1. CHAPTER 10 – Gases Lecture 1 – KMT, Graham’s & Dalton’s Law

2. Gases – Kinetic Molecular Theory
Gas particles are very small and are very far apart.
Attractive/repulsive forces are negligible
All gases have similar PHYSICAL properties even though they have different CHEMICAL properties

3. IDEAL GASES We assume most gases have “ideal” behavior:
Gases are tiny particles that are far apart and in constant, random motion.
Collisions between gas particles and other objects are elastic (no kinetic energy is lost and no attractions between particles).

4. IDEAL GASES
The average kinetic energy of gases depends on temperature, not the identity of the particle.
All gases have similar PHYSICAL properties even though they have different CHEMICAL properties

5. Kinetic Energy of Gas Molecules
m = mass
v = velocity
**At any given temperature, the molecules of all gases have the same AVERAGE kinetic energy.

6. Kinetic Molecular Theory
Only a change of temperature will change the average kinetic energy of gases.

7. Kinetic-Molecular Theory
At the same temp, larger molecules (with more mass) move more slowly.

8. Diffusion of Gases
Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout due to constant random motion of gases.

9. Diffusion of Gases
Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout due to constant random motion of gases.

10. Diffusion of Gases
Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout due to constant random motion of gases.

11. Effusion
The escape of gas molecules through a tiny hole into an evacuated space.

12. Effusion

13. GRAHAM’S LAW OF EFFUSION
Note: GRAHAM’S LAW CALCULATIONS ARE NOT INCLUDED ON THE AP TEST. ONLY QUALITATIVE ANALYSIS.
The rate of effusion or diffusion is inversely proportional to the molar mass of the molecule. (OR SMALLER MOLECULES DIFFUSE/EFFUSE MORE RAPIDLY)
Mixture of He (4 g/mol) and Ne (20.2 g/mol):
only until 2:38 min.

14. Review:
1. At the same temperature, all molecules have the (same, higher, less) average kinetic energy.
same
If you double the pressure of a gas, at constant temperature, how does the average kinetic energy change?
no change

15. Review: N2 CH4 No, only a change in temp. changes velocity
3. If you compress the gas to half the original volume at constant temperature, will the molecules increase in speed?
No, only a change in temp. changes velocity
4. At constant temperature, which molecule or atom will have the higher speed?
a) O2 or N2
b) CH4 or H2S N2
CH4

16. What Causes Pressure of a Gas?
Pressure = Force / Area
Gas particles exert pressure when they collide with the walls of an object.

17. Pressure
Atmospheric pressure is the weight of air per unit of area.

18. Dalton’s Law of Partial Pressures
In a mixture of gases, the total pressure is the sum of the partial pressures of the component gases.
Three gases are combined in container T. The pressure that each gas exerts is independent of the pressure exerted by the other two gases. The pressure in container T is the sum of the pressures in containers A, B, and C. Interpreting Diagrams What is the relationship between the number of particles in containers A and C and the partial pressures of the gases in A and C?

19. Dalton’s Law – Partial Pressures
The contribution each gas in a mixture makes to the total pressure is called the partial pressure exerted by that gas.
Ptotal = PA + PB + PC
Ptotal = 100 kPa kPa kPa
Ptotal = 550 kPa

20. Partial Pressures and Mole Fractions

21. Mole Fractions Problem
Calculate the partial pressure of each gas in the mixture if the total pressure is 745 torr. XO2 = 0.18, XCO2 = 0.015, XAr = 0.805