1. For your lab… What should you have in your conclusion? What is the point of “concluding”?

2. Review: Behavior of Gases Low density Compression & Expansion Diffusion & Effusion

3. Diffusion Tendency of molecules to move from areas of high concentration to low concentrations Perfume Example Lighter particles diffuse faster (KE = ½ mv 2 ) Think back to the 5 assumptions we make about gases… All have the same amount of KE

5. Effusion Gas escaping through a tiny hole

6. Graham’s Law of Effusion A gas will effuse at a rate that is inversely proportional to the square root of its molar mass Rate A/Rate B = √(Molar mass B/Molar Mass A)

7. Which gas diffuses faster? Ammonia or hydrogen chloride?

8. Gas Pressure Pressure = Force/Area Barometer

9. Dalton’s Law of Partial Pressure Gas pressure results from the collisions of particles If more particles are present in an specific area, then there is more pressure. Depends on number of moles of gas, size of container, and temperature of mixture

10. If you have a gas that is a mixture of several gases (like our air), then each gas produces a certain pressure. Partial Pressure- each gas’s pressure

11. Dalton’s Law of Partial Pressures P total = P 1 + P 2 + P 3 … Total pressure equals the sum of each pressure OR (moles gas/ total moles) x P total

12. Practice A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. If the partial pressure of carbon dioxide is 0.70 atm and the partial pressure of nitrogen is 0.12 atm, what is the partial pressure for oxygen?