1. States Of Matter!

2. Gases – Kinetic Molecular Theory Explains the forces between molecules and the energy the molecules possess.

3. Kinetic Molecular Theory 1. Matter in all its form (solid, liquid, gas) is composed of extremely small particles. The space occupied by the gas particles themselves is ignored in comparison with the volume of space they occupy. 2. The particles of matter are in constant motion. In solids, this motion is restricted to a small space. In liquids, the particles have a more random matter. In gas, the particles are in continuous, random, straight line motion. 3. When these particles collide with each other or with the walls of a container there is no energy loss.

4. Behavior of Gases Low density Compressed and Expand Diffusion: Random motion of gas particles cause gases to mix until they are evenly distributed Effusion: Gas escapes through a tiny opening. Thomas Graham did experiments to measure rate of effusion for different gases at the same temperature.

5. Graham’s Law of Effusion Discovered inverse relationship between effusion rates and molar mass. Rate of effusion: Grahams Law relates to diffusion as well. Heavier particles diffuse more slowly then lighter particles. We can set up a proportion to compare diffusion rates. Rate A = Sqrt(molar mass B/ Molar Mass A) Rate B

6. Example Ammonia has a molar mass of 17.0 g/mol; hydrogen chloride has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

7. Gas Pressure Pressure is defined as force per unit area. A barometer can be used to measure the pressure of a gas in a confined container.

8. Gas Pressure Cont… One can also used an apparatus called the manometer. A manometer is a U-tube containing mercury or some other liquid.

9. Daltons Law of Partial Pressure Dalton discovered that each gas in a mixture exerts pressure independently of the other gases present. Law of Partial Pressure: total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. P total = P 1 + P 2 + P 3 + … P n

10. Example What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600mm Hg and the partial pressure of helium is 439 mm Hg?

11. Forces of Attraction Dispersion Forces: Weak forces that result from temporary shifts in the density of electrons in electron clouds. Dipole-Dipole forces: Attractions between oppositely charge regions of polar molecules

12. Forces Of Attraction Cont… Hydrogen Bonds: Intermolecular forces that result from a specific interaction between a hydrogen atom in one molecule and a fluorine, oxygen, or nitrogen atom in another molecule.

13. Liquids Kinetic Molecular Theory predicts constant motion of the liquid particles. Attraction between liquid particles limit their range of motion. Denser than gases. Liquids can be compressed, but change in volume is smaller then gas, because liquid particles are already tightly packed together. A liquid can diffuse through another liquid, but is slower then gas at the same temperature.

14. Liquids Cont… Viscosity: Friction/resistance to motion that exists between the molecules of a liquid when they move past one another. The stronger the attraction between the molecles of a liquid, the greater its resistance to flow – greater resistance to flow. Surface tension: Uniformly distributed attractive forces. Uneven forces creates a “film” on the surface. Capillary action: Cohesion/Adhesion

15. Solids The particles in a solid are more closely packed than those in a liquid. Exception: ICE! Crystalline Solid: solid whose atoms, ions, or molecules are arranged in a orderly, geometric, 3D structure. Example: ICE. Atomic, molecular, covalent, ionic, and metallic.

16. Phase Changes When energy is added or removed from a system, one phase can change into another. Melting (E) Vaporization (E) Sublimation (E) Condensation (-E) Deposition (-E) Freezing (-E)

17. Melting Lets think about ICE Energy absorbed by ice is used to disrupt the hydrogen bonds holding the water molecules together. Amount of energy required to melt one mole of solid depend on the strength of the forces holding the molecules together.

18. Vaporization Particles that escape from the liquid enters the gas phase. AKA, Vaporization When vaporization occurs only at the surface of a liquid, the process is called evaporation We control our body temperature by evaporation.

19. Sublimation The process which a solid changes directly to a gas without first becoming a liquid. i.e Freeze drying

20. Condensation Process by which a gas or a vapor becomes a liquid. Formation of hydrogen bonds  energy is released. Causes? Contact with cold surface  from droplets or dew Layer of new air near the ground cools  produces fog.

21. Deposition and Freezing Deposition Substance changes from a gas/vapor to a solid without becoming a liquid. Energy is released Example: FROST, SNOWFLAKES Freezing Heat is removed, molecules loose kinetic energy. Hydrogen bonds keep molecules fixed/frozen. Converting into a crystalline solid.

22. Phase Diagrams Graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure.