# 1: The Properties of Gases

• 1.1: The Empirical Gas Laws
A number of important relationships describing the nature of gas samples have been derived completely empirically (meaning based solely on observation rather making an attempt to define the theoretical reason these relationships may exist. These are the empirical gas laws.
• 1.2: The Ideal Gas Law
The ideal gas law combines the empirical laws into a single expression. It also predicts the existence of a single, universal gas constant, which turns out to be one of the most important fundamental constants in science.  As derived here, it is based entirely on empirical data. It represents “limiting ideal behavior.” As such, deviations from the behavior suggested by the ideal gas law can be understood in terms of what conditions are required for ideal behavior to be followed (or approached).
• 1.3: The Kinetic Molecular Theory of Gases
The gas laws were derived from empirical observations. Connecting them to fundamental properties of the gas particles is subject of great interest. The Kinetic Molecular Theory is one such approach. In its modern form, the Kinetic Molecular Theory of gasses is based on five basic postulates.
• 1.4: Kinetic Energy
It is also important to recognize that the most probable, average, and RMS kinetic energy terms that can be derived from the Kinetic Molecular Theory do not depend on the mass of the molecules. As such, it can be concluded that the average kinetic energy of the molecules in a thermalized sample of gas depends only on the temperature. However, the average speed depends on the molecular mass. So, for a given temperature, light molecules will travel faster on average than heavier molecules.
• 1.5: Graham’s Law of Effusion
An important consequence of the kinetic molecular theory is what it predicts in terms of effusion and diffusion effects. Effusion is defined as a loss of material across a boundary
• 1.6: Collisions with Other Molecules
A major concern in the design of many experiments is collisions of gas molecules with other molecules in the gas phase. For example, molecular beam experiments are often dependent on a lack of molecular collisions in the beam that could degrade the nature of the molecules in the beam through chemical reactions or simply being knocked out of the beam.
• 1.7: Real Gases
While the ideal gas law is sufficient for the prediction of large numbers of properties and behaviors for gases, there are a number of times that deviations from ideality are extremely important.
• 1.8: Intermolecular Forces
Intermolecular forces are the attractive or repulsive forces between molecules. They are separated into two groups; short range and long range forces. Short range forces happen when the centers of the molecules are separated by three angstroms (10-8 cm) or less. Short range forces tend to be repulsive, where the long range forces that act outside the three angstroms range are attractive. Long range forces are also known as Van der Waals forces. They are responsible for surface tension, friction,
• 1.9: Specific Interactions
Intermolecular forces are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the intramolecular forces, which keep a molecule together (e.g., covalent and ionic bonding).
• 1.E: Gases (Exercises)
Exercises for Chapter 2 "Gases" in Fleming's Physical Chemistry Textmap.
• 1.S: Gases (Summary)
Summary for Chapter 2 "Gases" in Fleming's Physical Chemistry Textmap.

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