1. Chapter 4
Heat and Temperature

2. Definition of Heat and Temperature
Heat is a measure of the internal energy of the molecules that make up a substance.
- Heat may be transferred from one substance to another.
We say that substances can absorb or radiate heat.
Temperature is a measure of the average kinetic energy of the molecules in a substance.
Energy, heat, and temperature are all related.

3. The Kinetic Molecular Theory
The Greek philosopher Democritus, in the 5th century B.C., described matter as empty space filled with many, many very small particles.
-He called these indivisible particles atoms.
-His theory meant that matter was discontinuous and not continuous as his contemporaries thought.
This early idea was dismissed by Aristotle and others. The science of chemistry, more than 2000 years later, provided the evidence for atoms.

4. Kinetic Molecular Theory and Molecules
Matter is made up of tiny units called atoms.
- Each element has a unique combination of subatomic particles (p+, N, e-) that make up its atoms.
- During any type of chemical or physical change
atoms are neither created or destroyed .
- Arrangements of atoms determines the type of matter.
Elements are pure substances made up of only one type of atom.
Compounds are made up of one type of atom, but have more complex structures.

5. Kinetic Molecular Theory and Molecules
Pure substances are combinations of 2 or more elements in definite proportions.
A molecule is the smallest particle of a compound in which all of the atoms maintain their identity.
They maintain the physical and chemical and properties of the compound.
Many atoms form monatomic molecules. (He)
Some atoms naturally form diatomic molecules: O2, H2 and Cl2 are examples.

6. Where these two gold crystals meet they are joined by a complex arrangement of atoms, forming a nanobridge that accommodates their different orientations. The gold atoms are 2.3 angstroms apart. It is possible to distinguish individual atoms and, at the edges of the two crystals, deduce their position in three dimensions. Credit: Image courtesy of DOE/Lawrence Berkeley National Laboratory

7. Interactions of Molecules
Some solids and liquids attract each other and cling to each other.
- Cohesion is when this attractive force is between like molecules. Example: Water clings to water.
- Adhesion is when one type of molecules is
attracted to another type of molecules.
Example: Water clings to skin.

8. Phases of Matter Solids Definite shape and definite volume.
Molecules are fixed distances apart and have strong cohesive forces.
Liquids
Close together.
Cohesive forces not as strong as in a solid.
Definite volume, but not a definite shape.

9. Phases of Matter Gases Cohesive forces are weak.
Kinetic energy is high.
Molecules move in random motion and are relatively far apart from each other.
They have no fixed shape or volume.
A vapor is a gas that is normally a liquid.
Example: Evaporated water!

10. Molecules Move
-All molecules have kinetic energy due to movements.
-This kinetic energy can be in the form of:
Vibrational energy.
Rotational energy.
Translational energy where the entire molecule has motion.

11. Temperature
The kinetic energy of a substance is measured as the temperature of that substance.
Temperature is actually a measure of the average kinetic energy and has nothing to do with heat until there is a transfer of energy.

12. Thermometers
We say a thermometer measures the hotness or coldness of an object.
What a thermometer really measures is the average kinetic energy of an object.
There is a physical transfer of kinetic energy to the thermometer which responds due to the increase in its kinetic energy.

13. Thermometer Scales
Fahrenheit
Boiling point of water is 212 OF and freezing point of water is 32 OF.
180 degrees between these two.
Celsius
Sets boiling point of water at 100 OC and freezing point of water at 0 OC.
100 divisions between these two points.

14. Kelvin or absolute scale
Begins at absolute zero; the temperature at which all kinetic energy is changed into potential energy, or all molecular motion ceases.
Boiling point of water is 373 K and freezing point of water is 273 K.
Divisions are same as the Celsius scale

15. The Fahrenheit, Celsius, and Kelvin scales

16. Conversions.
TF = 1.8 TC + 32 OC
TC = (TF - 32 OF)
1.8
1.8 accounts for the divisions between freezing point of water and boiling point of water.
There are 1.8 divisions in the F scale for every 1 division in the C scale.
TK = TC

17. Example:
The temperature of Lake Superior in August averages 34 OF. What is the temperature in OC.
Use: TC = (TF - 32 OF)
1.8
TC = (34 OF - 32 OF) / 1.8
TC = (2 OF) / 1.8
about 1.1 OC

18. Example:
What is the equivalent Celsius temperature of K? The equivalent Fahrenheit temperature?
Use: TK = TC + 273
Rearrange to : TC = TK
TC = K = OC
TF = 1.8 (127.0 OC) + 32 OC =

19. Heat Internal and External Energy
External energy is total potential and kinetic energy of everyday sized objects.
Internal energy is the total kinetic and potential energy of an object’s molecules.

20. External energy is the kinetic and potential energy that you can see
External energy is the kinetic and potential energy that you can see. Internal energy is the total kinetic and potential energy of molecules. When you push an object across a surface, you do work against friction. Some of the external mechanical energy goes into internal kinetic and potential energy, and the surfaces in contact become warmer.

