1. Phases and Heat
Chapters 13 & 17

2. Phases of Matter
Chapter 13

3. Phases There are three phases, or states, that we will discuss Solid
Liquid
Gas

4. Phases Solid(s) Liquid(l) Gas(g)
form of matter that has a definite shape and definite volume.
Liquid(l)
form of matter that has a definite volume, indefinite shape, and flows.
Gas(g)
form of matter that takes both the shape and volume of its container

5. Phases
In most solids the atoms, ions, or molecules are packed tightly together (regular geometric pattern)
The particles in solids only have enough space to vibrate
In liquids the atoms or molecules are able to slide past each other.
The particles in a gas are spread very far apart

6. Heating When you heat a solid, the particles vibrate more rapidly.
Eventually the particles gain enough energy to push past each other forming a liquid
When you heat a liquid the particles vibrate more rapidly and start moving past each other faster.
Eventually the particles gain enough energy to overcome the intermolecular forces of attraction and become a gas

7. Phase Changes Six Changes Solid  Liquid Melting
Liquid  Solid Freezing
Liquid  Gas Vaporization
Gas  Liquid Condensation
Solid  Gas Sublimation
Gas  Solid Deposition

8. Phase Changes
During a phase change, both phases can exist together in equilibrium
Example
At 0°C, water can exist in both the liquid and solid phases in equilibrium

9. Energy
When energy is added to a reaction, or phase change, it is called Endothermic
When energy is released during a reaction, or phase change, it is called Exothermic

10. Phase Changes
Which phase changes are endothermic, requiring the addition of energy?
Melting
Vaporization
Sublimation

11. Phase Changes Which phase changes are exothermic, releasing energy?
Freezing
Condensation
Deposition

12. Short Cut S L g
Endothermic
Exothermic

13. Phase Diagrams A phase diagram is a graph of Pressure vs. Temperature
Shows the phases and boundary lines between them
Solids to the left side, gases to the right and bottom, liquid between them
Crossing a boundary line means undergoing a phase change
Examples on next 2 slides

15. Phase Diagram of CO2

16. Allotropes
Two or more different molecular forms of the same element in the same physical state (phase)
Different properties because they have different molecular structures
O2 vs. O3
Diamond, Graphite, Fullerenes (pictured on next slide)

17. Allotropes

19. Energy What is energy? Two main types
Capacity to do work (textbook def)
Ability to do something (reworded def)
Two main types
Kinetic
Potential

20. Types of Energy Kinetic Energy Potential Energy Energy of motion
Related to the speed and mass of molecules
Potential Energy
Stored energy

21. Temperature
Temperature is related to heat, but it doesn’t measure heat
Temperature is a measure of the average kinetic energy
When a substance is heated
Particles move faster
Temperature increases

22. Temperature Scales Kelvin (K) and Celsius (°C) scales
Kelvin scale is called the absolute scale
Directly related to the kinetic energy of a substance
Celsius scale is a relative scale
based on the boiling and freezing points of water

23. Temperature Conversion
K = °C + 273

24. Pressure Physics – Force per unit area
Chemistry – related to the number of collisions between particles and container walls

25. Pressure Conversion
1 atm = kPa

26. Vapor Pressure
Pressure exerted by vapor that has evaporated and remains above a liquid
Related to temperature
As temperature increases, vapor pressure increases

27. Boiling vs. Evaporation
Vapor pressure equals external, or atmospheric pressure
Evaporation
Some molecules gain enough energy to escape the liquid phase
At temp. less than boiling point

28. Normal Boiling Point Boiling Point at Standard Pressure
1 atm or kPa

29. Evaporation Why is evaporation considered a cooling process?
As the molecules with higher kinetic energy evaporate, the average kinetic energy of the substance decreases

