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KENYATTA UNIVERSITY
SCHOOL OF PURE AND APPLIED SCIENCES
DEPARTMENT OF CHEMISTRY
SCH 201 CHEMICAL THERMODYNAMICS
LEVEL : II SEMESTER: I ACADEMIC YEAR: 2019/2020
Series -1 (Introduction and First Law of Thermodynamics)
Prof Onindo Charles, Mr. Gitari Mugambi, Dr. Eric Masika
Lecture Days: Wednesday's 5:00 – 7:00
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Course Purpose
Purpose of the course
To provide a basic introduction to chemical thermodynamics BEd
(Sc), BSc, BSc. (Micro) BSc. Analytical with Management and for
BSc. Industrial with Management students with main focus on
interrelationships of heat and work with in both physical and
chemical reactions. Thus, several basic principles of chemical
thermodynamics will be considered with enhancement of
computational skills.
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Course Objective & Content
Course Content:
Chemical thermodynamic systems, states, state function and equilibrium
state. First law of thermodynamics, constant volume, constant pressure and
the reversible process. Isothermal and adiabatic expansion of an ideal gas.
Thermochemistry: Heat changes involved in chemical reactions and Hess’s
law. Heat capacities and enthalpy dependence on temperature.
Second law of thermodynamics: entropy and disorder. Temperature
dependence on entropy and the third law of thermodynamics.
Free energy (G), pressure dependence of G on ideal gases, relationship
between G and the equilibrium constant.
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S/No.
1
Week(s)
Topic(s)
1 -4
Introduction
Background


Subtopic(s)
The course examines the phase rule (variables required to
define state of a system
Heat and work.
It also reviews the first law of thermodynamics, characteristics
of a function of state, enthalpy
, Ideal gas calculations, isothermal and adiabatic changes are
examined.
Tutorial Questions: Set II
CAT 1 (16th October 2019)
and 



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5 -7
Thermochemistry and 
Kirchoff’s equations.



Thermochemistry and Hess’s law. Derivation of heat capacities
(Cv and Cp) from the Kinetic molecular theory of gases.
Kirchhoff’s equations and applications.
Tutorial Questions: Set II
CAT 2 (6th November 2019)
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8-9
Second Law, Entropy 
and third law

Brief introduction to the second law of thermodynamics.
Entropy and disorder. Temperature dependence on entropy and
the third law of thermodynamics.
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10 -11
Gibbs Free Energy

The Gibb’s free energy and equilibrium (free energy changes for
chemical reactions) are given.
Pressure dependence of G on ideal gases, relationship between
G and the equilibrium constant.
The Clausius-Clapeyron Equation


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12 -14
Revision/Exams
Examinations (Nov/Dec 2020)
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Methodology and Assessment
Teaching organization:
Lecturing, class discussion, tutorials, online work and
excursions.
Instructional Materials
White Board and white board marker, projector for power point
presentation, manuals, journals, e-materials, text books,
handouts and internet.
Mode of delivery: Face-to- face (full time, part time)
Assessment:
Mode of assessment
Marks
2 Continuous Assessment Tests
20 %
3 – 6 Pre-selected Practicals
10 %
End of Semester examination
70 %
Total
100 %
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Expected Learning Outcomes
At the end of this course, learners will be able to:
1. Define (using illustrations as necessary) key working terms: e.g.
system (Open, closed and isolated systems), surroundings,
universe; Energy, heat and work; Isothermal, adiabatic, reversible
and irreversible processes or changes; Thermodynamic functions
such as internal energy (E) and enthalpy (H)
2. Illustrate the relationship between heat capacity at constant volume
(Cv), and Heat capacity at constant pressure (Cp).
3. Solve problems relating to chemical thermodynamic
thermodynamics, for example: The enthalpy change, during
expansion or compression of an ideal gas, of Enthalpy, internal
energy and the work done. Include adiabatic changes in systems.
4. Formulate the first law of thermodynamics for a closed systems
and arrange the change in energy in the closed systems via heat and
work transfer
5. Apply first law of thermodynamics for closed systems and
construct conservation of mass and energy equations.
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Expected Learning Outcomes (Cont-1):
6. State the Second Law of Thermodynamics and use it to predict
the spontaneity of physical and chemical changes.
