Free Download Link (Copy URL): https://sites.google.com/view/varunpratapsingh/teaching-engagements Dear Students, Please find the Basic Mechanical Engineering (TME-101, 2020-21 Session) Unit One notes in this section. Topic cover in this section are: UNIT-1: Fundamental Concepts and Definitions Definition of thermodynamics, System, Surrounding and universe, Phase, Concept of continuum, Macroscopic & microscopic point of view. Density, Specific volume, Pressure, temperature. Thermodynamic equilibrium, Property, State, Path, Process, Cyclic and non-cyclic processes, Reversible and irreversible processes, Quasi-static process, Energy and its forms, Enthalpy.

1. UNIT-1
THERMODYNAMICS
NOTES
By
Mr. Varun Pratap Singh
Assistant Professor
Department of Mechanical Engineering
College of Engineering Roorkee

2. Syllabus
UNIT-1: Fundamental Concepts and Definitions
Definition of thermodynamics, System, Surrounding and universe, Phase, Concept of continuum,
Macroscopic & microscopic point of view. Density, Specific volume, Pressure, temperature.
Thermodynamic equilibrium, Property, State, Path, Process, Cyclic and non-cyclic processes,
Reversible and irreversible processes, Quasi static process, Energy and its forms, Enthalpy.

3. Introduction
The field of science, which deals with energies possessed and exchanged by a system of gases and
vapours, their conversion in terms of heat and work and their relationship with properties of system
is called thermodynamics. Thus in thermodynamics, we deal with conversion of energy from one
form to the other, mainly heat into work or vice versa. In this context, we must know the overall
scenario of conversion of available forms of energy into usable forms as shown below:
Figure: Energy conversion chart

4. Application of Thermodynamics
 Power Generation (Thermal/Nuclear Power plants)
 Automobiles (Petrol/ Diesel/LPG/CNG engines)
 Processing Industries (Steam generation, Refrigeration systems)
 Gas Compressors (Air compressors)
Thermodynamics Definition:
Is the science of energy transfer and its effect on the physical properties of substance.
Thermodynamics is primarily concerned with two forms of energy i.e. "Heat & Work". It is
basically based on four laws of thermodynamics known as zeroth, first, second and third law.
Thermodynamic principles are used in the designing of energy converting devices, such as, steam
engines, IC. engines, steam and gas turbines, refrigerators, air conditioners, fuel cells etc.
Thermodynamic System:
It is defined as a definite region or space (area) or a quantity of matter on which attention is
focus for the study of work and heat transfer and conversion.
Figure: A Thermodynamic System
Boundary:
It is defined as actual or hypothetical envelope enclosing the system. The boundary may be real
or imaginary. It may be fixed or movable in nature.
Surrounding :
It is defined as the space and matter extremal to the thermodynamic
system and outside boundary.
Universe:
System + Boundary + Surrounding

5. Types of System:
1. Open System:
A system is called an open system if the mass as well as the energy
transfers and its boundaries. Examples 1. IC engines 2 Turbines 3. Nozzles etc.
2. Closed System:
In this type of system no matter or mass may enter or leave across the system boundary, there
may be energy (heat or work) transfer across the system boundary
Examples 1.A certain quantity of fluid in a cylinder bounded by a piston constitutes
a closed system. 2. Thermal power plants
Figure 2: Types of Systems
3. Isolated System:
Neither mass nor energy transfers can take place through its boundaries. Example Fluid
enclosed Insulated in a perfectly insulated closed vessel (Thermos -Flask).
Figure 2.1: Types of Systems
4.Adiabatic System:
the system which is thermally insulated from its surrounding is called adiabatic system it can
power exchange work with the surrounding if it does not, it become an isolated system.

