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Unit I: Gases

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    The study of gases allows us to understand the behavior of matter at its simplest: individual particles, acting independently, almost completely uncomplicated by interactions and interferences between each other. This knowledge of gases will serve as the pathway to our understanding of the far more complicated condensed phases (liquids and solids) in which the theory of gases will no longer give us correct answers, but it will still provide us with a useful model that will at least help us to rationalize the behavior of these more complicated states of matter.

    • 2.1: Some Definitions
    • 2.2: An Operational Definition of Temperature
    • 2.3: Ideal Gases
      In an ideal gas, there are no interactions between the particles, hence, the particles do not exert forces on each other. However, particles do experience a force when they collide with the walls of the container.
    • 2.4: Real Gases
      So what happens when a real gas is subjected to a very high pressure? The outcome varies with both the molar mass of the gas and its temperature, but in general we can see the the effects of both repulsive and attractive intermolecular forces.
    • 2.5: Condensation of Gases and the Critical State
      The most striking feature of real gases is that they cease to remain gases as the temperature is lowered and the pressure is increased.
    • 2.6: Kinetic Theory of Gases
      The kinetic theory describes a gas as a large number of submicroscopic particles, all of which are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container. Kinetic theory explains macroscopic properties of gases, such as pressure, temperature, viscosity, thermal conductivity, and volume, by considering their molecular composition and motion.
    • 2.7: The Maxwell Distribution Laws
      The Maxwell-Boltzmann distribution is used to determine how many molecules are moving between velocities.
    • 2.8: Molecular Collisions and the Mean Free Path
      The collision theory states that when suitable particles of the reactant hit each other, only a certain percentage of the collisions cause any noticeable or significant chemical change; these successful changes are called successful collisions. The successful collisions have enough energy, also known as activation energy, at the moment of impact to break the preexisting bonds and form all new bonds. This results in the products of the reaction.
    • 2.9: Graham's Laws of Diffusion and Effusion
      Graham's Law states that the effusion rate of a gas is inversely proportional to the square root of the mass of its particles.
    • 2.E: Properties of Gases (Exercises)
      This are exercises that to accompany the TextMap organized around Raymond Chang's Physical Chemistry for the Biosciences textbook.

    Unit I: Gases is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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