instantaneous dipoles.

Capillary Action

  • Tendency of liquid to rise in narrow tubes or openings due to intermolecular attractions.

  • Combines cohesion (like molecules) and adhesion (unlike molecules).

Viscosity

  • Resistance of a liquid to flow; describes thickness or thinness of liquids.

  • Stronger intermolecular forces correlate with higher viscosity.

  • Examples include honey and oil.

Vapor Pressure

  • Pressure exerted by vapor in equilibrium with its liquid or solid state.

  • Strong intermolecular forces result in low vapor pressures (e.g., water vs. ethyl ether).

Boiling Point

  • Temperature at which a liquid's vapor pressure equals atmospheric pressure.

  • Higher intermolecular forces correspond to higher boiling points.

Unique Properties of Water

  1. Good solvent

  2. High specific heat - the heat required to change temperature.

  3. Anomalously high boiling point.

  4. Solid water (ice) is less dense than liquid water, making ice float.

Types and Properties of Solids

Section 1.4: Classification

  • Crystalline solids: Regular arrangement of particles.

  • Amorphous solids: Disordered structure (e.g., glass).

Phase Changes

Section 1.5: Transformations of Matter

  • Phase changes occur with added/removed energy, changing molecular order from solid (high order) to gas (high disorder).

Phase Change Types

  • Melting

  • Freezing

  • Vaporization

  • Condensation

  • Sublimation

Phase Diagrams

  • Graphical representation of physical states under temperature and pressure variations.

  • Indicates equilibrium states and transitions between solid, liquid, and gas phases.

Notable Points on Phase Diagrams

  1. Triple Point: Conditions for all three phases to coexist.

  2. Critical Point: Where liquid and gas merge into a superfluid state.

Activity: Constructing a Phase Diagram

Tasks:

  1. Sketch phase diagram with specified points.

  2. Describe states at pressures and temperatures beyond stated points.

  3. Explain phase changes across specified temperatures and pressures.