21. Heat as Energy Transfer
Temperature is a measure of the average kinetic energy of an object.
Heat is a measure of the internal energy that has been absorbed or transferred from one body to another.
Increasing the internal energy is called heating.
Decreasing the internal energy is called cooling.

22. Two ways to increase temperature:
From a temperature difference, with energy moving from a region of higher temperature to a region of lower temperature.
From an object gaining energy by way of an energy form conversion (like when some KE becomes heat energy due to friction.)

23. Heat and temperature are different concepts
Heat and temperature are different concepts. A liter of water cup and a of water are both at the same temperature. The liter of water contains more heat since it will require more ice cubes to cool it than will be required for the cup of water.

24. Measures of Heat
The metric unit of measuring work, energy, or heat is the joule.
The metric unit of heat is also the calorie.
A calorie is the amount of energy needed to increase the temperature of 1 gram of water 1 OC (from 14.5 OC to 15.5 OC.
A kilocalorie is the amount of energy needed to increase the temperature of 1 kg of water 1 OC.

25. The English unit of heating is the BTU.
A BTU is the amount of energy needed to increase the temperature of 1 lb of water 1 OF.
A Quad is 1 quadrillion BTU 1 X 1015 BTU.
778 ftlb = 1 BTU
4.184 ftlb = 1 calorie
4,184 J = 1 kcalorie

26. Example: a 2,200. 0 kg automobile is moving at 90. 0 km/hr (25. 0 m/s)
Example: a 2,200.0 kg automobile is moving at 90.0 km/hr (25.0 m/s). How many kilocalories are generated when the car brakes to a stop?
KE = 1/2mv2
KE = 1/2(2,200 kg)(25.0m/s)2
KE = (1,100 kg)(625.0 m2/s2)
KE = 687,500 m2/s2
KE = 687,500 J
Kcal = 687,500 J X 1 kcal/4,184J = 164 kcal

27. Specific Heat
Three variables that influence energy transfer.
The temperature change.
The mass of the substance.
The nature of the material being heated.
The amount of heat (Q) needed to increase the temperature (Ti) of a pot of water from the initial temperature to a final temperature (Tf) is proportional to (Tf-Ti).
Q  (Tf-Ti).
Q  T.

28. The quantity of heat (Q) absorbed or given off during a certain change in temperature is also proportional to the mass (m) of the substance.
Q  m
Putting this all together we get:
Q  mcT
c is the specific heat of the substance.
Specific heat is the energy needed to increase the temperature of 1 gram of a substance 1 OC.

29. When two materials of different temperatures are involved in heat transfer and are perfectly insulated from the surroundings, the heat lost by one is equal to the heat gained by the other.
heat lost = heat gained.
Qlost = Qgained
(mcT)lost = (mcT)gained

30. Example: How much heat must be supplied to a 500
Example: How much heat must be supplied to a g pan to increase its temperature from 20.0 OC to OC if the pan is made of a) iron and b) aluminum.
Iron from table 4.2 has a specific heat of 0.11 cal/gOC.
Q = mcT
Q = (500.0g)(0.11 cal/gOC)(80.0OC)
Q = 4,400 cal or 4.40 kcalories
Aluminum from table 4.2 has a specific heat of 0.22 cal/gOC
Q = (500.0g)(0.22 cal/gOC)(80.0OC)
8,800 calories or 8.80 kcalories

31. Heat Flow
Conduction
When there is a temperature difference; there is a natural tendency for temperature to flow from the area of high temp to the area of low temp.
Conduction is the transfer of energy from molecule to molecule.
The rate depends on the temperature difference, the area and thickness of the substance, and the nature of the material.

32. Some materials are good conductors while others are good insulators.
Conductors transfer energy very efficiently.
Insulators transfer energy very inefficiently.
The best conductors are usually metals which have very little air space between molecules.
The best insulators have a great deal of air space between molecules.
The vacuum is absolutely the best insulator as it has no molecules to pass on energy.

33. Convection
Transfer of heat by a large scale displacement of molecules with a relatively higher kinetic energy.
Molecules with higher kinetic energy are moved from one place to another place.
Happens only in liquids and gases where fluid motion can carry molecules with higher kinetic energy over a distance.

34. Radiation
Radiation involves the form of energy called radiant energy that moves through space.
All objects with a temperature above absolute zero give off radiant energy.
The absolute temperature of the object determines the rate, intensity, and kinds of radiant energy emitted.

35. Energy, Heat, and Molecular Theory

36. Phase Change
The motion of a molecule can be increased by:
Adding heat through a temperature difference.
The absorption of one of the five forms of energy.
Temperature increases according to the specific heat of the substance.

37. When a substance changes from one state to another, the transition is called a phase change.
A phase change always absorbs or releases energy, a quantity of heat that is not associated with a temperature change.
Latent heat is the hidden energy of a phase change, which is energy that goes in or comes out of internal potential energy.

