The Kinetic Molecular Theory

1. The Kinetic Molecular Theory

2. Kinetic Molecular Theory • We have been using an equation entitled the “ideal” gas law. • What is an “ideal” gas??

3. Kinetic Molecular Theory • No gas is ideal! • But, many gases come close to ideal under certain conditions

4. Kinetic Molecular Theory • The Kinetic Molecular Theory of Gases attempts to define what an “ideal” gas is.

5. Kinetic Molecular Theory • 1. The particles are so small compared with the distances between them, that the volume of the individual particles can be assumed to be zero.

6. Kinetic Molecular Theory • 2. The particles are in constant motion. The collisions of the particles with the walls of the container are the cause of pressure exerted by gas.

7. Kinetic Molecular Theory • 3.The particles are assumed to exert no forces on each other; they are assumed to neither to attract or repel each other.

8. Kinetic Molecular Theory • 4. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temp of the gas.

9. Real Gases • Real gas particles do have finite volumes and do exert forces. • Ideal behavior is what is approached by gases.

10. Deviations • When P is plotted vs. PV/nRT there is some deviation from an ideal gas. • What pressure is “most ideal”?

12. Deviations • When P is plotted vs. PV/nRT at different temps, which temp is closest to “ideal”

14. Where is it best? • The ideal gas law works best at LOW pressure and HIGH temperatures. • Why?

15. Correction Factors • Van der Waals developed an equation for real gases in 1873. • He corrected for the “ideal” gas assumptions.

16. Correction Factors • 1. Corrected for the volume by subtracting the volume of the molecules. nb

17. Correction Factors • 2. Accounted for the interactions by adding a “correction factor” that would help make the pressure higher.

18. Correction Factors • The size of the pressure correction factor depends on the number of particles present (represented as moles/L)

19. So... • Why high temps and low pressure??? • High temps- less time to interact

20. So... • Low pressure- take up less of available space

21. Effusion and Diffusion

22. Effusion • The passing of a gas through a tiny hole into an empty chamber

24. Graham’s Law • Found that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles

25. Diffusion • The mixing of gases

27. RMS Velocity • The root mean square velocity is an average velocity that we can find for particles

28. μRMS = (3RT/M)1/2 • The constants are: • R = 8.31 J/mol K • T = temp in Kelvin • M = mass of mole of gas in KILOGRAMS!