1. Unit 4 - Thermodynamics
Chapters 9 and 10

2. Chapter 9 - Heat Expectations:
Learn the difference between temperature and heat.
Learn how different substances change temperature or phase when energy is added to or removed from the substances.

3. Chapter 9 Preview Section 1: Temperature & Thermal Equilibrium
Defining Temperature
Measuring Temperature
Section 2: Defining Heat
Heat and Energy
Thermal Conduction
Heat and Work
Section 3: Changes in Temperature and Phase
Specific Heat Capacity
Latent Heat

4. Section 1 – Temperature & Thermal Equilibrium
Objectives:
Relate temperature to the kinetic energy of atoms and molecules.
Describe the changes in the temperatures of two objects reaching thermal equilibrium.
Identify the various temperature scales, and convert from one scale to another.

5. Defining Temperature
Energy must either be added to or removed from a substance in order to change its temperature.
Temperature is proportional to the average kinetic energy of particles in the substance.
The energies associated with atomic motion are referred to as internal energy, which is proportional to the substance’s temperature.
It is due to both the random motions of its particles and the potential energy that results from the distances and alignments between the particles.

6. Defining Temperature
For an ideal gas, the internal energy depends only on the temperature of the gas.
For nonideal gases, liquids, and solids, other properties contribute to the internal energy.
U = internal energy
ΔU = change in internal energy
Thermal equilibrium = the state in which two bodies in physical contact with each other have identical temperatures

7. Defining Temperature
Thermal equilibrium is the basis for measuring temperature with thermometers.
If the temperature of a substance increases, so does its volume = thermal expansion (ex. = bridges)
Different substances undergo different amounts of expansion for a given temperature change.
This is indicated by the coefficient of volume expansion, which is greatest in gases.

8. Measuring Temperature
In order for a device to be used as a thermometer, it must make use of a change in some physical property that corresponds to changing temperature – the volume of a gas or liquid, the pressure of a gas at constant volume.
Thermometers are calibrated using fixed temperatures.
One reference point marks when the thermometer is in thermal equilibrium with a mixture of water and ice at one atmosphere of pressure.

9. Measuring Temperature
This temperature is referred to as the ice point or melting point of water and is defined as 0ºC.
The second reference point marks when the thermometer is in thermal equilibrium with a mixture of steam and water at one atmosphere of pressure.
This temperature is referred to as the steam point or boiling point of water and is defined as 100ºC.

10. Measuring Temperature
The remainder of the scale is divided into equally spaced units that represent the degrees between 0 and 100.
The temperature scales most widely used today are the Fahrenheit, Celsius, and Kelvin scales.
The U.S. still uses the Fahrenheit scale, but the Celsius scale is used by countries that utilize the metric system and within the sciences.

11. Measuring Temperature
Celsius-Fahrenheit Temperature Conversion:
TF = TC
Celsius-Kelvin Temperature Conversion:
T = TC 9 5

12. Measuring Temperature
Scale
Ice Point
Steam Point
Fahrenheit
32ºF
212ºF
Celsius
0ºC
100ºC
Kelvin (absolute) K K
Practice A, p. 303

13. Section 2 – Defining Heat
Objectives:
Explain heat as the energy transferred between substances that are at different temperatures.
Relate heat and temperature change on the macroscopic level to particle motion on the microscopic level.
Apply the principle of energy conservation to calculate change in potential, kinetic, and internal energy.

14. Heat and Energy
To understand thermal processes we have to consider the behavior of atoms and molecules.
Mechanics explains what is happening at the molecular (microscopic) level and accounts for what we observe at the macroscopic level.
Heat = the energy transferred between objects because of a difference in their temperature
Energy transferred as heat tends to move from an object at higher temperature to an object at lower temperature.

15. Heat and Energy
The direction in which energy travels as heat can be explained at the atomic level.

16. Heat and Energy

17. Heat and Energy
Thermal equilibrium may be understood in terms of energy exchange between two objects at equal temperature.
At the same temperature, the net energy transferred between two objects is zero.

18. Heat and Energy So, the difference between temperature and heat:
The atoms of all objects are in continuous motion, so all objects have some internal energy
= Temperature is a measure of that energy
= All objects have some temperature

19. Heat and Energy
The energy transferred from one object to another because of the temperature difference between them is heat
= No temperature difference means no net energy transferred as heat

20. Heat and Energy Heat, like work, is energy in transit.
Heat units can be converted into joules, the SI unit for energy.

21. Thermal Conduction
Thermal conduction occurs when energy transfer increases the temperature of an object.
The rate of thermal conduction depends on the properties of the substance being heated.
Substances that rapidly transfer energy as heat = thermal conductors
Substances that slowly transfer energy as heat = thermal insulators

22. Thermal Conduction
Convection is the transfer of energy involving the movement of cold and hot matter.
It involves the combined effects of heat, pressure differences, conduction, and buoyancy.
Electromagnetic radiation occurs when objects reduce their internal energy by giving off radiation of particular wavelengths or when objects are heated by the radiation.

