Chapter 10

The Kinetic-Molecular Theory of Matter

Objectives

State the Kinetic-molecular theory of matter, and describe how it explains certain properties of matter.

List the five assumptions of the kinetic-molecular theory of gases. Define the terms ideal gas and real gas.

Describe each of the following characteristic properties of gases: expansion, density, fluidity, compressibility, diffusion, and effusion.

Describe the conditions under which a real gas deviates from "ideal" behavior.

  1. Section I The Kinetic-Molecular Theory of Matter
    1. Definition- particles of matter are always in motion
    2. The Kinetic-Molecular Theory of gases
      1. Ideal gas- an imaginary gas that perfectly fits all the assumption of the kinetic-molecular theory.
      2. Kinetic molecular theory is based on five assumptions
        1. Gases consist of large numbers of tiny particles that are far apart relative to their size.
        2. Collision between gas particles and between particles and container walls are elastic collision (no loss of energy)
        3. Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy of motion.
        4. There are no forces of attraction or repulsion between gas particles.
        5. The average kinetic energy of gas particles depends on the temperature of the gas KE=1/2 mv2
    3. The Kinetic-molecular Theory and the Nature of Gases
      1. Expansion- no definite shape or volume
      2. Fluidity- no attractive force so particles glide past each other. This makes it a fluid.
      3. Low density- about 1/1000 that of the liquid or solid.
      4. Diffusion and Effusion
        1. Diffusion- spontaneous mixing of the particles of two substances caused by their random motion.
        2. Effusion- is the process by which gas particles pass through a tiny opening.
    4. Deviations of Real Gases from Ideal Behavior.
      1. Real gas is a gas that does not behave completely according to the assumption of the kinetic-molecular theory
      2. Ideal gases- He, Ne, N2, H2 and are non-polar.

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    Objectives

    Define pressure and relate it to force.

    Describe how pressure is measured.

    Convert units of pressure.

    State the standard conditions of temperature and pressure.

  2. Section 10-2 Pressure
    1. Pressure and Force
      1. Pressure is defined as the force per unit area on a surface
        1. Pressure= Force / area
        2. Force is defined as a push or a pull
          1. SI unit Newton
          2. The force increases the speed of a particles.
          3. Newton’s second law
      2. Measuring Pressure
        1. Barometer- is a device used to measure atmospheric pressure
        2. Discovered by Evangelista Torricelli
        3. U shaped tube

         

         

         

      3. Units of Pressure
        1. Millimeters of mercury
          1. Symbol (mm Hg)
          2. 1 mm Hg is 1 torr
        2. Atmosphere of pressure
          1. Symbol (atm)
          2. 1 atm = 760 mm Hg
        3. SI unit is Pascal
          1. Symbol (Pa)
          2. Defined as 1 newton acting on 1 meter squared.
          3. 1 atm = 101.325 KPa.
      4. Standard temperature and Pressure
        1. Standard temperature= 0 degrees Celsius
        2. Standard temperature= 1atm
        3. Conversions
        4. Practice

     

     

     

     

     

    Objectives

    Using the kinetic-molecular theory to explain the relationships between gas volume, temperature, and pressure.

    Use Boyle’s law to calculate volume-pressure changes at constant temperature.

    Use Charle’s law to calculate volume-temperature changes at constant pressure.

    Use Gay-lussac’s law to calculate pressure-temperature changes at constant volume.

    Use the combined gas law to calculate volume-temperature-pressure changes.

    Use Dalton’s law of partial pressure to calculate partial pressures and total pressure.

  3. Section 10-3 The Gas Laws.
    1. The gas laws- are simple mathematical relationships between the volume, temperature, pressure, and amount of gas.
      1. Boyles’s law
        1. Discovered by Robert Boyle
        2. States that the volume of a fixed mass of gas varies inversely with the pressure at constant temperature
        3. Equation p1v1=p2v2
        4. Pressure and volume is inversely proportional.
        5. Practice
      2. Charles Law
        1. Discovered by Jacques Charles in 1787
        2. States that the volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature.
        3. Equation v1/t1=v2/t2
        4. Volume and temperature are directly proportional
        5. Practice
      3. Gay-lussac’s law
        1. Discovered by Gay-Lussac in 1802
        2. States that the pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature.
        3. Equation p1/t1=p2/t2
        4. Pressure and temperature are directly proportional.
        5. Practice

         

         

         

      4. The combined Gas Law
        1. States that the relationship between pressure, volume, and temperature or a fixed amount of gas.
        2. Equation p1v1/t1=p2v2/t2
        3. Practice
    2. Dalton’s Law of Partial Pressures
      1. States that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gas.
      2. Equation Pt= P1+P2+P3
      3. Gases collected by water displacement
        1. A gas collected by water displacement is not pure gas it is mixed with water vapor
        2. Equation Patm = Pgas + PH2O
        3. Practice