PHY 352_Lecture 3 Heat&Thermodynamics II_REVISED.ppt

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University of Education-Winneba
University of Education-Winneba
Faculty of Science Education
Faculty of Science Education
Department of Physics Education
Department of Physics Education
Lecturer Contact
Lecturer Contact
Details
Details
Dr. Desmond Appiah
Dr. Desmond Appiah
PHY 352
PHY 352
First Law of Thermodynamics
First Law of Thermodynamics

2
Objectives
• Introduce the concept of energy and define its various forms.
• Discuss the nature of internal energy.
• Define the concept of heat and the terminology associated with energy
transfer by heat.
• Discuss the three mechanisms of heat transfer: conduction,
convection, and radiation.
• Define the concept of work, including electrical work and several forms
of mechanical work.
• Introduce the first law of thermodynamics, energy balances, and
mechanisms of energy transfer to or from a system.
• Determine that a fluid flowing across a control surface of a control
volume carries energy across the control surface in addition to any
energy transfer across the control surface that may be in the form of
heat and/or work.
• Define energy conversion efficiencies.
• Discuss the implications of energy conversion on the environment.

3
INTRODUCTION
• If we take the entire room—including the air and the refrigerator (or
fan)—as the system, which is an adiabatic closed system since the
room is well-sealed and well-insulated, the only energy interaction
involved is the electrical energy crossing the system boundary and
entering the room.
• As a result of the conversion of electric energy consumed by the
device to heat, the room temperature will rise.
A refrigerator
operating with its
door open in a well-
sealed and well-
insulated room
A fan running in a
well-sealed and
well-insulated room
will raise the
temperature of air in
the room.

4
FORMS OF ENERGY
• Energy can exist in numerous forms such as thermal, mechanical,
kinetic, potential, electric, magnetic, chemical, and nuclear, and their
sum constitutes the total energy, E of a system.
• Thermodynamics deals only with the change of the total energy.
• Macroscopic forms of energy: Those a system possesses as a whole
with respect to some outside reference frame, such as kinetic and
potential energies.
• Microscopic forms of energy: Those related to the molecular
structure of a system and the degree of the molecular activity.
• Internal energy, U: The sum of all the microscopic forms of energy.
The macroscopic energy of an
object changes with velocity and
elevation.
• Kinetic energy, KE: The energy
that a system possesses as a result
of its motion relative to some
reference frame.
• Potential energy, PE: The energy
that a system possesses as a result
of its elevation in a gravitational
field.

5
TOTAL ENERGY OF A SYSTEM
• Sum of all forms of energy (i.e., thermal, mechanical,
kinetic, potential, electrical, magnetic, chemical, and
nuclear) that can exist in a system
• For systems we typically deal with in this course, sum
of internal, kinetic, and potential energies
E = U + KE + PE
E = Total energy of system
U = internal energy
KE = kinetic energy
PE = potential energy

6
Total energy
of a system
Energy of a system
per unit mass
Potential energy
per unit mass
Kinetic energy
per unit mass
Potential energy
Total energy
per unit mass
Kinetic energy
Mass flow rate
Energy flow rate

7
Some Physical Insight to Internal Energy
The internal energy of a
system is the sum of all forms
of the microscopic energies.
The various forms of
microscopic
energies that make
up sensible energy.
Sensible energy: The portion
of the internal energy of a
system associated with the
kinetic energies of the
molecules.
Latent energy: The internal
energy associated with the
phase of a system.
Chemical energy: The internal
energy associated with the
atomic bonds in a molecule.
Nuclear energy: The
tremendous amount of energy
associated with the strong
bonds within the nucleus of the
atom itself.
Internal = Sensible + Latent + Chemical + Nuclear
Thermal = Sensible + Latent
We recall that, temperature is a
measure of the average K.E of
particles in a system

8
The macroscopic kinetic energy is an
organized form of energy and is much
more useful than the disorganized
microscopic kinetic energies of the
molecules.
• The total energy of a system, can
be contained or stored in a system,
and thus can be viewed as the
static forms of energy.
• The forms of energy not stored in a
system can be viewed as the
dynamic forms of energy or as
energy interactions.
• The dynamic forms of energy are
recognized at the system boundary
as they cross it, and they represent
the energy gained or lost by a
system during a process.
• The only two forms of energy
interactions associated with a
closed system are heat transfer
and work.
• The difference between heat transfer and work: An energy interaction is
heat transfer if its driving force is a temperature difference. Otherwise it is
work.

