Kinetic Molecular Theory and Gas Laws Day 1

1. Kinetic Molecular Theory and Gas LawsDay 1

2. Kinetic-Molecular Theory – explains how particles in matter behave • All matter is composed of small particles that are far apart. Gas is mostly empty space. • Particles are in constant, random motion. • Particles collide with each other and walls of their containers: collisions create pressure 4. Collisions are elastic = no KE lost 5. No attractive/repulsive forces between particles. Molecules move in straight lines.

4. EXIT QUESTIONS: 1) Each of these flasks contains the same number of molecules. Which container has the highest pressure? Explain your answer.

5. EXIT QUESTIONS: • Which of the following changes to a system will NOT result in an increase in pressure? Explain why you chose your answer. • Increasing the volume of container • adding more gas molecules • Decreasing the volume of the container • Raising the temperature

6. Kinetic Molecular Theory and Gas LawsDay 2

7. Factors affecting gases • Volume – amount of space an object occupies • Measured in milliliters (mL) or Liters (L) • 1000 mL = 1 L • We already have heard that 1 mol = 22.4 L @ STP • The more moles we have the bigger the balloon will need to be!

8. Example 1 How many moles of nitrogen gas are in 89.6 L at STP? 1 mol N2 89.6 L N2 22.4 L N2 = 4.00 mol N2

9. Example 2:What volume does 76 grams of fluorine (F2) occupy at STP (normal conditions)? 76 g F2 1 mole F2 22.4 L F2 1 mole F2 37.996 g F2 44.8 L F2 **45 L =

10. 2. Temperature • Average kinetic energy of particles (how fast they go) • Measured in Kelvin • K = oC + 273 Ex: Convert 17oC to Kelvin: 17oC + 273 = 290 K

11. *C= 5/9 (*F-32) • *F= 9/5 (*C) + 32 • K= *C + 273 • *C= K- 273

12. 3. Pressure • Force exerted by a gas per unit area on a surface. Example: Pounds/in2 or psi • Results from the simultaneous collisions of billions of gas particles with the walls of the vessel containing the gas. Standard pressure: 760 mm Hg = 1 atmosphere = 101.3 kPa = 29.92 in. Hg = 14.7 psi = 760 torr

13. Measuring Atmospheric Pressure • Measured with a barometer. • A barometer uses a column of mercury that rises to an average height of 760 mmHg at sea level. • 1 atmosphere (1 atm)

14. Standard Temperature and Pressure (STP) • The conditions of standard temperature and pressure are = 1.0 atm pressure and = 273 K (or 0C). @STP 1 mole of gas = 22.4 L of gas

15. Example 1 The atmospheric pressure in Denver, CO is 0.830 atm on average. Express this pressure in mm Hg.

16. 0.83 atm  mm Hg 1 atm = 760 mm Hg 0.83atm 760 mm Hg 1 atm = 630.8 mm Hg **631 mmHg

17. Example 2 Convert a pressure of 175 kPa to atmospheres.

18. 175 kPa  atm • 101.3 kPa = 1 atm 175 kPa 1 atm 101.3 kPa 1.72 atm

19. Gas Law Foldable • Fold the left and right to the middle. • Cut along solid lines (but only to the crack!)

20. Dalton’s Law of Partial Pressure The pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases. PTotal = P1 + P2 + P3….

21. Example A balloon is filled with air (O2, CO2, & N2) at a pressure of 1.3 atm. If PO2 = 0.4 atm and PCO2 = 0.3 atm, what is the partial pressure of the nitrogen gas?

22. PTotal = P1 + P2 + P3…. Ptotal = PO2 + PCO2 + PN2 1.3 atm = 0.4 atm + 0.3 atm + PN2 PN2 = 0.6 atm