23. Heat and Work
When objects collide in-elastically, not all of their initial kinetic energy remains as kinetic energy after the collision.
Some of the energy is absorbed as internal energy by the objects.
If changes in internal energy are taken into account along with changes in mechanical energy, the total energy is a universally conserved property.

24. Heat and Work Conservation of Energy: ΔPE + ΔKE + ΔU = 0
Practice B, p 311

25. Section 3 – Changes in Temperature & Phase
Objectives:
Perform calculations with specific heat capacity.
Interpret the various sections of a heating curve.

26. Specific Heat Capacity
Specific heat capacity = the energy required to change the temperature of 1 kg of a substance by 1ºC
specific heat capacity =
cp =
energy transferred as heat
mass x change in temp Q
mΔT

27. Specific Heat Capacity
The p indicates that the specific heat capacity is measured at constant pressure.
The equation for specific heat capacity applies to both substances that absorb energy from their surroundings and those that transfer energy to their surroundings.

28. Specific Heat Capacity
To measure the specific heat capacity of a substance, it is necessary to measure mass, temperature change, and energy transferred as heat.
Qw = -Qx
Qw + Qx = 0
cp,wmwΔTw = -cp,xmxΔTx

29. Specific Heat Capacity
The w indicates water.
Calorimetry = experimental procedure used to measure the energy transferred from one substance to another as heat
Calorimeter = device that contains a thermometer for measuring temperature and a stirrer to ensure a uniform mixture of energy throughout water
Practice C, p 315

30. Latent Heat
When substances melt, freeze, boil, condense, or sublime, the energy added or removed changes the internal energy of the substance without changing the substance’s temperature.
These changes in matter are called phase changes.
Sublime = change from a solid to a vapor or from vapor to a solid

31. Latent Heat
Phase changes result from a change in the potential energy between particles of a substance.
When energy is added to or removed from a substance that is undergoing a phase change, the particles of the substance rearrange themselves to make up for their change of energy.
This rearrangement occurs without a change in the average kinetic energy of the particles.

32. Latent Heat
The energy that is added or removed per unit mass is called latent heat (L).
Q = mL
During melting, the energy that is added to a substance equals the difference between the total potential energies for particles in the solid and liquid phases = heat of fusion (Lf)

33. Latent Heat
During vaporization, the energy that is added to a substance equals the difference in the potential energy of attraction between the liquid particles and between the gas particles = heat of vaporization (Lv)

34. Chapter 10 - Thermodynamics
Expectations:
Learn how two types of energy transfer – work and heat – serve to change a system’s internal energy.
Learn a new form of the law of energy conservation and see how machine efficiency is limited.

35. Chapter 10 Preview Section 1: Relationships Between Heat & Work
Heat, Work, and Internal Energy
Thermodynamic Processes
Section 2: The First Law of Thermodynamics
Energy Conservation
Cyclic Processes
Section 3: Changes in Temperature & Phase
Efficiency of Heat Engines
Entropy

36. Section 1 – Relationships Between Heat & Work
Objectives:
Recognize that a system can absorb or release energy as heat in order for work to be done on or by the system and that work done on or by a system can result in the transfer of energy as heat.
Compute the amount of work done during a thermodynamic process.
Distinguish between isovolumetric, isothermal, and adiabatic thermodynamic processes.

37. Heat, Work, & Internal Energy
Work can increase the internal energy of a substance.
The internal energy can then decrease through the transfer of energy as heat.
Energy can be transferred to a substance as heat and then be used to do work.
On a microscopic scale, heat and work are similar – they refer to energy in transit.

38. Heat, Work, & Internal Energy
This energy exists in systems – sets of particles or interacting components considered to be a distinct physical entity for the purpose of study.
Systems are rarely isolated from their surroundings, so we have to account for interactions with the environment as well.
Environment = combination of conditions and influences outside a system that affect the behavior of the system

39. Heat, Work, & Internal Energy
In thermodynamic systems, work is defined in terms of pressure and volume change.
Work done by a gas:
W = PΔV
work = pressure x volume change

40. Heat, Work, & Internal Energy
If a gas expands:
ΔV is positive
Work done by the gas is positive
If a gas is compressed:
ΔV is negative
Work done by the gas is negative
When the gas volume remains constant, there is no displacement and no work is done on or by the system.

41. Heat, Work, & Internal Energy
Practice A, p 338

42. Thermodynamic Processes
Most processes transfer energy as both heat and work.
In many processes, one type of energy transfer is dominant and the other type is negligible.
In these instances, the real process can be approximated with an ideal process.

43. Thermodynamic Processes
In general, when a gas undergoes a change in temperature but no change in volume, no work is done on or by the system.
This type of process is called a constant-volume process, or isovolumetric process.

44. Thermodynamic Processes
Isothermal process = system’s temperature remains constant and internal energy doesn’t change when energy is transferred to or from the system as heat or work
These processes must
happen slowly.

