U9 Notes-1 Ch 14 States of Matter notes HCHem

States of Matter

• Matter exists in three primary states: solids, liquids, and gases.

Liquids & Solids Overview

• Liquids: Definite volume, take shape of the container, incompressible.

• Solids: Definite shape and volume, particles in fixed positions.

• Gases: Total disorder, particles are far apart and compressible.

Key Comparisons of States

Properties of Matter

• Gas

  • Assumes volume and shape of container

  • Highly compressible

  • Flows readily

  • Rapid diffusion

• Liquid

  • Assumes shape of the container it occupies

  • Virtually incompressible

  • Flows readily

  • Slow diffusion

• Solid

  • Retains own shape and volume

  • Virtually incompressible

  • Does not flow

  • Extremely slow diffusion

Kinetic Theory of Gases

• All matter consists of tiny particles in constant motion.

• Key assumptions of gas particles:

  • Small hard spheres with insignificant volume.

  • No attractive or repulsive forces.

  • Rapid, constant, straight-line motion until colliding.

  • Elastic collisions with no loss in kinetic energy (KE).

Gas Properties

• Compressibility: Gases are compressible with low densities (1 g/L).

• Uniform pressure: Gases exert uniform pressure on surfaces.

• Vacuum: Empty space without particles has no pressure.

• Atmospheric Pressure: Decreases with elevation due to lesser density.

Measuring Pressure

• Barometer: Measures atmospheric pressure. At sea level, supports a 760 mm Hg column.

Kinetic Energy & Temperature

• Kinetic energy increases with temperature,

• Absolute Zero: Theoretical temperature (0K or -273°C) where particles cease movement.

• At a given temperature, all substances have the same average kinetic energy.

Liquids

• Exist in a narrow range of temperature and pressure.

• Intermolecular forces (IMF) hold particles together.

• Liquids are denser than gases, incompressible, and have a definite volume.

Key Terms

• Intermolecular Forces (IMF): Attractive forces between particles, strongest in solids and weakest in gases.

• Equilibrium: Dynamic state where opposing changes occur simultaneously.

Intermolecular Forces

• Responsible for the attraction between liquid particles and vary in strength among states of matter.

Vaporization

• Evaporation: Transition from liquid to gas occurring at the surface of non-boiling liquids and is a cooling process.

• Temperature affects evaporation rates; higher temperatures increase KE and evaporation.

Equilibrium Vapor Pressure

• Occurs when vapor pressure equals the rate of condensation above a liquid in a closed system (e.g., sealed container).

Gas-Liquid Equilibrium

• Molecules continuously vaporizing and condensing until dynamic equilibrium.

Vapor Pressure Characteristics

• Increases with temperature.

• High vapor pressure (volatile) liquids have weak IMF and lower boiling points.

• Measured with a manometer.

Boiling and Boiling Point

• Boiling: conversion of liquid to vapor at a given temperature (boiling point).

• Normal Boiling Point: Occurs at 1 atm; for water, it is 100°C.

• During boiling, the temperature remains constant as heat is used to overcome attractions.

Effects of Atmospheric Pressure on Boiling Point

• At higher elevations (lower atmospheric pressure), boiling occurs at lower temperatures.

• Cooking in a pressure cooker raises boiling point and cooks food faster.

Viscosity and Surface Tension

• Viscosity: Resistance to flow, increases with stronger IMF and decreases with temperature.

• Surface Tension: Inward pull by surface particles minimizes liquid surface area; contributes to the spherical shape of droplets.

Properties of Solids

• Solid particles are in fixed positions and densely packed, vibrating but not flowing.

• Melting Point: Temperature at which solids change to liquids.

Crystalline Solids and Allotropes

• Most solids are crystalline with a defined arrangement of particles (crystal lattice).

• Allotropes: Different forms of the same element (e.g., carbon as diamond or graphite).

Amorphous Solids

• Lack a well-defined arrangement and long-range molecular order (e.g., glass, rubber).

Phase Changes

• Key Transitions:

  • Melting: Solid to liquid

  • Freezing: Liquid to solid

  • Vaporization: Liquid to gas

  • Condensation: Gas to liquid

  • Sublimation: Solid to gas

  • Deposition: Gas to solid

• Changes of state involve energy changes (absorbed or released).

Phase Diagrams

• Graphs showing the relationship between pressure and temperature and the phases of substances.

• Triple Point: Conditions where all three phases exist in equilibrium.

Critical Points and Measurements

• Critical Temperature: Above which a substance cannot exist as a liquid, irrespective of pressure.

• Critical Pressure: Lowest pressure for a substance to exist as a liquid at its critical temperature.