States of Matter

1. States of Matter • I will use the kinetic-molecular theory to explain the physical properties of gases, liquids and solids • I will compare types of intermolecular forces • I will explain how kinetic energy and intermolecular forces combine to determine the state of a substance • I will describe the role of energy in phase changes

2. 13.1 Gases • I will use the Kinetic-molecular theory to explain the behavior of gases • I will describe how mass affects the rates of diffusion and effusion • I will explain how gas pressure is measured and calculate the partial pressure of a gas • Vocabulary kinetic-molecular theory elastic collision • temperature Graham’s Law of effusion diffusion • Pressure barometer pascal atmosphere • Dalton’s law of partial pressures

3. Gases • Substances that are gases at room temperature • usually display similar physical preperties despite • their different compositions. • Why is there so little variation in behavior among gases?

4. Kinetic-Molecular Theory

5. Kinetic-Molecular Theory • Describes the behavior of gases in terms of • particles in motion • Makes several assumptions about gas particles: • Size • Motion • Energy

6. Gas Assumptions • Particle Size • Separated by empty space • Volume of particle = small • Volume of empty space = large • NO significant attractive or repulsive forces • Particle Motion • Constant random motion • Move in a STRAIGHT line • Collide with walls or other particles • Collisions are elastic (NO kinetic energy is lost, just transferred)

7. Gas Assumptions Cont... • Particle Energy • Affected by mass and velocity • KE = ½ mv2 • Single gas • Particles have same mass • Particles have different velocity • Particles have different KE • Temperature = measure of the average KE of the particles in a sample of matter • At a given temperature, ALL gases have the SAME average KE

8. Gas Behavior

9. Gas Behavior • Low Density • Compressible • Expandable • Diffuse • Effuse

10. Gas Behavior • Low Density • Density = mass/volume • Large space between gas particles • Fewer gas molecules than solid or liquid molecules in the same volume

11. Gas Behavior Compression & Expansion • Large amount of empty space between gas particles • Allows particles to be squished into a smaller volume • Stop squishing; random motion of particles fills the available space, expands to original volume

12. Gas Behavior Diffusion Particles diffuse from areas of high concentration to areas of low concentration Rate of diffusion Depends on mass of particles Lighter = faster Heavier = slower • The movement of one material through another material • Gas particles have no significant forces of attraction • Particles can slide past each other • Mix until evenly distributed

13. Gas Behavior

14. Gas Behavior Effusion • Related to diffusion • When a gas escapes through a tiny opening • Ex. Puncture a balloon or tire • Inverse relationship between effusion rates and molar mass

15. Gas Behavior Graham’s Law of Effusion Also applies to rate of diffusion Set up a proportion to compare the diffusion rates of 2 gases • The rate of effusion for a gas is inversely proportional to the square root of its molar mass

16. Finding a Ratio of Diffusion Rates • Let’s Try!

17. Practice Problems • RN2/RNe = 0.849 • RCO/RCO2 = 1.25 • 2.5mol/min Calculate the ratio of effusion rates for nitrogen (N2) and Neon (Ne). Calculate the ratio of diffusion rates for carbon monoxide (CO) and carbon dioxide (CO2) What is the rate of effusion for a gas that has a molar mass twice that of a gas that effuses at a rate of 3.6 mol/min? Use to find ratios

18. Gas Pressure

19. Pressure Gas Pressure • Force per unit area • Gas particles exert pressure when they collide with: • The walls of their container • Each other • Exert pressure in ALL directions (because particles move in all directions) Individual particles can only exert little pressure Many particles colliding can exert substantial pressure pressure of gases is what keeps our tires inflated, makes our basketballs bounce, makes hairspray come out of the can, etc.

20. Gas Pressure • Pressure increases when temperature increases because the molecules are moving with greater speed and colliding against the sides of their containers more often. • Therefore, the pressure inside that container is greater, because there are more collisions.

21. Gas Pressure Atmospheric Pressure Varies by elevation Mountains = less Sea Level = more • Air pressure at Earth’s surface • Equal to: • Pressure exerted by 1 kg mass on a square centimeter

22. Gas Pressure Measuring Air Pressure Air pressure varies because of: Changes in air temperature Changes in humidity Increase in air pressure = Hg rises Decrease in air pressure = Hg falls • Barometer • Measures atmospheric pressure • Mercury in it is always about 760mm • Exact amount determined by 2 forces • Gravity- downward force • Air pressure- upward force (air presses on surface of Hg)

23. Gas Pressure Measuring Enclosed Gas Pressure • Manometer • Measures pressure of enclosed gas • The difference in height of the mercury in the 2 arms is used to calculate the pressure of the gas in the flask

24. Gas Pressure Units of Pressure SI Unit = Pascal (Pa) Atmosphere = atm (used to report air pressure)