6. Other classification of system:
Homogeneous system:
system or quantity of mass which consists of a single phase is called homogeneous system.
Example mixture of air and water vapor, solution of NH3 in water.
Heterogeneous System:
A system which consists of two or more phase, called heterogeneous system. Example 1 water
and steam 2. ice and water 3 water and oil etc.
Phase:
A phase is a quantity of matter which is h homogeneous throughout in chemical composition and
physical structure.

7. Macroscopic Vs Microscopic Viewpoint
Concept of continuum - thermodynamics, Mechanical Engineering
Even the simplification of matter into molecules, atoms, electrons, etc. is too complex a picture for
many problems of thermodynamics. Thermodynamics doesn't make any hypotheses about the
structure of the matter of system. The volumes of system considered are quite large as compared to
molecular dimensions. The system can be regarded as continuum. The system is supposed to
contain continuous distribution of matter. There are no voids and cavities present. The pressure,

8. temperature, density and other properties are average values of action of several molecules and
atoms. This kind of idealization is a must for solving most of the problems. The laws and concepts
of thermodynamics are not dependent of structure of matter.
In accordance to this concept there is minimum limit of volume up to which the property remains
continuum. Below this volume, there is sudden change in the value of the property. This type of
region is called as region of discrete particles and the region for which the property are maintain is
called as region of continuum. The volume up to which continuum properties are maintained is
called as continuum limit.
For Example: If we measure the density of a substance for a large volume (ν1), the value of density
is(ρ1). If we go on reducing the volume by v', below which ratio äm/äv deviates from its actual
value and value of äm/äv is large or small.
Thus according to this concept design could be defined as:
Here Δ is the volume of the fluid element and m is the mass.
Thermodynamic equilibrium:
A system is said to be in a state of thermodynamic equilibrium if the values of the microscopic
properties (such as temperature pressure etc.) at all the point are same. thermodynamic
equilibrium is a complete equilibrium and it is including the following equilibrium:
1. Mechanical equilibrium
there are no unbalanced forces within the system or between the surrounding and system is known
as mechanical equilibrium. In this case magnitude of of resultant force is zero.

9. 2.Chemical equilibrium
in this equilibrium NO2 chemical reaction take place in the system and the chemical composition
which is same throughout the system does not vary with time.
3. Thermal equilibrium
the temperature of the system does not change with time and has the same value at the all point
in the system its mean temperature gradient is not available or can be considered as zero in thermal
equilibrium.
4. Electrical equilibrium
If there exist and uniformity of electrical potential throughout the system it can be treated as in
electrical equilibrium.
Property of substance
Any characteristics of a substance which can be observed on major is called property of
substance. Example pressure temperature mass specific volume etc. A Property is any
observable characteristic of a system.
(a) intensive property:
these properties do not depend on the mass of the system. Intensive properties are generally
denoted by lowercase letter with major exceptions pressure” P” and temperature “T”, which are
always intensive.
Example temperature (T), pressure (P), specific volume (v), density (ρ) etc.
(b) extensive property:
these properties depend on the mass of the system. If mass changes, there value also changes.
These are generally denoted by upper case letter with major exceptions like mass “m”, number
of moles “n”.
Examples volume (V), energy (E), total enthalpy (H), entropy (S), potential energy (PE), kinetic
energy (KE) etc.
The ratio of two extensive property of a homogeneous system is an intensive property. Like ratio
of mass and volume is density which is intensive property. when an extensive property is divided

