1. States of Matter and Phase Changes

2. Kinetic Theory of Matter: Matter is made of particles that are in constant motion – Describes how close together the molecules are in a solid, liquid, and gas – The relative motion of the molecules – and the attractive forces between the molecules

4. Solid (Lowest kinetic energy) Particle Arrangement, energy, and relative motion Particles close together Often arranged in an orderly 3-dimensional pattern Strong attractions between particles Only motion of particles is to vibrate in place energy that binds particles is great

5. Solids cont’d Shape Definite will NOT take the shape of container strong attractive forces and the closeness and rigid arrangement of particles do not allow particles the freedom to move around one another

6. Volume Fixed (definite) volume strong chemical bonds or intermolecular forces holding particles together

7. Solids cont’d Density Highest* density (particles are very close together) *Water is an exception! Ice is actually less dense than liquid water- that’s why it floats!

8. Solids cont’d Compressibility (ability to force to a smaller volume) Not compressible

9. Liquid Particle Arrangement, energy, and relative motion Particles close together Somewhat randomly arranged Weaker attractions between particles Particles able to move around each other-vibrate and rotate/ not locked in position energy that attracts particles together is weaker than in solids

10. Liquids cont’d Shape Indefinite Attractive forces weak enough to allow particles to move around each other (flow) to take shape of container

11. Liquids cont’d Volume Fixed (definite) volume because particles are close together particles have strong enough attractive forces between them to hold them close together

12. Liquids cont’d Density High density (compared to gases) because of closeness of particles Compressibility Not very compressible

13. Gas (Highest kinetic energy) (Vapor-the gaseous state of a substance that is normally a liquid or solid at room temperature)

14. Gases cont’d Particle Arrangement, energy, and relative motion Particles are far apart randomly arranged very weak attractions between particles that have little to no effect Particles are free to vibrate, rotate, and translate travel in straight line paths until a collision with another particle energy of attraction between particles is minimal to none

15. Gases cont’d Shape Indefinite Very weak attractive forces and distance between particles allow them to freely flow and take shape of container

16. Gases cont’d Volume Indefinite volume particles are not close together and experience very weak or no attractive forces between them, so they will move apart to fill container

17. Gases cont’d Density Low density particles are far apart Compressibility Very compressible

18. Plasma- the fourth state of matter; not found on earth under normal circumstances, but is the most abundant state of matter in the universe; found only at extremely high temperatures (core of star); electrons are stripped off of nuclei; very energetic particles Bose-Einstein Condensate-fifth state of matter existing at temperatures close to absolute zero (0 Kelvin or -273 Celsius). Groups of atoms behave as though they are a single particle; very low energy.

19. Phase Changes (Changes of State) Phase change- a reversible physical change that occurs when a substance changes from one state of matter to another.

21. Energy is either released or absorbed during a phase change – Changes that release energy to the surroundings are called exothermic changes. Freezing, deposition, and condensation are exothermic changes – Changes that absorb energy from the surroundings are called endothermic changes. Melting, vaporization, and sublimation are endothermic

22. Phase Changes Endothermic Changes (Energy is required to overcome intermolecular forces) – Solid  Liquid (melting or fusion) – Liquid  Gas (vaporization, evaporation, or boiling) – Solid  Gas (sublimation) Exothermic Changes (Energy is released as intermolecular forces are formed) – Liquid  solid (freezing) – Gas  Liquid (condensation) – Gas  Solid (deposition)

23. MELTING and FREEZING occur at the same temperature CONDENSATION and VAPORIZATION occur at the same temperature SUBLIMATION and DEPOSITION occur at the same temperature

24. The temperature of a substance does NOT change DURING a phase change- the energy is being used to overcome or form attractive forces between the particles.

25. Heating and Cooling Curves Graphs of temperature vs time for the heating or cooling of a substance

26. Heating Curve Areas where slope is positive kinetic energy of molecules change / heat energy speeds up the molecules. Areas where slope is zero phase change potential energy of molecules change the heat energy is used to overcome intermolecular forces to separate the molecules from solid to liquid or liquid to gas.

27. Heating Curves Describe what is happening in each of the following time intervals: 1-2 mins: The temperature of the solid is rising 2-5 mins: PHASE CHANGE: The solid is melting into a liquid 5-10 mins: The temperature of the liquid is rising 10-14 mins: PHASE CHANGE: The liquid is vaporizing into a gas 14-16 mins: The temperature of the gas is rising.

28. Cooling Curves Note that cooling a substance has the opposite effect Areas where slope is negative kinetic energy of molecules change / heat energy is lost to surroundings Areas where slope is zero phase change potential energy of molecules change the heat energy is used to form intermolecular forces to join the molecules the molecules from gas to liquid or liquid to solid

29. Cooling Curve Describe what is happening in each of the following time intervals: 1-3 mins: 3-7 mins: 7-12 mins: 12–15 mins: 15-16 mins:

30. Energy in phase changes solid  liquid or liquid  solid ΔH f =q=L f m – ΔH f is called the Heat or enthalpy of fusion (q=ΔH f ) and represents the amount of heat given off when the substance melts – L f is called the Latent Heat of Fusion The amount of heat it takes to melt one gram of a substance

31. liquid  gas or gas  liquid ΔH f =q=L v m – ΔH v is called the Heat or enthalpy of vaporization (q=ΔH f ) and represents the amount of heat given off when a substance vaporizes – L v is called the Latent Heat of vaporization The amount of heat it takes to vaporize one gram of a substance

32. Note that heat of fusion is smaller than heat of vaporization – Takes less energy to allow particles to move around each other than to separate them completely

33. To calculate the total energy change…. On the parts of the curve that have a positive or negative slope, use q=mcΔT On the zero slope parts of the curve (the phase changes), use q=L v m or q=L f m Then add together your q values for each section

34. Phase Diagrams A graph that gives the conditions of temperature and pressure at which a substance in a sealed container exists as a solid, liquid, or gas.

35. Solid phase- on left Liquid phase- in the middle Gas phase- on the right

36. Phase boundaries – The conditions of temperature and pressure at which two phases exist in equilibrium are indicated by a line separating the phases. – Both phases are present

37. Triple Point – The only set of conditions at which all three phases can exist in equilibrium with one another – The point where all three phases meet

38. Normal Boiling Point: the temperature where liquid changes to gas at normal atmospheric pressure (1 atm or 760mm Hg) – This is also the normal condensation point Normal Freezing Point: the temperature where liquid changes to solid at normal atmospheric pressure – This is also the normal melting point

39. critical temperature: temperature above which the vapor cannot be liquefied. critical pressure: pressure required to liquefy AT the critical temperature. critical point: critical temperature and pressure coordinates (for water, T c = 374°C and 218 atm).

41. Most substances have a solid-liquid line that has a positive slope since their solid phase is more dense than the liquid. This one is for carbon dioxide.

42. Water is a freak! The solid-liquid line tilts to the left [negative slope] since its solid is less dense than its liquid phase—ice floats. Usually the solid sinks as it is more dense.

43. http://treefrog.fullerton.edu/chem/LS/phased.h tml