Phases of Matter and Phase Changes

1. Phases of Matter and Phase Changes

2. Phase • Depends on strength of forces of attraction between particles. .

3. Solids • Definite shape, volume. • Regular crystalline lattice structure. • Most dense phase (exception is water!). • Difficult to compress. • Highest attraction between particles. • Atoms vibrate in fixed positions • Note: Amorphous solids include glass, plastic, wax, and silly putty

4. Liquids • Definite volume • No definite shape • Hard to compress • Particles can slide past each other • Forces of attraction between particles still high

5. Gases • No definite shape or volume • Expands to fill container • Lowest density • Density depends on pressure • Little attraction between particles • “Vapor” = a gaseous state of something that is normally liquid (Ex: water vapor)

6. Phases Applet • Short Summary video on phases: • http://www.youtube.com/watch?v=s-KvoVzukHo&safe=active • http://www.harcourtschool.com/activity/states_of_matter/

7. Changes in Phase Gas Condensation Vaporization (Boiling or Evaporating) Liquid Solidification Melting (fusion) Solid

8. Let’s Skip a Phase • Sublimation • Directly from the solid phase to the gas phase. • Happens with substances with weak intermolecular forces of attraction • They separate easily! • Ex: CO2(s) dry ice, Iodine CO2(s) → CO2 (g) http://www.youtube.com/watch?v=J8mDGwf-5x0&safe=active

9. Energy • Energy = capacity to do work or produce heat. It can be anything that causes matter to move or change direction. • Ex: electrical, radiant, atomic, mechanical, magnetic, sound, chemical • Energy and the 4 states of matter: • http://www.youtube.com/watch?v=88tK5c0wgH4&safe=active

10. Law of Conservation of Energy • Energy can’t be created or destroyed, just transferred from one form to another

11. PE vs. KE • Potential Energy stored energy • Energy can be stored in bonds between atoms and released during chemical rxns. • Kinetic Energy energy of motion • All atoms are moving and vibrating unless at absolute zero

12. Heat Energy • A form of energy that increases the random motion of particles • Measured in Joules or calories. http://www.youtube.com/watch?v=f1eAOygDP5s&safe=active

13. Heat Flow • Heat energy travels from an object of higher temp. to one of lower temp. until both reach the same temp.

14. Temperature • Measure of the average kinetic energy (motion) of all the particles in a sample. • Not a form of energy!!! • But if you add heat energy or take it away, it causes particles to move faster or slower and thus changes the temp.

15. Temperature Scales Used in Chemistry Celsius • Fixed points of scale based on the freezing point and boiling point of water • 0 °C = water freezes, 100 °C = water boils Kelvin • Scale based on lowest temperature possible • 0 K = absolute zero

16. Temperature Scales and ConversionsK = ˚C + 273

17. Absolute Zero • Temperature at which particles have slowed down so much they no longer possess any kinetic energy. 0 Kelvin -273° Celsius

18. Heat vs. Temperature • Teacup vs. Bathtub • Both at 25˚C • Which one contains more heat energy? • Which one has the greater average KE?

19. Exothermic vs. Endothermic • All changes in matter are accompanied by changes in energy. • Exothermic Change:A + B → C + D + energy • Energy is released • Energy “ex”its • Endothermic Change:A + B + energy → C + D • Energy is absorbed • Energy “en”ters

20. Energy During Phase Changes • Endothermic: (s→l, or l→g) • Energy overcomes attractive forces between particles • PE increases • Exothermic: (g→l, or l→s) • As particles come closer together energy is released • PE decreases

21. Heating & Cooling Curves • Graphically represents temp. changes as heat energy is added or taken away.

22. Label This Graph

23. The slanted portions = temp is changing Single phase is heating up or cooling down KE is changing The flat portions = temp not changing Substance undergoing a phase change PE is changing Interpreting the Graph • http://mutuslab.cs.uwindsor.ca/schurko/animations/waterphases/status_water.htm

24. Heating Curve for Water

25. What is Melting Pt? Boiling Pt?

26. Heat Equations • Calculates the energy involved when a substance changes in temperature or undergoes a phase change.

27. Physical Constants for WaterTable B Use these constants in Heat Equations Hf = heat of fusion = 334J/g Hv = heat of vaporization = 2260J/g Specific Heat Capacity (“c”) = 4.18 J/g x K

28. When temp. of substance changes use this formula: Q = mcΔT

30. What is Specific Heat Capacity? Specific Heat: Joules of heat needed to raise 1 gram of a substance 1°C. • Substances have different abilities to absorb heat when energy is applied depending on their composition. Ex: Piece of Iron vs. Water.

31. When Undergoing Phase Change (Temp. constant) use one of these formulas: • Q = mHf Use when changing from solid to liquid (melting) or liquid to solid (freezing) • Q = mHv Use when changing from liquid to gas (vaporization) or gas to liquid (condensing)

32. Calorimeters • Instrument used to determine amount of heat lost or gained in a reaction by measuring changes in the temp. of water surrounding the system. Q = mcΔT

33. Multi-step Heat Problems (Honors) • Need to use more than one of the heat equations and add up the total heat. • Ex: Calculate the heat energy to raise 10 grams of water at -25°C to 80°C. • Draw a heating curve. Figure out # of steps. • 1.) Heat ice from -25° to 0° q = mcΔT • 2.) Melt ice to liquid at 0° q = mHf • 3.) Heat liquid water from 0° to 80° q = mcΔT

34. Heat Lost = Heat Gained (Honors) • When two objects of different temperatures are placed together in a closed system, heat will flow from the hotter to the colder object until they reach the same temperature. • The total heat lost = total heat gained mcΔT = mcΔT