7. calculate the standard entropy change (ΔSo ) for a physical or
chemical process given standard entropy values, So , for
reactants and products
8. State the Third Law of Thermodynamics and describe its
significance
9. Describe Helmholtz and Gibb’s Free Energies.
10. Apply principle in determining the Gibb’s Free Energy changes
for different physical and chemical transformations.
11. Relate changes in Gibb’s Free Energy with Temperature,
Enthalpy, Entropy and Equilibrium Constant.
12. Apply the Clausius-Clapeyron Equation in determining the
vapour pressure.
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Suggested References
References/instructional materials/Equipment:
CORE REFERENCES
1. Atkins, P.W., Physical Chemistry, any edition, OUP
2. Smith, E.B., Basic Chemical Thermodynamics, any edition, ELBS-OUP
3. Warn, J.R.W, Concise chemical thermodynamics, VNR, London
4. Laidler, K. J., John H. M. and Bryan C. S. (2003), Physical Chemistry with
student CD (4th Ed.) Houghton Mifflin.
5. kusoma.ku.ac.ke LMS SCH 201 online module
RECOMMENDED REFERENCE MATERIALS
1. H. D. B. Jenkins, Chemical Thermodynamics at a Glance, 1st ed. (Blackwell
2008).
2. I. M. Klotz and R. M. Rosenberg, Chemical Thermodynamics: Basic
Concepts and Methods, 7th ed. (Wiley 2008).
3. D. Kondepudi and I. Prigogine, Modern Thermodynamics, 1st ed. (Wiley
1998).
4. K. J. Laidler, J. H. Meiser, and B. C. Sanctuary, Physical Chemistry, ebooks.
(Houghton Mifflin 2010).
5. Website based materials
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1.0 INTRODUCTION
The term THERMODYMAMICS means flow of heat. The study
of energy and its transformations.
(Greek: therm, ‘heat’ dynamis ‘power’). In general it deals with
the inter conversion of various kinds of energy in physical and
chemical systems.
Thermodynamics
1. Predicts the feasibility of a physical process or chemical
reaction under given condition of temperature and pressure
2. Predicts whether a chemical reaction would occur
spontaneously or not under a given set of conditions
3. Helps to determine the extent to which a reaction would take
place.
.
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Introduction ….
Limitations of Thermodynamics
1. It predicts the extent to which a reaction can take place.
However it does not say anything about the rate.
2. It applies only to matter in bulk and not to individual atoms or
molecules
Importance of thermodynamics
Thermodynamics is important not only to chemistry but to other
areas of science and engineering as well. It touches our daily lives
as we use energy for manufacturing, travel and communications.
.
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Pictorial Application Areas of Thermodynamics
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1.1 THE NATURE OF ENERGY
Definitions
FORCE
A force is any kind of push or pull exerted on an object. For
example force of gravity ‘pull’ us to the ground. In chemistry,
positively charged nuclei exert a ‘pull’ on negatively charged
electrons.
WORK
When we increase the distance between a proton and an
electron, we perform work to overcome the attractive force
between them. The work, w, that we do in moving objects
against a force equals the product of the force, F, and the
distance,d, that the object is moved. w=F x d
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1.1 THE NATURE OF ENERGY
Definitions
ENERGY
Energy in the form of work, must be used to move an object
against a force. As we perform work, heat may be generated.
HEAT
Heat is the energy that is transferred from one object to another
because of difference in temperature.
WORK and HEAT are the two ways that we experience energy
changes in our MACROSCOPIC environment. Energy is the
capacity to do work or to transfer heat
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1.1 THE CONCEPT OF ENERGY
ENERGY may take various forms: kinetic energy (due to motion)
 potential energy (due to position)
 heat energy (as measured by temperature)
 chemical energy, etc
A basic definition given by the great physicist, Albert Einstein is:
E = mc2
where
E = energy, m = mass and c = velocity of light.
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Some SI derived Units
Physical Quantity Unit
Symbol Definition of unit
force
N
newton
pressure
1N = 1mkgs
-2
-2
-1 -2
pascal
Pa
Celcius temperature degree Celcius ˚C
1Pa = 1Nm = 1kgm s
t/˚C = T/K - 273.15
energy
joule
J
1J = 1 NM = 1m kgs
power
watt
W
1W =1Js
frequency
electric charge
hertz
coulomb
Hz
C
1Hz = 1s
1C = 1As
electric potential
volt
V
1V = 1JC = 1m kgs A
electric resistance
ohm
Ω
1Ω = 1VA =1m kgs A
-1
2
-2
2
-3
2
-3 -1
2
-3 -2
= 1m kgs
-1
-1
-1
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1.2 BASIC DEFINITIONS
System and surroundings:
• Any part of the universe which is selected for
thermodynamic study is called a system and the rest
of the Universe in its neighborhood is known as
surroundings .
