1. Kinetic Theory: all particles of matter are in constant motion. Particles of Matter: Smallest unit of pure substances, atoms or molecules.

2. Temperature: Average Kinetic Energy of a substance’s particles. SI Temperature Unit: Kelvin (K) No degree symbol Absolute Temperature scale, no negative values. 0 K = no motion of molecules No maximum value Other Temperature Scales Celsius: Based on freezing/boiling of water (0 and 100 Celsius) Fahrenheit: Arbitrary scale, limited use in science.

3. Temperature Conversions Fahrenheit/Celsius °C = 5/9 (°F – 32.0) °F= 9/5 (°C) + 32.0° Sample Problems: 1)Convert 20 °C to Fahrenheit 2)Convert 20 °F to Celsius Celsius/Kelvin K = °C + 273 Sample Problems: 1) Convert 20 °C to Kelvins 2) Convert 20 K to Celsius

4. States of Matter: Solids, Liquids, and Gases

5. Solids Strong attractions between particles in a solid Particles have orderly arrangement, vibrate in fixed locations Have a definite shape and volume

6. Liquids Particles are close together, but their arrangement is more random. More kinetic energy than solids. Low attraction between particles. Definite volume but no definite shape.

7. Gases Lots of space between particles. No attraction between particles. Particle move at high speeds, over 1,000 mi/h at room temperature. Have neither a definite volume nor shape.

8. Plasma Exist at high temperatures, make up 99% of all matter in the universe Gases consisting of atomic nuclei and electrons

11. Phase Change When a substance changes from one state of matter to another.

12. Phase Changes: Movement of Energy Exothermic = energy released Endothermic = energy absorbed Temperature does not change during a phase change.

13. Solid/Liquid Phase Changes Melting (solid to liquid) Endothermic Heat of fusion = energy absorbed as a substance melts Used to break bonds/attractions between molecules, NOT to raise temperature Freezing (liquid to solid) Exothermic Freezing water makes rigid bonds and makes ice less dense

15. Vaporization (liquid to gas) Endothermic Heat of vaporization: energy absorbed as a substance vaporizes, used to break bonds Vapor pressure: pressure caused by collisions of vapor with walls of a container

16. Liquid to Gas Phase Changes Evaporation Occurs at the surface of a liquid Below the liquid’s boiling point Some particles have enough energy to escape the liquid Boiling Occurs below the surface Bubbles form as vapor pressure exceeds atmospheric pressure.

17. Evaporation

18. ) Other Phase Changes Condensation gas to liquid, Exothermic (ex. Dew) Sublimation solid to gas, Endothermic (ex. Dry ice) Deposition gas to solid, Exothermic (ex. Frost on window)

20. The Gas Laws Gas Pressure: P = F/A P = Pressure measured in Pascals (Pa) F = Force A = Area More force = more pressure. More area = less pressure

21. Gas Pressure Caused by collisions of molecules on surface/walls of a container. More molecules = more collisions, more pressure.

22. Effects of Temperature on Gas Pressure If temperature is increased: increased kinetic energy increases speed of particles Pressure increases because number of collisions increases

25. Boyle’s Law If temperature is kept constant, pressure and volume are inversely proportional (If one increases, the other decreases) P 1 V 1 =P 2 V 2

26. Effects of volume on gas pressure If volume is decreased then: - More collisions with walls of container, more pressure. If volume is increased then: - Less collisions with walls of container, less pressure.

27. Charles’s Law If pressure is kept constant, then volume is directly proportional to temperature (in Kelvin) V 1 /T 1 =V 2 /T 2

29. Combined Gas Law