31. Table H
Shows the relationship between temperature and vapor pressure for four specific substances

32. Table H
Shows the boiling points at different pressures

35. Thermochemistry
Chapter 17

36. Thermochemistry
Heat involved with chemical reactions and phase changes

37. Heat
Energy transferred from one object to another, usually because of a temperature difference
Measured in Joules (J) or calories (cal)
Heat flows from hot to cold

38. Heat Transfer Endothermic Exothermic Energy being added
Energy being released

39. Specific Heat Capacity
Amount of heat needed to raise the temperature of 1 g of a substance by 1°C
Unique for each phase of each substance
4.18 J/(g•K) for liquid water
Listed in Table B of Reference Tables

40. Heat What factors affect the amount of heat transferred?
Specific Heat Capacity
Mass
Temperature difference between objects

41. Heat Equation Heat, q Mass, m Specific Heat Capacity, C
Change in Temperature, ΔT
q=mCΔT

42. Example
200g of water is heated from 10°C to 30°C, how much heat is needed?
q = mCΔT
q = (200g) (4.18J/g•K) (20°C)
q = J

43. Example
How much energy is required to raise the temperature of 50g of water from 5°C to 50°C?
q = mCΔT
q = (50g) (4.18J/g•K) (45°C)
q = 9405 J

44. Another Example
What is the Specific Heat Capacity of Fe, if it takes 180J of energy to raise 10g of Fe from 10°C to 50°C?
q = mCΔT
180J = (10g) C (40°C)
C = 0.45 J/(g•K)

46. Phase Change At what temperature does ice melt?
At what temperature does water freeze?
Melting point and freezing point are the same

47. Phase Change What happens to temperature during phase changes?
Temperature remains constant
Temperature remains CONSTANT during a phase change

48. Phase Change If energy is being added, what kind of energy is it?
Energy being added is potential energy, not kinetic energy
Potential energy is being used to separate or spread the particles apart

49. Heat of Vaporization, Hv
Amount of energy needed to vaporize 1g of a substance
Water = 2260 J/g
q=mHv
Use for Liquid  Gas or
Gas  Liquid

50. Heat of Fusion, Hf Amount of energy needed to melt 1g of a substance
Water = 334 J/g
q=mHf
Use for Solid  Liquid or
Liquid  Solid

51. Examples How much energy is needed to melt 10g of ice at 0°C? q = mHf
q = (10g) (334J/g)
q = 3340 J

52. Example How much energy is needed to vaporize 10g of water at 100°C?
q = mHv
q = (10g) (2260J/g)
q = J

53. Phase Change Which requires more energy melting or vaporization? Why?
Molecules are spread farther apart as a gas
It takes more energy to get gas particles spread apart

55. Heating (Cooling) Curves
Shows relationship between temperature and time during constant heating or cooling.
Also shows phases, and the phase changes between them.

56. Heating Curves Diagonal lines are phases
Horizontal lines are phase changes
Time (s)
Temp (˚C)
Gas
Liquid
Solid

57. Heating Curves Diagonal lines are phases
Horizontal lines are phase changes
Vaporization
Condensation
Time (s)
Temp (˚C)
Melting
Freezing

59. Conservation of Energy
Energy can not be created or destroyed, only transferred or converted from one form to another.
Energy lost by one object must be gained by another object or the environment
qlost = qgained

60. Example
A chunk of iron at 80°C is dropped into a bucket of water at 20°C.
What direction will heat flow?
From the iron to the water
Hot to cold

61. Example
A chunk of iron at 80°C is dropped into a bucket of water at 20°C.
What could be the final temperature, when they both come to equilibrium?
Between 20°C and 80°C

62. Example
A 100g block of aluminum, c=0.90J/gK at 100°C is placed into 50g of water at 20°C, what will be the final temperature when the aluminum and water reach equilibrium?
qlost = qgained
mCΔT = mCΔT
(100g)(0.90J/gK)(100°C-Tf) = (50g)(4.18J/gK)(Tf-20°C)
90(100-Tf) = 209(Tf-20)
Tf = 209Tf-4180
13180 = 299Tf
Tf = 44°C