General Chemistry II - Chapter 1

General Chemistry II - Chapter 1

• This chapter covers the kinetic molecular model and intermolecular forces of attraction in matter.
• The kinetic molecular theory explains the properties of solids and liquids in terms of intermolecular forces of attraction and the kinetic energy of the individual particles.
• All matter is made up of tiny particles.
• These particles are in constant motion.
• The speed of particles is proportional to temperature. Increased temperature means greater speed.
• Solids, liquids, and gases differ in distances between particles, in the freedom of motion of particles, and in the extent to which the particles interact.

Section 1.1: Kinetic Molecular Theory of Solids and Liquids

• Solids, liquids, and gases differ in their properties due to differences in intermolecular forces and particle motion.
• The kinetic molecular theory provides a framework to understand these properties.
• Gases have no fixed volume and shape; liquids have a fixed volume but not shape; solids have fixed volume and shape.

Kinetic Molecular Theory Details

• All matter is composed of tiny particles.
• These particles are in constant motion.
• The speed of particles is proportional to temperature; higher temperature leads to faster movement.
• Solids, liquids, and gases differ in their arrangement and motion.

States of Matter

• Solids have particles tightly packed in a fixed arrangement.
• Liquids have particles close together but can move past each other.
• Gases have particles far apart and are in constant, random motion

Activity 1

• Compare distances between gas, liquid, and solid molecules and rank phases by increasing distance.
• Describe the characteristic movement of gas, liquid, and solid particles.
• Describe the arrangement of molecules in gas, liquid, and solid phases.
• Rank the phases in order of increasing volume of empty space.

Molecular Behavior in Different States

• Volume/Shape: Gas – assumes volume and shape of its container; Liquid – fixed volume; assumes shape of occupied part of container; Solid – fixed volume and shape.
• Density: Gas – low; Liquid – high; Solid – high.
• Compressibility: Gas – easy to compress; Liquid – cannot be appreciably compressed; Solid – cannot be appreciably compressed.
• Motion of Molecules: Gas – random, fast, covering large distances; Liquid – random, medium speed, limited distances; Solid – vibration in place.

Intermolecular Forces of Attraction

• Intermolecular forces (IMFs) are attractive forces between molecules or particles in the solid or liquid states.
• IMFs are weaker than intramolecular forces (forces within a molecule).

Types of Intermolecular Forces

• Dipole-dipole forces: Attractive forces between polar molecules.
• Hydrogen bonding: A special type of dipole-dipole force between hydrogen bonded to a highly electronegative atom (N, O, or F).
• Ion-dipole forces: Attractive forces between an ion and a polar molecule.
• London dispersion forces: Weakest type of IMF; occurs between nonpolar molecules due to temporary dipole moments.
• Dipole-induced dipole forces: Occurs when a polar molecule induces a temporary dipole in a nonpolar molecule.

Section 1.3: Intermolecular Forces and Properties of Liquids

• Intermolecular forces significantly influence the properties of liquids, including surface tension, viscosity, heat of vaporization, boiling point, and vapor pressure.

Liquid Properties

• Surface tension: The elastic force in the liquid surface; caused by IMFs.
• Viscosity: Liquid resistance to flow. Stronger IMFs lead to higher viscosity.
• Heat of vaporization: The energy required to convert 1 mole of liquid to gas.
• Boiling point: Temperature at which vapor pressure equals atmospheric pressure. Stronger IMFs lead to higher boiling points.
• Vapor pressure: Pressure exerted by liquid’s vapor when in equilibrium with the liquid. Lower vapor pressure means stronger IMFs.

Section 1.4: Types and Properties of Solids

• Solids are classified as crystalline or amorphous based on the arrangement of their particles.
• Crystalline solids have a highly regular arrangement of particles, leading to well-defined crystal lattices. A lattice is a three-dimensional system of points that indicates the positions of components (atoms, ions, or molecules) of the crystal.
• Amorphous solids have a disordered arrangement of particles.

Crystalline Solid Types

• Ionic: Hard, high melting point, good insulators but conduct when dissolved. (e.g., NaCl)
• Molecular: Soft, low melting point. (e.g., ice, dry ice)
• Metallic: Wide range of hardness and melting points. (e.g., silver, iron, brass).
• Network: Hard, high melting points. (e.g., diamond)
• Group 8A (Noble gases): Very low melting points. (e.g., argon)

Section 1.5: Phase Changes and Phase Diagrams

• Phase changes occur when energy is added or removed from a substance.
• Phase changes are transformations from one physical state to another (e.g., solid to liquid, liquid to gas).
• The molecular order changes during phase changes.

Types of Phase Changes

• Melting: Solid to liquid
• Freezing: Liquid to solid
• Vaporization: Liquid to gas
• Condensation: Gas to liquid
• Sublimation: Solid to gas
• Deposition: Gas to solid

Phase Diagrams

• Phase diagrams visualize the conditions of pressure and temperature under which different phases occur.
• Key points include: triple point, critical point, solid, liquid, and gas areas, and lines that separate regions representing equilibrium between phases (melting, vaporization, and sublimation).