38. Three major types of phase change.
Solid-liquid.
Liquid-gas.
Solid-gas

39. Solid-liquid
The temperature at which a substance changes from a liquid to a solid is called the freezing point.
The temperature at which a solid changes to a liquid is the melting point.
Both of these occur at the same temperature.

40. Liquid-gas
The temperature at which a liquid changes from the liquid phase to the gaseous phase is the boiling point.
The temperature at which a gas or vapor changes to the liquid phase is the condensation point.
Both of these occur at the same temperature.

41. Solid-gas
A phase change directly from a solid to a gas or vapor is called sublimation.

42. Latent heat of fusion
The latent heat of fusion is the heat involved in a solid-liquid phase change in melting or freezing.
A melting solid absorbs energy and a freezing liquid releases this same amount of energy, warming the surroundings.
The total heat involved in a solid-liquid phase change depends on the mass of the substance involved.
Q = mLf
Where Lf is the latent heat of fusion for the substance involved.

43. Latent heat of vaporization
The amount of heat involved during a phase change from a liquid to a gas or vapor is called the latent heat of vaporization.
The latent heat of vaporization is the heat involved in a liquid-gas phase change where there is evaporation or condensation.
The escaping molecules absorb energy from the surroundings, and a condensing gas releases this exact same amount of energy.

44. The total heating depends on the amount of water vapor that condenses so that:
Q = mLV
Where LV is the latent heat of vaporization for the substance involved.

45. Example:
How much energy does a refrigerator remove from g of water at 20.0 OC to make ice at OC
Three steps.
Q1 = mcT to cool from 20.0 OC to 0.0 OC
Q1 = (100.0g)(1.00cal/gOC)(0.0OC-20.0OC)
= 2,000 cal = 2.00 X 10 3 cal.

46. Q2 = mLf to remove latent heat of fusion.
(100.0g)(80.0cal/g)
8,000 cal = 8.00 X 103 cal
Q3 = mcT to go from 0.0 OC to -10 OC
(100.0g)(0.500cal/g)(10.0OC-0.0OC)
500 cal = 5.00 X 102 cal
Qtotal = Q1 + Q2 + Q3
= 2.00 X 103 cal X 103 cal X 102 cal
= 1.05 X 104 cal

47. Evaporation and Condensation
Evaporation occurs when enough energy is put into a system to cause liquid molecules to overcome attractive forces near the surface, escape, and become a gas or vapor.
In evaporation, more molecules are leaving the liquid state than are returning.
In condensation, more molecules are returning to the liquid state than are leaving.
When the condensation rate is equal to the evaporation rate, the air above the liquid is saturated (holds all the vapor that it is capable of holding).

48. Four ways to increase the rate of evaporation.
An increase in temperature of the liquid will increase the average kinetic energy of the molecules and thus increase the number of high energy molecules capable of escaping from the liquid state.
Increase the surface area of the liquid in contact with the air.

49. Removal of water vapor from near the surface will prevent the return of molecules to the liquid phase.
Reducing atmospheric pressure will reduce one of the forces holding molecules in a liquid.

50. Relative Humidity
The ratio of how much water vapor is in air to how much water vapor it could hold at a certain temperature is the relative humidity
Usually expressed as a percent.

51. The inside of a closed bottle is isolated so the space above the liquid becomes saturated. When it is saturated, the evaporation rate equals the condensation rate. When cooled, condensation exceeds evaporation and liquid droplets form inside the bottle.

52. Thermodynamics

53. Introduction
The laws of thermodynamics describe what happens to energy as it is transformed into work and to other forms.
Thermodynamics is concerned with internal energy, which is the total internal kinetic and potential energy of a system.

54. The system is the component we want to describe.
The state of the system are the variables under which it exists, temperature, pressure, volume, heat, etc…
Everything outside of the system is the surroundings.

55. The First Law of Thermodynamics
The energy supplied to a system is equal to the change in internal energy of the system.
The Second Law of Thermodynamics.
Heat flows from objects with a higher temperature to objects with a cooler temperature.

56. The Second Law and Natural Processes
Energy can be viewed from two considerations of scale:
The observable external energy of an object.
The internal energy of the molecules, or particles that make up an object.

57. Two kinds of motion that the particles of an object can have.
A coherent motion where they move together.
An incoherent, chaotic motion of individual particles.
Work on an object is associated with coherent motion, while heating an object is associated with its internal incoherent motion.

58. Entropy
Energy is always degrading toward a more disorderly state.
The total entropy of the universe is continually increasing.
The natural process is for the state of order to degrade into a state of disorder with a corresponding increase in entropy.

59. Eventually all of the useable energy in the universe will diminish to unusable forms.
The universe will at some time reach a limit of disorder called the heat death of the universe.
The heat death of the universe is the theoretical limit of disorder, with all molecules spread far, far apart, vibrating slowly with a uniform low temperature.