• The system is separated from the surroundings by a
real or imaginary boundary through which exchange
of energy or matter may take place.
• Interactions between the system and its surrounding
can be classified as: (i) Open , (ii) Closed or (iii)
Isolated system.
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1.2.1 Types of System
A. Open system: Exchange both energy and matter
with its surroundings.
B. Closed system: Can exchange energy but not matter
with the surroundings.
C. Isolated System: There is no transfer of either
energy or matter with the surroundings.
(A)
(B)
(C)
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1.3 SAQ-1
Classify the following systems as open, closed or
isolated.
1) Nitrogen and hydrogen reacting to form ammonia in
a sealed tube.
2) Potassium chlorate and manganese dioxide are
heated in an unsealed test tube to form potassium
chloride and oxygen.
3) A glass vial containing conc. Hydrochloric acid is
broken inside water in a beaker.
4) A glass vial containing hydrochloric acid is broken
inside a solution of sodium hydroxide kept in a
closed polythene bottle insulated with a cork.
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1.5 Properties of a System
The properties associated with a macroscopic system are called
thermodynamic properties or variables
Two main categories:
1) EXTENSIVE PROPERTIES:- Depend upon the amount/quantity
of matter present in the system e.g.mass, volume, energy etc.
2) INTENSIVE PROPERTIES: depends on characteristics of matter
but independent of its amount e.g. pressure, temperature, density,
viscosity, surface tension, refractive index, electromotive force,
chemical potential, specific heat, etc.
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1.4 State functions
Internal energy, a state function, depends only on the present
state and NOT on the path by which it arrived at that state.
Example:
The internal energy of 50g of water at 25ºC is the same whether
the water is cooled from a higher temperature to 25ºC or is
obtained by melting 50 g of ice and then warming to 25ºC
The value of a state function does not depend on the particular
history of the sample, only on its present condition. The properties
that depend only on the state of the system and not on its past
history are called STATE FUNCTIONS.
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STATE FUNCTION
1.
Must have a definite value for a given system in a
given state, independent of the history of that system
and
2.
Change by a definite amount in a given change of state,
independent of the path along which that change proceeds.
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Thermodynamic process:
The method by which the state of a system is changed is called
a ‘process’. It can be effected by changing any one of the state
variables viz P, T, C etc.
Isothermal process: It is a process carried out at constant
temperature dT=0
Adiabatic process: It is a process in which no exchange of heat
takes place. The temperature of the system may increase or
decrease during adiabatic process. dq=0
Isobaric process: Carried out at constant pressure
Isochoric process: Carried out at constant volume
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Reversible process:
A process which occurs infinitely slowly, can be reversed along
the same path and performs maximum work is referred to as
being thermodynamically reversible.
It is a process which takes place infinitesimally slowly so that
the system is in thermodynamic equilibrium at any instant of
the change.
Since the process is carried out extremely slowly the properties
of the system remain virtually unchanged and the direction may
be reversed by small change in a variable like temperature,
pressure etc.
The driving force is greater than the opposing force only by a
infinitesimal quantity and hence the process would require
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infinite time for completion.
Irreversible process:
It is a process which takes place rapidly or spontaneously so
that it is in equilibrium with the surroundings. The driving force
differ from the opposing force by a large amount and hence it
cannot be reversed unless some external force is applied. All
the natural processes are irreversible.
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Reversible process:
.
S/No
Reversible Process
Irreversible process
1
Driving force and opposing
force differ by a large amount
It is a rapid process
The work obtained is less
It is a real process
It has only two steps i.e.
initial and final
It occurs in only one direction
It cannot be reversed
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3
4
5
6
7
Driving force and opposing
force differ by small amount
It is a slow process
The work obtained is more
It is an imaginary process
It consists of many steps
It occurs in both directions
It can be reversed by
changing thermodynamic
variables
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SAQ-2:
Which of the following are extensive properties
and which are intensive properties?
a. Viscosity
b. Weight
c. Mass
d. Volume of a solid
e. Temperature
f. Pressure
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1.6 The First law of Thermodynamics:
Define first law of thermodynamics both verbally and by means
of a equation
In general, energy can be converted from one form to another.
