(
课件网) Chapter 11 Gases Gases 11 11.1 Properties of Gases 11.2 The Kinetic Molecular Theory of Gases Molecular Speed Diffusion and Effusion 11.3 Gas Pressure Definition and Units of Pressure Calculation of Pressure Measurement of Pressure 11.4 The Gas Laws Boyle’s Law: The Pressure-Volume Relationship Charles’s and Gay-Lussac’s Law: The Temperature-Volume Relationship Avogadro’s Law: The Amount-Volume Relationship The Gas Laws and Kinetic Molecular Theory The Combined Gas Law: The Pressure-Temperature-Amount-Volume Relationship 11.5 The Ideal Gas Equation Applications of the Ideal Gas Equation Gases 11 11.6 Real Gases Factors That Cause Deviation from Ideal Behavior The van der Waals Equation van der Waals Constants 11.7 Gas Mixtures Dalton’s Law of Partial Pressures Mole Fractions 11.6 Reactions with Gaseous Reactants and Products Calculating the Required Volume of a Gaseous Reactant Determining the Amount of Reactant Consumed Using Change in Pressure Using Partial Pressures to Solve Problems Properties of Gases Gases differ from solids and liquids in the following ways: A sample of gas assumes both the shape and volume of the container. Gases are compressible. The densities of gases are much smaller than those of liquids and solids and are highly variable depending on temperature and pressure. Gases form homogeneous mixtures (solutions) with one another in any proportion. 11.1 The Kinetic Molecular Theory The kinetic molecular theory explains how the molecular nature of gases gives rise to their macroscopic properties. The basic assumptions of the kinetic molecular theory are as follows: A gas is composed of particles that are separated by large distances. The volume occupied by individual molecules is negligible. Gas molecules are constantly in random motion, moving in straight paths, colliding with perfectly elastic collisions. Gas molecules do not exert attractive or repulsive forces on one another. The average kinetic energy of a gas molecules in a sample is proportional to the absolute temperature: 11.2 The Kinetic Molecular Theory Gases are compressible because molecules in the gas phase are separated by large distances (assumption 1). Pressure is the result of the collisions of gas molecules with the walls of their container (assumption 2). Decreasing volume increases the frequency of collisions. Pressure increases as collision frequency increases. The Kinetic Molecular Theory Heating a sample of gas increases its average kinetic energy (assumption 4). Gas molecules must move faster. Faster molecules collide more frequently and at a greater speed. Pressure increases as collision frequency increases. The Kinetic Molecular Theory The total kinetic energy of a mole of gas is equal to: The average kinetic energy of one molecule is: For one mole of gas: m is the mass is the mean square speed rearrange and Take the square root (m x NA = M) The Kinetic Molecular Theory The root-mean ... ...