9
More on Nuclear Energy
The fission of uranium and the fusion of
hydrogen during nuclear reactions, and
the release of nuclear energy.
• The best known fission reaction
involves the split of the uranium atom
(the U-235 isotope) into other elements
and is commonly used to generate
electricity in nuclear power plants (440
of them in 2004, generating 363,000
MW worldwide), to power nuclear
submarines and aircraft carriers, and
even to power spacecraft as well as
building nuclear bombs.
• Nuclear energy by fusion is released
when two small nuclei combine into a
larger one.
• The uncontrolled fusion reaction was
achieved in the early 1950s, but all the
efforts since then to achieve controlled
fusion by massive lasers, powerful
magnetic fields, and electric currents to
generate power have failed.

10
Mechanical Energy
Mechanical energy: The form of energy that can be converted to
mechanical work completely and directly by an ideal mechanical device such
as an ideal turbine.
Kinetic and potential energies: The familiar forms of mechanical energy.
Mechanical energy of a
flowing fluid per unit mass
Rate of mechanical energy
of a flowing fluid
Mechanical energy change of a fluid during incompressible flow per unit mass
Rate of mechanical energy change of a fluid during incompressible flow

11
ENERGY TRANSFER BY HEAT
Energy can cross the
boundaries of a closed system
in the form of heat and work.
Temperature difference is the driving
force for heat transfer. The larger the
temperature difference, the higher is the
rate of heat transfer.
Heat: The form of energy that is
transferred between two
systems (or a system and its
surroundings) by virtue of a
temperature difference.
Note: Internal Energy (U) α
Temperature (T)
Substances DONOT contain heat but
contains internal energy so when the
particles move from higher temperature
to lower temperature is used to
determine heat
Joule's second law states that the
internal energy of an ideal gas is
independent of its volume and pressure,
depending only on its temperature.
So what is the Joule’s first law?
So what is the Joule’s first law?

12
Energy is
recognized as
heat transfer
only as it
crosses the
system
boundary.
During an adiabatic process, a system
exchanges no heat with its surroundings.
Heat transfer
per unit mass
Amount of heat transfer
when heat transfer rate
changes with time
Amount of heat transfer
when heat transfer rate
is constant

13
Historical Background on Heat
• Kinetic theory: Treats molecules
as tiny balls that are in motion and
thus possess kinetic energy.
• Heat: The energy associated with
the random motion of atoms and
molecules.
Heat transfer mechanisms:
• Conduction: The transfer of energy
from the more energetic particles of
a substance to the adjacent less
energetic ones as a result of
interaction between particles.
• Convection: The transfer of energy
between a solid surface and the
adjacent fluid that is in motion, and
it involves the combined effects of
conduction and fluid motion.
• Radiation: The transfer of energy
due to the emission of
electromagnetic waves (or
photons).
In the early nineteenth century, heat was
thought to be an invisible fluid called the
caloric that flowed from warmer bodies to
the cooler ones.

14
ENERGY TRANSFER BY WORK
• Work: The energy transfer associated with a force acting through a distance.
 A rising piston, a rotating shaft, and an electric wire crossing the
system boundaries are all associated with work interactions
• Formal sign convention: Heat transfer to a system and work done by a
system are positive; heat transfer from a system and work done on a system
are negative.
• Alternative to sign convention is to use the subscripts in and out to indicate
direction. This is the primary approach in this text.
Specifying the directions
of heat and work.
Work done per
unit mass
Power is the
work done per
unit time (kW)