But energy can neither be created nor destroyed: Energy is
conserved.
• The total energy lost by a system equals the total energy
gained by its surroundings. This is the law of conservation of
energy also known as the first law of thermodynamics
Alternatively:
The total energy of a system and its surroundings must remain
constant, although it may be changed from one form to
another"
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Internal Energy
• The total energy of a system is the sum of all the kinetic and
potential energies of its component parts. For the system of
H2 and O2 molecules, the total energy include not only the
motions and interactions of the H2 and O2 molecules
themselves but also of the component nuclei and electrons.
• The total energy is called the internal energy of the system
• Because there are so many types of motions and interactions,
we cannot determine the exact energy of any system of
practical interest. We can, however, measure the change in
internal energy that accompany chemical and physical
processes.
• ΔE = Efinal - Einitial
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In terms of the work performed (w), the heat (q) transferred and the
change in internal energy (ΔE) that is stored in the system; first law
of Thermodynamics is mathematically expressed as:
ΔE = q – w or q = ΔE + w
For very small changes: dE = q - w
Where dE is the small change in the internal energy, q is the small
amount of heat absorbed and w is the small amount of work done
by the system.

This implies that "of the total heat, q, absorbed from the
surroundings a portion raises the internal energy of the system
by E and the rest is used as the work of expansion".

Heat transfer may occur by conduction, convection or radiation
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Summary of sign conventions
Positive (+)
Negative (-)
Heat (q)
System gains
energy
System loses
energy
Work (w)
Work done by the Work done on the
system
system
Change in
internal energy
(ΔE)
Internal energy of Internal energy of
the system
the system
increases
decreases
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Expansion or compression work :
• The external pressure Pex is
equivalent to a weight pressing
on the piston, and the force
opposing expansion is:
ork (w) = F x d
But F=PexA, then:
w = PexAdz ⇒ w = PexΔV
When a piston of area A moves out
through a distance dz it sweeps out a
volume dV = Adz.
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Question-2:
When one mole of water, at 100 0C and 1 atm. pressure, is
converted to steam (at 100 0C) the amount of heat absorbed is
40670 J. Calculate ΔE for the change.
Solution 2:
Work done in expanding against atmospheric pressure:
w = Pext (V2 – V1)
Given: V1 = Volume of one mole of a liquid water
= 18 cm3 = 18 x 10-6 m3
V2 = volume of 1 mole of water vapour at 100 oC
From Gas laws and at constant pressure: V1/T1 = V2/T2
Also, at STP (273 K and 1 atm pressure) 1 mole of ideal gas
occupies ≈ 0.0224 m3
Therefore: V2 = 0.0224 * (373/273) = 0.0306 m3
Then; Pext = 101325(0.0306 – 18x10-6) = 3099 J
From First Law: ΔE = q – w = (40670-3099) = 37571 J.
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Question-3:
A balloon is filled with helium is heated and expands, performing
370 J of work on the surrounding gases in the atmosphere. The
internal energy of helium increases by 1220 J. How much heat did
helium absorb?
Solution 3:
ΔE = q - w
Given: w = 370 J, ΔE = 1220 J
1220 J= q - 370 J
q = 1590 J
SAQ-3
Two moles of an ideal gas at 273 K and 101325 Nm-2 pressure
expand from 0.056 m3 to 0.28 m3 at the same temperature.
Determine the work done.
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Question-4:
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INTERNAL ENERGY OF AN IDEAL GAS
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Question-5:
0.1 mole of an ideal gas is expanded isothermally at 273 K from 3
dm3 to 5 dm3. Determine the energy (q) absorbed from the
surroundings.
Solution 5:
Given: n= 0.1 mole, T= 273, V1 = 3 dm3 and V2 = 5 dm3
Using equation (7)
q = 0.1 * 8.314 * 273 ln (5/3)
= 115.94 J
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Graphical Illustration
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Molar heat capacities at constant pressure and
constant volume (more equations):
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Relationship between Cv and Cp
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Reversible adiabatic expansion of an ideal gas
Derives the temperature-volume and pressure-volume relations
when a fixed amount of an ideal gas is expanded or compressed
without heat.
Therefore, an adiabatic change is a change that can be carried
out reversibly such that no heat enters or leaves the system. In
an adiabatic expansion of a gas the system does external work
(of expansion), and since no heat can be taken up, the necessary
energy comes from the kinetic energy of the molecules. The
decrease in the value of the latter results in a fall in the
temperature of the system. On the other hand, in an adiabatic
compression the temperature of the gas will rise.
