1. Phases of Matter, Energy and Phase Changes

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

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

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

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

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

7. Phase Changes Short Summary video on phases: (1 min)
Applet: (Excellent)

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)

9. Energy
Energy = capacity to do work or produce heat. It can be anything that causes matter to move or change direction.
Many different types of energy
Ex: electrical, thermal, atomic, mechanical
“Chemical” energy is the potential energy stored in the bonds between atoms

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
Kinetic Energy energy of motion
All atoms are moving and vibrating unless at absolute zero

12. Energy and Changes to Matter
Exothermic Change: A + B → C + D + energy
Energy is released or “ex”its
Endothermic Change: A + B + energy → C + D
Energy is absorbed or “en”ters

13. Energy During Phase Changes
Solid Liquid or Liquid Gas
Endothermic
Energy is absorbed and overcomes attractive forces between particles
Add heat

14. Gas Liquid, Liquid Solid
Exothermic
As particles come closer together energy is released
Remove heat

16. Heat Energy
Also called Thermal energy, it makes particles move more as it is added
Measured in Joules or calories.

17. Heat Flow or Transfer
Heat energy travels from an object of higher temp. to one of lower temp. until both reach the same temp.

19. 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.
Heat vs Temp.

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

21. 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

22. Temperature Scales and Conversions K = ˚C + 273

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

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

25. Label This Graph

26. Interpreting the Graph
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

27. Heating Curve for Water

28. What is Melting Pt? Boiling Pt?

29. Heat Equations
Calculates the energy involved when a substance changes in temperature or undergoes a phase change.
Use this when temperature of substance changes use this formula:

30. When Undergoing Phase Change use one of these formulas: TEMPERATURE CONSTANT
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)

31. Physical Constants for Water Table 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

32. What is Specific Heat Capacity?
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.

33. Calorimeter
Instrument to determine amount of heat lost or gained in a reaction by measuring changes in the temp. of water surrounding the system.
Virtual Calorimetry
Q = mcΔT

34. Multi-step Heat Problems (Honors)
Need to use more than one of the heat equations and add up the total heat.
Note: Specific Heat of different phases of water!
H2O(s) = 2.10 J/gx°C
H2O (l) = 4.18 J/gx°C
H2O(g) = 1.84 J/gx°C
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

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

37. Try This!! Online App Demonstrates Specific Heat and Calorimetry