15
Heat vs. Work
• Both are recognized at the boundaries of
a system as they cross the boundaries.
That is, both heat and work are
boundary phenomena.
• Systems possess energy, but not heat
or work.
• Both are associated with a process, not
a state.
• Unlike properties, heat or work has no
meaning at a state.
• Both are path functions (i.e., their
magnitudes depend on the path followed
during a process as well as the end
states). Properties are point functions; but
heat and work are path functions
(their magnitudes depend on the
path followed).
Properties are point functions
have exact differentials (d ).
Path functions
have inexact
differentials ( )

16
THE FIRST LAW OF THERMODYNAMICS
• The first law of thermodynamics (the conservation of energy principle)
provides a sound basis for studying the relationships among the various forms of
energy and energy interactions.
• The first law states that energy can be neither created nor destroyed during a
process; it can only change forms.
• The First Law: For all adiabatic processes between two specified states of a
closed system, the net work done is the same regardless of the nature of the
closed system and the details of the process.
Energy cannot be
created or
destroyed; it can
only change
forms.
The increase in the energy of a
potato in an oven is equal to the
amount of heat transferred to it.

17
THE FIRST LAW OF THERMODYNAMICS
ΔU = Q – W ,
where ΔU is the increase of internal energy of the
system, Q is the heat entering or leaving the system
heat entering or leaving the system,
and W is the work done by the system
work done by the system.
This means that the internal energy of a system is equal
to the heat energy transferred to or from the system
minus the work done to or from the system
Q
W
ΔU
All of these quantities will be measured in Joules

19
EXAMPLES
If 100 J of compression work is done on a system and as a result the internal
energy increases by 74J. How much of the energy is transferred as heat and
which direction?
ΔU = Q – W
However, compression work done = positive
ΔU = Q +W
Q = (+74 J)-(+100 J)
Q = -26 J
Thus, the heat energy leaves the system (direction) or 26 J of the total work done is
lost in the process by the system as heat
A sample of gas does 150 J of work against the surroundings and loses 90J
of internal energy in the process. Does the gas gain or lose heat and how
much?
ΔU = Q + W
However, work done BY the system against the surrounding = (-)
ΔU = - 90 J
Q = (-90 J)+(150J)
Q = 60 J
Thus, 60 J of heat energy is absorbed for the process to occur

20
Work done by closed system
In order for work to be accomplished by the gas, it must expand or
contract (change volume)
A change in pressure, only, will not result in any work being
accomplished
In most examples a piston or object atop the gas must be moved for work
to be accomplished
Since we are considering the work done by the gas, as the piston moves,
the gas loses energy

21
Work done by closed system
•If system volume expands against a force, work is done by the system.
•If system volume contracts under a force, work is done on the system.
At the equilibrium state

22
Thermodynamic Processes
A thermodynamic process is represented by a change in
one or more of the thermodynamic variables describing
the system.
Each point on the curve
represents an equilibrium state
of the system.
Our equation of state, the ideal
gas law (PV = nRT), only
describes the system when it is
in a state of thermal equilibrium.

23
Reversible and Irreversible
Processes
A process is reversible if it does not violate any law of
physics when it is run backwards in time.
For example an ice cube placed on a countertop in a warm
room will melt.
The reverse process cannot occur: an ice cube will not form
out of the puddle of water on the countertop in a warm room.

24
Reversible Thermodynamic Process
For a process to be reversible each point on the curve
must represent an equilibrium state of the system.
The ideal gas law
(PV = nRT), does
not describe the
system when it is
not in a state of
thermal
equilibrium.
Reversible Process
Irreversible Process

25
A PV diagram can be used to represent the state changes of a
system, provided the system is always near equilibrium.
The area under a PV curve
gives the magnitude of the
work done on a system. W>0
for compression and W<0 for
expansion.
Thermodynamic Processes
Prove!