10. by the mass or number of moles of a substance forming the system, the resulting intensive
property is called a specific property. Example: specific volume, specific energy etc.
State, Change of State, Path, Process, Cycle
State: State is the condition of the system at an instant of time as described on measured by it
properties for each unique condition of a system.
Change of State:
Any operation in which one or more of the properties of a system changes is called a change of
state.
Path:
thermodynamic system passing through a series of state constitutes a path.
Process:
when a system changes its state from one equilibrium state to another equilibrium state, then the
path of successive States through which the system has passes is known as thermodynamic
process.
Cyclic process:
cyclic process consists of number of processes incorporated to achieve the target in
thermodynamics cycle the initial and final state of process are same.
Reversible and Irreversible Process
In reversible process, the initial state and final states are same, together with all energy
transfer or transfer during the process can be completely restored in both system and
surroundings. No heat loss, friction loss and traces of path cannot be observing after reversing
the process. In reality no thermodynamic process is completely reversible. Same path and traces
can be followed to achieve reversibility.
The reversal process having following condition which must be satisfied:
1. The process processed at an extremely slow rate so that it goes inside with the Quasi static
process.
2. The system should be free from dissipative forces like friction, inelasticity, viscosity,
electrical resistance etc. This is because energy spent against such forces cannot be recovered.
3. No heat, work and friction losses should occur during process.
4. A Reversible process should not leave any traces to show that the process had ever occurred.
5. It should not leave any history behind its mean no change should be created on environment.

11. Figure: (a) Reversible Process (b) Irreversible Process
In Irreversible process,
If the energy transfer is not restored in both system and surroundings, then process is called
irreversible. Heat transfer friction losses and traces of path can be observed in irreversible
process. All real or natural process are irreversible in nature. In irreversible process same path
cannot be followed if try to achieve reversibility.
Quasi-static process
In thermodynamics, a quasi-static process is a thermodynamic process that happens slowly
enough for the system to remain in internal equilibrium. An example of this is quasi-static
compression, where the volume of a system changes at a slow rate enough to allow the pressure
to remain uniform and constant throughout the system.
Quasi means ‘as if/ almost constant’, This process is a succession of equilibrium states and
infinite slowness is its characteristic feature. The definition given above is closer to the intuitive
understanding of the word “quasi-” (as if) “static” and remains technically different from
reversible processes.
Only in Quasi-static process can we define intensive quantities (like Pressure, Temperature,
Specific volume, Specific entropy) of the system at every instant during the whole process;
Otherwise, since no internal equilibrium is established, different parts of the system would have
different values of these quantities.
Any reversible process is a quasi-static one. However, quasi-static processes involving entropy
production are not reversible.

12. Path Function and Point Function
Path function
A Path function is a function whose value depends on the path followed by the thermodynamic
process irrespective of the initial and final states of the process. Their magnitudes depend on
the path followed during a process as well as the end states. Work (W), heat (Q) are path
functions.
An example of path function is work done in a thermodynamic process. Work done in a
thermodynamic process is dependent on the path followed by the process. A path function is an
inexact or imperfect differential.
In the P-V diagram given above we can easily see that for the same initial and final states of the
system, work done in all the three process is different.
For process A work done is b2A1a
For process B work done is b2B1a
For process C work done is b2C1a
Another example of path function is heat.

13. Point function
A Point function (also known as state function) is a function whose value depends on the final and
initial states of the thermodynamic process, irrespective of the path followed by the process.
Example of point functions are density, enthalpy, internal energy, etc.
A point function is a or we can say all the properties of the system are point functions.
Point functions are exact or perfect differential.
Note: Since a point function is only dependent on the initial or final state of the system, hence in a
cyclic process value of a thermodynamic function is zero, or change in thermodynamic property is
zero.
Difference between point function and path function
Sr.
no. Point Function Path Function
1
Its values are based on the state of the system
(i.e. pressure, volume, temperature etc.)
Its values are based on how that particular
thermodynamic state is achieved.
2
No matter by which process the state is obtained,
its values will always remain the same.
Different processes to obtain a particular
state will give us different values.
3
Only initial and final states of the process are
sufficient
We need to know exact path followed by
the process
4 Its values are independent of the path followed
Its values are dependent on the path
followed
5 It is an exact or perfect differential It is an inexact or imperfect differential.
6 Its cyclic integral is always zero Its cyclic integral may or may not be zero
7 It is property of the system It is not the property of the system
8
Its examples are density, enthalpy, internal
energy, entropy etc. Its examples are Heat, work etc.

