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. Solids form of matter that has a definite shape and definite volume.
Use (s) to denote solids in chemical reactions

5. Solids
In most solids the atoms, ions, or molecules are packed tightly together
The particles in solids tend to vibrate around fixed points
When you heat a solid, its particles vibrate more rapidly, eventually the solid breaks down and melts.

6. Types of Solids Crystalline Solids
In a crystal the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice. There are many different shapes of crystalline solids, pg 397

7. Types of Solids Non-Crystalline Solids
Amorphous solids lack an orderly internal structure.
Ex – Rubber, plastic, and asphalt.
Glass – transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. Sometimes called super-cooled liquids.

9. Allotropes
Two or more different molecular forms of the same element in the same physical state
Different properties because they have different structures

10. Allotropes of Carbon

11. Liquids
form of matter that has a definite volume, indefinite shape, and flows.
Use (l) to denote liquids in chemical reactions

12. Liquids
In liquids the atoms or molecules are able to slide past each other.
In liquids there are intermolecular attractions between the atoms or molecules, which determine the liquid’s physical properties.
When you heat a liquid the particles vibrate more rapidly and start moving past each other faster.

13. Gases
form of matter that takes both the shape and volume of its container
Use (g) to denote gases in chemical reactions

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

15. Phase Changes
During any given 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

16. 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

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

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

20. Phase Diagram of CO2

22. Energy What is energy? Two main types Capacity to do work
Ability to do work
Two main types
Kinetic
Potential

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

24. Temperature How is energy related to Temperature?
What happens to the temperature of a substance when you add energy?
Particles move faster
Temperature increases

25. Temperature
Relationship between energy, particle speed, and temperature
Temperature Definition
Average Kinetic Energy

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

27. Temperature Conversion
K = °C + 273

28. Pressure What is pressure? Physics – Force per unit area
Chemistry – related to the number of collisions between particles and container walls

29. Pressure Conversion
1 atm = kPa

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

31. 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

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

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

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

37. Thermochemistry
Chapter 17

38. Thermochemistry
Heat involved with chemical reactions and phase changes

39. 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

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

41. 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*°C) for liquid water
Listed in Table B of Reference Tables

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

43. Heat Equation Heat, Q Mass, m Specific Heat Capacity, c
Change in Temperature, ΔT
Q=m*c*ΔT

44. Example
200g of water is heated from 20°C to 40°C, how much heat is needed?
Q = m*c*ΔT
Q = (200g) * (4.18J/g°C) * (20°C)
Q = J

45. Example
How much energy is required to raise the temperature of 50g of water from 5°C to 50°C?
Q = m*c*ΔT
Q = (50g) * (4.18J/g°C) * (45°C)
Q = 9405 J

46. 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 = m*c*ΔT
180J = (10g) * c * (40°C)
c = 0.45 J/(g*°C)

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

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

50. 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

51. 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

52. 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

53. Examples How much energy is needed to melt 10g of ice at 0°C? Q = m*Hf
Q = (10g) * (334J/g)
Q = 3340 J

54. Example How much energy is needed to vaporize 10g of water at 100°C?
Q = m*Hv
Q = (10g) * (2260J/g)
Q = J

55. 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

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

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

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

61. 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

62. 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

63. 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

64. Example
A 100g block of aluminum, c=0.90J/g*°C 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
m*c*ΔT = m*c*ΔT
100g*0.90J/g°C*(100°C-Tf) = 50g*4.18J/g°C*(Tf-20°C)
90*(100-Tf) = 209*(Tf-20)
Tf = 209Tf-4180
13180 = 299Tf
Tf = 44°C

65. Conservation of Energy
System
Work Done (Energy)
Energy In
Energy Out

66. Conservation of Energy
Work Done (Energy)
Energy In
Energy Out

67. Conservation of Energy
Metabolism
Food In
Waste Out