PV Diagram basics
P(Pa)
V(m3
)

PV Diagram basics
P(Pa)
V(m3
)
Given a point on the
diagram we can use
PV = nRT
to find the gas’s
temperature (K)

PV Diagram basics
P(Pa)
V(m3
)
If we see a line or curve
connecting points then
we know the gas has
changed its properties in
some way

PV Diagram basics
P(Pa)
V(m3
)
When we see an arrow
on that line, then we
know the original and
final states of the gas

PV Diagram basics
P(Pa)
V(m3
)
Movement to the right
shows expansion
(remember this is
negative work)

PV Diagram basics
P(Pa)
V(m3
)
Movement to the left
shows compression
(remember this is
positive work)

PV Diagram basics
P(Pa)
V(m3
)
Movement directly
up or down shows
no change in
volume (remember
this is zero work)

PV Diagram basics
P(Pa)
V(m3
)
When the path
closes then we
have one complete
cycle
The gas has
returned to its
original pressure,
temperature, and
volume

34
• Isobaric (Isopiestic) Process
A thermodynamic process that occurs while the pressure remains
constant
From Greek isos "equal“ and baros "weight"
Here, no pressure can be done by or on the system
The work done is the area under the
process
Note: Take into consideration the arrow
of the process
Practical Example: Boiling of water to steam or freezing of water to ice. In
the process, a gas either expands or contracts to maintain constant
pressure and hence the net amount f work is done by or on the system

35
Example
A gas expands at a constant pressure of 5 atm from 4L to 12L. How much work is
done by the gas in Joules?

36
• Isochoric (Isolvolumetric) Process
A thermodynamic process that occurs while the volume remains constant.
Look for gases that are contained in a closed or fixed container
From Greek isos "equal“ and choro "place“
For a constant volume, W = 0
Thus, no work is done on or
by the system
ΔU = Q – W ΔU = Q
An example is a bomb calorimeter
Combustion occurs in a rigid
container such that only heat
transfer occurs

37
Example
The pressure of a gas inside a rigid 5L container changes from 4 atm to 2 atm. How
much work is performed during this process?

38
The work done on a system depends on the path taken in
the PV diagram. The work done on a system during a
closed cycle can be nonzero.
To go from the state (Vi, Pi) by the path (a) to the state (Vf,
Pf) requires a different amount of work then by path (b). To
return to the initial point (1) requires the work to be nonzero.

39
• Isothermal Process
A thermodynamic process that occurs while the temperature remains
constant. Usually a relatively slow process to allow the gas to maintain
its temperature.
From Greek isos "equal“ and therme "heat"
ΔU = Q – W ΔU = 0
For a constant temperature, T= 0
Thus, there cannot be any change in
the internal energy since it is
proportional to temperature ΔU = 0.
This implies that any heat supplied to
the system is completely used to do
work
Q = W
An ideal example is a car engine,
as the piston ought to convert all
the heat energy from the
combustion reaction to expansion
work that moves the car.
For the isothermal graphs
recall Boyle’s law

40
For a constant temperature (isothermal) process, U = 0 and the
work done on an ideal gas is
.
ln
i
f









V
V
nRT
W Prove!
If the gas was compressed so that dl pointed into the gas, the volume
would decrease and dV < 0. The work done by the gas in this case
would then be negative, which is equivalent to saying that positive
work was done on the gas, not by it.

41
Example
5 moles of N2 gas expands at a constant temperature of 400K
form 3L to 6 atm to 9L at 2 atm. How much work is performed
by the gas?
Answer : 18269 J
Calculate the Q, W and U for the reversible isothermal
expansion of one mole of an ideal gas at 37ºC from a volume
of 20dm3 to a volume of 30dm3
Answers : U = 0, W = -1045.02J Q = 1045.02J

Molar Specific Heat Capacities of Gases

43
Molar Specific Heat Capacities of Gases

44
• Adiabatic Process
A thermodynamic process that occurs without gain or loss of heat and without a
change in entropy (No heat transfer). Look for gases that are insulated from the
environment
From Greek a “not“, dia “through”, and batos “passable”
ΔU = Q – W Q = 0
This means that the internal energy
of a system changes as a result of
doing work on its surrounding or the
surrounding doing work on the
system
ΔU = -W
An example is as masses of air
change position as result of
pressure difference in the Earth
atmosphere

45
• Adiabatic Process
 Physical an adiabatic process can occur through two ways
1. Either the process takes pace in an insulated system or container
2. Either the process is performed quickly

