1. Ch. 13 States of Matter 13.1 Nature of Gases

2. I. Kinetic Theory A. Kinetic energy (K.E.): energy related to motion B. Kinetic theory assumptions about gases: –1. Small, far apart particles, no attractive or repulsive forces –2. Particles move fast and straight in random directions, only change directions during collision – 3. During collision, all K.E. exchanged, none lost (“elastic collision”)

3. II. Gas Pressure A. Force exerted by a gas on a surface B. Vacuum: space with no particles, no pressure C. Atmospheric pressure: gravity pulling air particles down D. Barometer: measures atmospheric pressure

4. III. Measuring Atm. Press. A. Pascal (Pa): SI unit for atmos. press. B. Standard atmosphere (Atm): unit for pressure based on sea level = 1 atm C. mm Hg: unit based on mercury barometer 1 atmosphere = 760 millimeters Hg = 101.3 kiloPascals

5. IV. Temperature A. Increased particle motion with increased Kelvin temperature B. 200 Kelvin has twice K.E. of 100 Kelvin C. No motion at 0 Kelvin “absolute zero” D. Temp. of a sample represents average of particles

6. 13.2 The Nature of Liquids

7. I. Liquid Model A. Particles closer, more dense than gas B. Intermolecular forces (between molecules) hold liquid particles together

8. II. Liquid to Gas A. Vaporization: changing to gaseous state B. Evaporation: vaporization below boiling pt. C. Cooling process: when fastest particles removed remaining particles have lower KE (cooler)

9. III. Vapor Pressure B. Vapor pressure increases with more heat C. Manometer: measures vapor pressure - If atm pressure greater: P vapor = P atmosphere – ΔP - If vapor pressure greater: P vapor = P atmosphere + ΔP A. Force of gas above a liquid or solid

10. IV. Boiling Point A. Temp. when vapor pressure ≥ external pressure B. Lower external pressure, lower B.P. C. B.P. depends on strength of intermolecular forces D. Adding particles increases boiling pts. due to interrupting molecular attractions

11. 13.3 The Nature of Solids

12. I. Solid Model A. Particles vibrate in fixed points B. Highly organized structures C. Melting point: temp. of solid to liquid D. K.E. breaks attractions keeping particles in fixed positions

13. II. Crystal Structure A. Most solids form organized crystals

14. III. Unit Cell A. Smallest group of particles within the crystal retaining the crystal shape B. Three types: Simple Cubic Body- centered Cubic Face-centered cubic

15. IV. Determining Density from Unit Cell A. Density is mass/volume B. Mass of unit cell: (molar mass/6.02x10 23 ) x # atoms C. Volume of unit cell: (side of unit cell) 3 D. Side of unit cell can be determined from atomic radius E. Determine side from radius using Pythagorean theorem (A 2 + B 2 = C 2 )

16. V. Other Solid Structures A. Allotropes: multiple forms of same element in same state B. “C” (diamond, graphite, Buckminsterfullerene) C. Amorphous solids: no crystal structure, random D. Ex. Glass, plastic, rubber, asphalt

17. 13.4 Changes of State

18. I. Phase Diagrams A. Shows conditions of temp. and press. at which substance is solid, liquid or gas B. Lines represent equilibrium between phases C. Triple point: when three states exist together D. Normal boiling/ melting pts. Based on 1 atm pressure

19. II. Sublimation A. Change of a solid to a gas B. Opposite process called “crystallization” C. When vapor pressure of solid is high enough to overcome atmospheric pressure D. Process used to separate/purify mixtures, freeze- drying foods

20. III. Water vs. CO 2 A. At sea level, H 2 O changes to all states at diff. temp. B. CO 2 doesn’t become a liquid at standard pressure C. H 2 O only substance with negative slope of solid/liquid line: as pressure increases ice melts