14. Work, Heat, and Heat Capacity
Work in Thermodynamics
In thermodynamics, work performed by a system is the energy transferred by the system to its
surroundings. Kinetic energy, potential energy and internal energy are forms of energy that are
properties of a system. Work is a form of energy, but it is energy in transit. A system contains
no work; work is a process done by or on a system. In general, work is defined for mechanical
systems as the action of a force on an object through a distance.
Below mentioned figure shows a gas confined to a cylinder that has a movable piston at one end. If
the gas expands against the piston, it exerts a force through a distance and does work on the piston.
If the piston compresses the gas as it is moved inward, work is also done—in this case, on the gas.
The work associated with such volume changes can be determined as follows: Let the gas pressure
on the piston face be p. Then the force on the piston due to the gas is pA, where A is the area of the
face. When the piston is pushed outward an infinitesimal distance dx, the magnitude of the work
done by the gas is:
This integral is only meaningful for a quasi-static process, which means a process that takes place
in infinitesimally small steps, keeping the system at thermal equilibrium. (We examine this idea in
more detail later in this chapter.) Only then does a well-defined mathematical relationship (the
equation of state) exist between the pressure and volume. This relationship can be plotted on

15. a pV diagram of pressure versus volume, where the curve is the change of state. We can
approximate such a process as one that occurs slowly, through a series of equilibrium states. The
integral is interpreted graphically as the area under the pV curve (the shaded area of (Figure)). Work
done by the gas is positive for expansion and negative for compression.
When a gas expands slowly from V1 to V2 the work done by the system is represented by the shaded
area under the pV curve.
Consider the two processes involving an ideal gas that are represented by
paths AC and ABC in (Figure). The first process is an isothermal expansion, with the volume of the
gas changing its volume from V1 to V2. This isothermal process is represented by the curve between
points A and C. The gas is kept at a constant temperature T by keeping it in thermal equilibrium
with a heat reservoir at that temperature. From (Figure) and the ideal gas law,
PΔV Work
PΔV Work is equal to the area under the process curve plotted on the pressure-volume
diagram.
Pressure-volume work (or pΔV Work) occurs when the volume Vof a system changes.
The pΔV Work is equal to the area under the process curve plotted on the pressure-volume
diagram. It is known also as the boundary work. Boundary work occurs because the mass of the
substance contained within the system boundary causes a force, the pressure times the surface area,
to act on the boundary surface and make it move. Boundary work (or pΔV Work) occurs when
the volume V of a system changes. It is used for calculating piston displacement work in a closed

16. system. This is what happens when steam, or gas contained in a piston-cylinder device expands
against the piston and forces the piston to move.
Heat in Thermodynamics
While internal energy refers to the total energy of all the molecules within the object, heat is the
amount of energy flowing from one body to another spontaneously due to their temperature
difference. Heat is a form of energy, but it is energy in transit. Heat is not a property of a system.
However, the transfer of energy as heat occurs at the molecular level as a result of a temperature
difference.
In general, when two objects are brought into thermal contact, heat will flow between
them until they come into equilibrium with each other. When a temperature difference does
exist heat flows spontaneously from the warmer system to the colder system. Heat transfer
occurs by conduction or by thermal radiation. When the flow of heat stops, they are said to be
at the same temperature. They are then said to be in thermal equilibrium.
The amount of heat transferred depends upon the path and not simply on the initial and final
conditions of the system. There are actually many ways to take the gas from state i to state f.
Also, as with work, it is important to distinguish between heat added to a system from its
surroundings and heat removed from a system to its surroundings. Q is positive for heat added to
the system, so if heat leaves the system, Q is negative. Because W in the equation is the work done
by the system, then if work is done on the system, W will be negative and Eint will increase.
The symbol q is sometimes used to indicate the heat added to or removed from a system per unit
mass. It equals the total heat (Q) added or removed divided by the mass (m).