In mathematical terms (From First Law)
dE = -dw
Equating: CvdT = -PdV
(20)
(21)
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Adiabatic changes (cont.)
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Adiabatic changes (Summary.)
Cv/ R
V1T1
 V2T2
Cv/ R
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Question-6:
Two moles of an ideal monatomic gas at 300 K are compressed
adiabatically to one quarter of the original volume. What is the
temperature of the gas after compression.
SOLUTION :
For a monatomic gas;
Cv=3/2R
Then we are given;
T1=300 K, V2=V1/4, T2=?
Equation (24)
CVln T2/T1 = -RlnV1/V2
Substitution gives
3/2RlnT2/300 = -RlnV1/4V1
lnT2/300 = 2/3ln4
Hence: T2 = 755.95 K Ans
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Question-7:
To what pressure must a given volume of Helium originally at 100oC
and 1 atm. Pressure be adiabatically compressed in order to raise its
temperature to 400oC?
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Solution 6
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Tutorial Question. 1…
c) One mole of an ideal gas at 25 °C and a pressure of 10 x 105
Nm-2 expands isothermally to a pressure of 1 x 105 Nm-2
i. Calculate the final volume, w and q if the expansion
were carried out reversibly.
ii. Calculate w and q for an irreversible expansion
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Tutorial Question.
Question 1
a) Explain clearly what is meant by a thermodynamically
reversible process.
Solution:
a) A thermodynamic reversible process occurs in infinitely slowly,
can be reversed along the same path and performs maximum
worm.
b) Explain the thermodynamic meaning of a system,
distinguishing among the open, closed and isolated systems.
Solution:
b) A system is part of the universe, the properties of which are
under investigations and has a definite boundary.
 Open system: Exchange both energy and matter with
surroundings.
 Closed system: Exchange energy with surroundings,
matter NOT exchanged.
 Isolated system: No transfer of either energy or matter
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with surroundings
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Tutorial Question. 2
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Tutorial Descriptive Questions
a) Explain clearly what is meant by a thermodynamically
reversible process.
Solution:
a) A thermodynamic reversible process occurs in infinitely slowly,
can be reversed along the same path and performs maximum
worm.
b) Explain the thermodynamic meaning of a system,
distinguishing among the open, closed and isolated systems.
Solution:
b) A system is part of the universe, the properties of which are
under investigations and has a definite boundary.
 Open system: Exchange both energy and matter with
surroundings.
 Closed system: Exchange energy with surroundings,
matter NOT exchanged.
 Isolated system: No transfer of either energy or matter
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with surroundings
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SAQ-3
1) To what pressure must a given volume of Helium originally
at 100oC and 1 atm. Pressure be adiabatically compressed in
order to raise its temperature to 400oC?
2) At 0oC and 1 atm. Pressure the volume of 1 mole of an ideal
monatomic gas is 22.415 litres. The gas is expanded until its
pressure is 0.4 atm. by a reversible adiabatic process.
Determine the final volume.
3) The volume of a sample of an ideal monatomic gas at 0oC is
44.83 litres. To what volume must the gas be compressed
adiabatically so as to attain a temperature of 30oC
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Summary
 Energy cannot be created or destroyed but can be transformed
from one form to another.
 The heat absorbed by the system is partly stored within it as
internal energy and partly spent on doing work on the
surroundings.
 Heat absorbed by the system is taken to be positive.
 Internal energy is a state function.
 Transformations of systems can take place reversibly or
irreversibly.
 Transformations taking place in an infinite number of steps
permitting equilibrium, to prevail at every stage are called
reversible transformations. When the transformation is rapid
and sudden, it is said top be spontaneous or irreversible.
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Summary (cont.)
 Work is calculated as a product of the pressure against which
work is done and the change in volume. For an irreversible
process the pressure (external) is often constant. For a
reversible process, the pressure changes during the course of
the process.
 The heat absorbed by a system at constant volume is equal to
the internal energy change while that absorbed by the system
at constant pressure is the enthalpy change.
 The temperature coefficient of the internal energy change at
constant volume is the heat capacity at constant volume and
the temperature coefficient of the enthalpy change at constant
pressure is the heat capacity at constant pressure.
 An adiabatic change is that which can be carried out
reversibly such that no heat enters or leaves the system.
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