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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.