17. Heat Capacity
Different substances are affected to different magnitudes by the addition of heat. When a given
amount of heat is added to different substances,
their temperatures increase by different amounts.
This proportionality constant between the heat
Qthat the object absorbs or loses and the
resulting temperature change T of the object is
known as the heat capacity C of an object.
C = Q / ΔT
Heat capacity is an extensive property of matter,
meaning it is proportional to the size of the
system. Heat capacity C has the unit of energy per
degree or energy per kelvin. When expressing the
same phenomenon as an intensive property, the heat
capacity is divided by the amount of substance,
mass, or volume, thus the quantity is independent of
the size or extent of the sample.
Specific Heat Capacity
The heat capacity of a substance per unit mass is
called the specific heat capacity (cp) of the substance. The subscript p indicates that the heat
capacity and specific heat capacity apply when the heat is added or removed at constant pressure.
cp = Q / mΔT
Specific Heat Capacity
In the Ideal Gas Model, the intensive properties cv and cp are defined for pure, simple
compressible substances as partial derivatives of the internal energy u(T, v) and enthalpy h(T, p),
respectively:
where the subscripts v and p denote the variables held fixed during differentiation. The
properties cv and cp are referred to as specific heats (or heat capacities) because under certain
special conditions they relate the temperature change of a system to the amount of energy added by
heat transfer. Their SI units are J/kg K or J/mol K. Two specific heats are defined for gases, one
for constant volume (cv) and one for constant pressure (cp).

18. Pressure
Pressure is the force applied perpendicular to the surface of an object per unit area over which
that force is distributed. Gauge pressure is the pressure relative to the ambient pressure.
Mathematically:
P=F/A
where: P is the pressure,
F is the magnitude of the normal force,
A is the area of the surface on contact.
Pressure is a scalar quantity. It relates the vector area element (a vector normal to the surface)
with the normal force acting on it. The pressure is the scalar proportionality constant that
relates the two normal vectors:
Atmospheric Pressure, Gauge Pressure, Vacuum Pressure and Absolute Pressure
Absolute Pressure - The actual pressure at a given position is called the absolute pressure, and it is
measured relative to absolute vacuum (i.e., absolute zero pressure).
Gage Pressure –
Gage pressure is the pressure relative to the atmospheric pressure. In other words, how much above
or below is the pressure with respect to the atmospheric pressure.
Vacuum Pressure –
Pressures below atmospheric pressure is called vacuum pressures and are measured by vacuum
gages that indicate the difference between the atmospheric pressure and the absolute pressure.
Atmospheric Pressure –
The atmospheric pressure is the pressure that an area experiences due to the force exerted by the
atmosphere.
Equations
Pgage = Pabs − Patm Gage Pressure
Pvac = Patm − Pabs Vacuum Pressure
Pabs = Patm + Pgage Absolute Pressure

19. Manometer, Barometer and Anemometers
Barometers, manometers and anemometers are all scientific instruments. Scientists use
barometers and manometers to measure atmospheric pressure, while anemometers measure wind
speed.
Manometers
A manometer is a tube-like device which measures atmospheric measure. There are two types:
closed tube and open tube, but both measure pressure by comparing the pressure exerted by the
atmosphere at one end of the tube with a known pressure at the other. Manometer tubes are
typically filled with mercury.
Barometer
Barometers also measure atmospheric pressure. Mercury barometers are a type of closed-tube
manometer, while aneroid barometers use a small, spring balance to take the measurement. In the
past, mercury barometers were common in family homes where people used them to predict the
weather based on the air pressure reading. Rising air pressure meant good weather was on the
way, while falling pressure might bring rain.
Anemometers
Anemometers are a completely different type of instrument
used to measure wind speed. There are several different
types, but the most common—the cup anemometer—takes
the measurement by recording the number of times the wind
rotates a fan-shaped device.
Thank You