45. Thermodynamic Processes
In an isothermal process, small amounts of energy are removed as work.
Energy is added as heat.
Thermal equilibrium is restored.

46. Thermodynamic Processes
Adiabatic process = process in which changes occur but no energy is transferred to or from a system as heat
The decrease in internal energy must be equal to the energy transferred from the gas as work.
This process must happen rapidly.
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47. Section 2 – The First Law of Thermodynamics
Objectives:
Illustrate how the first law of thermodynamics is a statement of energy conservation.
Calculate heat, work, and the change in internal energy by applying the first law of thermodynamics.
Apply the first law of thermodynamics to describe cyclic processes.

48. Energy Conservation

49. Energy Conservation
On a roller coaster that experiences friction, mechanical energy isn’t conserved.
A steady decrease in the car’s total mechanical energy occurs because of friction.
Mechanical energy is transferred to the atoms and molecules throughout the entire roller coaster (cars and track).
The roller coaster’s internal energy increases by an amount equal to the decrease in the mechanical energy.

50. Energy Conservation
Most of this energy is gradually dissipated to the air surrounding the roller coaster as heat.
The principle of energy conservation that takes into account a system’s internal energy as well as work and heat is called the first law of thermodynamics.
This principle is often applied to systems, so values for work and heat provide information.

51. Energy Conservation

52. Energy Conservation The total change in the internal energy is:
ΔU = Uf – Ui
Uf = final internal energy
Ui = initial internal energy

53. Energy Conservation The first law of thermodynamics: ΔU = Q – W
ΔU = change in internal energy
Q = energy transferred to or from a system as heat
W = energy transferred to or from system as work
Practice B, p 345

54. Cyclic Processes
In a cyclic process, the system’s properties at the end of the process are identical to the system’s properties before the process took place.
The final and initial values of internal energy are the same, and the change in internal energy is zero.
ΔUnet = 0 and Qnet = Wnet

55. Cyclic Processes
These processes resemble isothermal ones in that all energy is transferred as work and heat.
Heat engines use heat to do work as part of a cyclic process.
They do work by transferring energy from a high-temperature substance to a lower- temperature substance.

56. Cyclic Processes
For each complete cycle, the net work done will equal the difference between the energy transferred as heat from a high-temperature substance to the engine (Qh) and the energy transferred as heat from the engine to a lower-temperature substance (Qc).
Wnet = Qh - Qc

57. Cyclic Processes
The larger the difference between the energy transferred as heat into the engine and out of the engine, the more work the engine can do in each cycle.
No heat engine works perfectly – only part of the available internal energy leaves the engine as work.
Most of the energy is removed as heat.

58. Section 3 – The Second Law of Thermodynamics
Objectives:
Recognize why the second law of thermodynamics requires two bodies at different temperatures for work to be done.
Calculate the efficiency of a heat engine.
Relate the disorder of a system to its ability to do work or transfer energy as heat.

59. Efficiency of Heat Engines
It is impossible to construct a heat engine that, operating in a cycle, absorbs energy from a hot reservoir and does an equivalent amount of work.
This is the basis for the second law of thermodynamics: no cyclic process that converts heat entirely into work is possible.
This tells us that some energy must always be transferred as heat to the system’s surroundings (Qc > 0).

60. Efficiency of Heat Engines
Cyclic processes can be made to approach ideal situations.
A measure of how well an engine operates is given the engine’s efficiency (eff).
Efficiency = a measure of the useful energy taken out of a process relative to the total energy put into the process.

61. Efficiency of Heat Engines
Equation for the efficiency of a heat engine:
eff = = = 1 –
Wnet = net work done by engine
Qh = energy added as heat
Qc = energy removed as heat
Wnet Qh - Qc Qc Qh Qh Qh

62. Efficiency of Heat Engines
Efficiency is a unitless quantity that can be calculated using only the magnitudes for the engines added to and taken away from the engine.
The efficiencies of all engines are less than 1.0.
The smaller the fraction of usable energy that an engine can provide, the lower its efficiency is.
The efficiency equation gives only a maximum value for an engine’s efficiency.
Practice C, p. 355

63. Entropy
In thermodynamics, a system left to itself tends to go from a state with a very ordered set of energies to one in which there is less order.
Entropy = the measure of a system’s disorder
The greater the entropy of a system is, the greater, the system’s disorder.
The entropy of systems tends to increase; once a system has reached a state of greatest disorder, it will tend to remain in that state and have maximum entropy.

64. Entropy
The motion of the particles of a system is not well ordered and is less useful for doing work.

65. The entropy of the universe increases in all natural processes.
Because of the connection between a system’s entropy, its ability to do work, and the direction of energy transfer, the second law of thermodynamics can also be stated as:
The entropy of the universe increases in all natural processes.

66. Entropy
Entropy can decrease for parts of systems provided that the decrease is offset by a greater increase in entropy elsewhere in the universe.