1. States of matter Solids and Liquids 1

2. Gases, Solids, and Liquids Phase Particle Properties SpacingEnergyMotionVolumeShape Solid Liquid Gas closelowvibrationaldefinite closemoderaterotationaldefiniteindefinite far aparthightranslationalindefinite 2

3. Other States of Matter Amorphous Solids Most solids with particles in repeating geometric patterns are crystals. Those with particles arranged randomly are amorphous. Glasses are one type of amorphous solid Plasmas a. Hot, ionized gas particles. b. Electrically charged. c. Most common state in universe. 3

4. Chumbler - Properties of Matter 4 Examples of Plasmas 4

5. 5 Plasmas Microscopic Explanation for Properties of Plasmas  Plasmas have an indefinite shape and an indefinite volume because the particles can move past one another.  Plasmas are easily compressible because there is a great deal of free space between particles.  Plasmas are good conductors of electricity and are affected by magnetic fields because they are composed of ions (negatively charged electrons and positively charged nuclei). 5

6. PHASE CHANGES Description of Phase Change Term for Phase Change Heat Movement During Phase Change Solid to liquid Melting Heat goes into the solid as it melts. Liquid to solid Freezing Heat leaves the liquid as it freezes. 6

7. PHASE CHANGES Description of Phase Change Term for Phase Change Heat Movement During Phase Change Liquid to gas Vaporization, which includes boiling and evaporation Heat goes into the liquid as it vaporizes. Gas to liquidCondensation Heat leaves the gas as it condenses. Solid to gasSublimation Heat goes into the solid as it sublimates. 7

8. Heating Curves The temperature of most pure substances is constant during a phase change. 8

9. Cooling Curves The temperature of most pure substances is constant during a phase change. 9

10. Heat of Fusion The heat required to convert a substance from the solid to the liquid phase is known as the heat of fusion The heat of fusion is a property of the substance. For water the heat of fusion is 335 Joules per gram 10

11. Heat of Vaporization The heat required to convert a substance from the liquid to the gas phase is known as the heat of vaporization The heat of vaporization for a substance depends on the temperature For water the heat of vaporization is about 2240 Joules per gram The heat required to vaporize a substance is generally much higher than the heat it takes to melt it. 11

12. Evaporation The molecular velocities of the particles in the liquid phase vary according to a Maxwell- Boltzman distribution The faster moving particles at the surface may escape the confines of the liquid entirely. Some particles in the vapor phase may be recaptured by the liquid. Since the higher energy particles are more likely to escape the average energy of the liquid particles is reduced. Evaporation is a cooling effect, while condensation is a warming effect 12

13. Vapor Pressure Explaining Vapor Pressure on the Molecular Level Explaining Vapor Pressure on the Molecular Level Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium. 13

14. Vapor Pressure and the Boiling Point Liquids boil when the external pressure equals the vapor pressure. The vapor pressure of a liquid increases with temperature The temperature of boiling point increases as pressure increases. There are two ways to get a liquid to boil: increase temperature or decrease pressure. Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. Normal boiling point is the boiling point at 760 torr (1 atm). 14

15. Gas-Liquid Equilibration 15

16. Vapor Pressure Volatility, Vapor Pressure, and Temperature 16

17. Phase Diagrams A Phase Diagram is a graph of pressure vs. Temperature summarizing all equilibria between phases. Given a temperature and pressure, phase diagrams tell us which phase(s) will exist. Key Features of a phase diagram: Vapor-pressure curve: generally as pressure increases, temperature increases. Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid Triple point: temperature and pressure at which all three phases are in equilibrium. Normal boiling and melting points (I.e. at 1 atm) Critical point: critical temperature and pressure for the gas. 17

18. Phase Diagrams Any temperature and pressure combination not on a curve represents a single phase. 18

19. Phase Diagram A phase diagram shows the relationship between the three phases of matter The boiling point of a substance depends on the pressure. The melting point is not significantly affected by the pressure 19

20. Phase Diagram The boiling point of a liquid is the temperature at which the vapor pressure of the liquid is equal to atmospheric pressure At the triple point all three phases are in equilibrium 20

21. Phase Diagram of H 2 O 21 The melting point curve slopes to the left because ice is less dense than water. Triple point occurs at 0.0098  C and 4.58 mmHg. Normal melting (freezing) point is 0  C. Normal boiling point is 100  C. Critical point is 374  C and 218 atm.

22. Phase Diagram of CO 2 Carbon Dioxide: Triple point occurs at -56.4  C and 5.11 atm. Normal sublimation point is -78.5  C. (At 1 atm CO 2 sublimes it does not melt.) Critical point occurs at 31.1  C and 73 atm. 22

23. Critical Temperature and Critical Pressure Gases liquefied by increasing pressure at some temperature. Critical temperature: the minimum temperature for liquefaction of a gas using pressure. Critical pressure: pressure required for liquefaction. 23

24. Critical Temperature 24

25. Specific Heat The ability of a material to absorb and release heat depends on its composition and makeup The heat required to raise the temperature of 1 gram of a material 1 o C is called the specific heat. For water the specific heat is 4.184 J g - 1o C -1

26. Phase Change --Problem 1 20.0 g of ice at -10.0 o C is heated until it melts and the is further heated to a final temperature of 40.0 o C. Calculate the total heat change for the ice. The heat of fusion of ice is 335 Jg -1. The specific heat of ice is 2.05 Jg -1 o C -1 and that of liquid water is 4.18 J g -1 o C -1.  Q = (20.0g)(10.0 o C)(2.05 Jg -1 o C -1 )+ (20.0g)(335 J g -1 ) + (20.0g)(4.18J g -1 o C -1 )(40.0 o C)  Q = 10454 joules or 10.5 kJ

27. Phase Change –Problem 2 50.0 g of water at 12.0 o C is added to 120.0 g of water at 84.0 o C. Calculate the final temperature of the water. Let T = final temperature Then (50.0g x (T- 12.0 o C)(4.18Jg -1 o C -1 ) =(120g)(84.0 o C -T)(4.18Jg -1 o C -1 ) 50T-600 = 10080 – 120 T 170 T = 10680 T = 62.8 o C