46
• Adiabatic Equations of an Ideal Gas
 Let us consider one mole of an ideal gas having a volume, V
at pressure, P and thermodynamic temperature, T
measured in Kelvin.
 Suppose it undergoes a small adiabatic expansion, it does
the necessary external work at the cost if its own internal
energy which decreases hence its temperature falls.
 Let dV be the infinitesimally small change in the volume of
the gas at pressure, P. Then, the external work done by the
gas in its expansion will be
dW = -pdV

47
• Adiabatic Equations of an Ideal Gas

48
Summary of Thermodynamic
Processes

49
Summary of Thermal Processes
f i
W = -P(V -V )
W
Q
U 


The First Law of Thermodynamics
 
 
 
 
i
f
V
W = nRT ln
V
 
 
 
 
i
f
V
+ nRT ln
V
f i
3
+ nR ( T - T )
2

50
ENERGY CONVERSION EFFICIENCIES
Efficiency is one of the most frequently used terms in thermodynamics, and it
indicates how well an energy conversion or transfer process is accomplished.
Efficiency of a water
heater: The ratio of the
energy delivered to the
house by hot water to
the energy supplied to
the water heater.
The definition of
performance is not limited to
thermodynamics only.

51
Heating value of the fuel: The amount of heat released when a unit
amount of fuel at room temperature is completely burned and the
combustion products are cooled to the room temperature.
Lower heating value (LHV): When the water leaves as a vapor.
Higher heating value (HHV): When the water in the combustion gases is
completely condensed and thus the heat of vaporization is also
recovered.
The definition of the heating value of
gasoline.
The efficiency of space heating
systems of residential and
commercial buildings is usually
expressed in terms of the annual
fuel utilization efficiency
(AFUE), which accounts for the
combustion efficiency as well as
other losses such as heat losses
to unheated areas and start-up
and cooldown losses.

52
• Generator: A device that converts mechanical energy to electrical
energy.
• Generator efficiency: The ratio of the electrical power output to the
mechanical power input.
• Thermal efficiency of a power plant: The ratio of the net electrical
power output to the rate of fuel energy input.
A 15-W
compact
fluorescent
lamp provides
as much light as
a 60-W
incandescent
lamp.
Lighting efficacy:
The amount of light
output in lumens
per W of electricity
consumed.
Overall efficiency
of a power plant

53
The efficiency of a cooking
appliance represents the
fraction of the energy
supplied to the appliance that
is transferred to the food.
• Using energy-efficient appliances conserve
energy.
• It helps the environment by reducing the
amount of pollutants emitted to the
atmosphere during the combustion of fuel.
• The combustion of fuel produces
• carbon dioxide, causes global warming
• nitrogen oxides and hydrocarbons,
cause smog
• carbon monoxide, toxic
• sulfur dioxide, causes acid rain.

54
Efficiencies of Mechanical and Electrical Devices
The mechanical
efficiency of a fan is the
ratio of the kinetic
energy of air at the fan
exit to the mechanical
power input.
The effectiveness of the conversion process between
the mechanical work supplied or extracted and the
mechanical energy of the fluid is expressed by the
pump efficiency and turbine efficiency,
Mechanical efficiency

55
Generator
efficiency
Pump-Motor
overall efficiency
Turbine-Generator
overall efficiency
The overall efficiency of a turbine–
generator is the product of the
efficiency of the turbine and the
efficiency of the generator, and
represents the fraction of the
mechanical energy of the fluid
converted to electric energy.
Pump
efficiency

56
ENERGY AND ENVIRONMENT
• The conversion of energy from one form to another often affects the environment
and the air we breathe in many ways, and thus the study of energy is not
complete without considering its impact on the environment.
• Pollutants emitted during the combustion of fossil fuels are responsible for smog,
acid rain, and global warming.
• The environmental pollution has reached such high levels that it became a
serious threat to vegetation, wild life, and human health.
Energy conversion processes are often
accompanied by environmental pollution.
Motor vehicles are the largest source of air
pollution.

57
Ozone and Smog
• Smog: Made up mostly of ground-level ozone (O3), but it also contains numerous other
chemicals, including carbon monoxide (CO), particulate matter such as soot and dust,
volatile organic compounds (VOCs) such as benzene, butane, and other hydrocarbons.
• Hydrocarbons and nitrogen oxides react in the presence of sunlight on hot calm days to
form ground-level ozone.
• Ozone irritates eyes and damages the air sacs in the lungs where oxygen and carbon
dioxide are exchanged, causing eventual hardening of this soft and spongy tissue.
• It also causes shortness of breath, wheezing, fatigue, headaches, and nausea, and
aggravates respiratory problems such as asthma.
Ground-level ozone, which is the primary component
of smog, forms when HC and NOx react in the
presence of sunlight in hot calm days.
• The other serious pollutant in smog is carbon monoxide,
which is a colorless, odorless, poisonous gas.
• It is mostly emitted by motor vehicles.
• It deprives the body’s organs from getting enough oxygen
by binding with the red blood cells that would otherwise
carry oxygen. It is fatal at high levels.
• Suspended particulate matter such as dust and soot are
emitted by vehicles and industrial facilities. Such particles
irritate the eyes and the lungs.

58
Acid Rain
• The sulfur in the fuel reacts with oxygen to form sulfur dioxide (SO2), which is an
air pollutant.
• The main source of SO2 is the electric power plants that burn high-sulfur coal.
• Motor vehicles also contribute to SO2 emissions since gasoline and diesel fuel
also contain small amounts of sulfur.
Sulfuric acid and nitric acid are formed
when sulfur oxides and nitric oxides react with
water vapor and other chemicals high in the
atmosphere in the presence of sunlight.
• The sulfur oxides and nitric oxides react
with water vapor and other chemicals high
in the atmosphere in the presence of
sunlight to form sulfuric and nitric acids.
• The acids formed usually dissolve in the
suspended water droplets in clouds or fog.
• These acid-laden droplets, which can be
as acidic as lemon juice, are washed from
the air on to the soil by rain or snow. This
is known as acid rain.

59
The Greenhouse
Effect: Global
Warming
• Greenhouse effect: Glass allows the solar
radiation to enter freely but blocks the
infrared radiation emitted by the interior
surfaces. This causes a rise in the interior
temperature as a result of the thermal
energy buildup in a space (i.e., car).
• The surface of the earth, which warms up
during the day as a result of the absorption
of solar energy, cools down at night by
radiating part of its energy into deep space
as infrared radiation.
• Carbon dioxide (CO2), water vapor, and
trace amounts of some other gases such as
methane and nitrogen oxides act like a
blanket and keep the earth warm at night by
blocking the heat radiated from the earth.
The result is global warming.
• These gases are called “greenhouse
gases,” with CO2 being the primary
component.
• CO2 is produced by the burning of fossil
fuels such as coal, oil, and natural gas.
The greenhouse effect on earth.

60
• A 1995 report: The earth has already warmed about 0.5°C during the last
century, and they estimate that the earth’s temperature will rise another 2°C by
the year 2100.
• A rise of this magnitude can cause severe changes in weather patterns with
storms and heavy rains and flooding at some parts and drought in others, major
floods due to the melting of ice at the poles, loss of wetlands and coastal areas
due to rising sea levels, and other negative results.
• Improved energy efficiency, energy conservation, and using renewable
energy sources help minimize global warming.
The average car produces several times its
weight in CO2 every year (it is driven 20,000
km a year, consumes 2300 liters of gasoline,
and produces 2.5 kg of CO2 per liter).
Renewable energies such as wind are
called “green energy” since they emit no
pollutants or greenhouse gases.

61
Summary
• Forms of energy
 Macroscopic = kinetic + potential
 Microscopic = Internal energy (sensible + latent + chemical + nuclear)
• Energy transfer by heat
• Energy transfer by work
• Mechanical forms of work
• The first law of thermodynamics
 Energy balance
 Energy change of a system
 Mechanisms of energy transfer (heat, work, mass flow)
• Energy conversion efficiencies
 Efficiencies of mechanical and electrical devices (turbines, pumps)
• Energy and environment
 Ozone and smog
 Acid rain
 The Greenhouse effect: Global warming

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