Thermodynamics

THERMODYNAMICS
Mrs. Praveen Garg
VITS College, Satna

INTRODUCTION
• Thermodynamics is the branch of physics that deals with the
relationships between heat and temperature and their relation to
energy and work.
• Thermo – Heat , Dynamics – Motion or move
• In particular, it describes how thermal energy is converted to and
from other forms of energy and how it affects matter.

Thermodynamics
Heat, work, different form of energy, heat changes
Heat flow from high temperature to low temperature
Example:
Gas
Pressure, Enthalpy
Temperature, Entropy
Volume, Gibbs energy
Thermodynamics

Branches of thermodynamics
 Classical thermodynamics
 Statistical mechanic
 Chemical thermodynamics
 Equilibrium thermodynamics

Classical Thermodynamics:
• In classical thermodynamics, the behavior of whole matter is
analyzed, such as temperature and pressure. It helps to calculate
properties and to predict the characteristics of the matter that is
undergoing the process.
Statistical Thermodynamics:
• In statistical thermodynamics, study the different properties of
each and every molecule and to characterize the behavior of a
each molecules.

Macroscopic analysis Microscopic analysis

Equilibrium Thermodynamics:
• Equilibrium thermodynamics is the study of transformations
of energy and matter as they approach the state of
equilibrium.
Chemical Thermodynamics:
• Chemical thermodynamics is the study of how work and heat
relate to each other both in chemical reactions and in changes
of states.

Basic Concepts of Thermodynamics –
Thermodynamic Terms
 Systems:
A system is a specific portion of
matter or substances, which
have to be study.
 Surrounding:
Except system everything in
universe is called surrounding.
 System boundary:
It separate system and
surrounding. The system
boundary may be real or
imaginary, fixed or deformable.
System + Surrounding = Universe

Boundary
Diathermic Adiabatic
Across heat
change easily
(Conducting)
heat can not
flow
(Insulating)
(∆q=0)
Heat flow
High temp.
Low temp.

There are three types of system as:
• Open System – In an open system, the mass and energy both
may be transferred between the system and surroundings.
• Closed System – Across the boundary of the closed system,
the transfer of energy takes place but the transfer of mass
doesn’t take place. Energy can exchange in the form of heat
and work.
Refrigerator, compression of gas in the piston-cylinder
assembly are examples of closed systems.
• Isolated System – An isolated system cannot exchange both
energy and mass with its surroundings. The universe is
considered an isolated system.

Real
Imaginary
Mass change – form of vapour
Due to open boundary
Conducting--
Heat change
Mass can not change
due to closed boundary
Real,
adiabatic
(Insulated)
∆q=0

SYSTEM
(variable/parameter/properties)
Extensive
 Its depend on amount of matter,
size of system or sample.
 Example:
• Mass (water)
• Volume (water)
• No. of moles (solution)
• Heat capacity,
• Energy- Heat, Work, Enthalpy,
Entropy, Internal Energy, Gibbs
free energy etc.
Intensive
 Does not depend on size of
sample/amount of matter.
 Example:
• Density (Glass of water)
• Specific heat capacity
• Temperature (Whole or small
part of room)
• Pressure, Viscosity (constant)
• Concentration (sugar solution)
• Molarity, Normality, Mole
fraction, Molar entropy, molar
enthalpy, pH etc.

Extensive (X)
Extensive (Y)
= Entensive (Z)
Example:
Molarity =
Molar entropy= entropy/mole
No. of moles
Volume

State variable/ State function/State parameter
• Condition or state of system is defined as variables.
• Example:
Volume
Gas Pressure Variables
Temperature
State
• Variables of state: pressure(P), temperature(T), volume(V), no. of
moles, enthalpy(H), entropy(S), internal energy (E), Gibbs free
energy (G) etc.
• Function of state can be change due to change in variable.
• These changes in variables depend on initial and final condition
of state and does not depend on path.

State function:
∆T = Tf – Ti
∆P = Pf – Pi
∆V = Vf – Vi
• Changes in temperature, pressure, volume depend on initial (i)
and final (f) condition.
• But Heat and Work depend only in path function, it does not
depend on initial and final condition.
• In the case of any cyclic process, change in state variables is
always zero (no change).
∆T = 0
∆P = 0
∆V = 0
But: ∆q ≠ 0
∆w ≠ 0

Types of thermodynamic process
 Isothermal process: Occur in constant/same temperature.
∆T = 0
If ideal gas, ∆T = 0, ∆E = 0
Internal energy (E = K.E. ∝ T)
 Isobaric process: Occur in constant pressure.
 Isochoric process: Occur in constant volume (closed vessel).
 Adiabatic process: In this process, heat can not change.
∆q = 0
Note: These all process can be reversible or irreversible.

Reversible
• Which occur in infinite
• Infinite time to finish
• Small steps
• Slow process
• Each step in equilibrium
exist always.
Irreversible
• At once
• Finite time to finish
• Fast process
• Equilibrium depend on
initial and final

1) Zeroth law of thermodynamics:
When two systems are each in thermal equilibrium with a third
system, the first two systems are in thermal equilibrium with
each other.
2) First law of thermodynamics:
It is the law of conservation of energy. The change in a system’s
internal energy is equal to the difference between heat added to
the system from its surroundings and work done by the system
on its surroundings.
•
LAWS OF THERMODYNAMICS

3) Second law of thermodynamics:
• Heat does not flow spontaneously from a colder region to
a hotter region, or, equivalently, heat at a given
temperature cannot be converted entirely into work.
• Consequently, the entropy of a closed system, or heat
energy per unit temperature, increases over time toward
some maximum value.
• Thus, all closed systems tend toward an equilibrium state
in which entropy is at a maximum and no energy is
available to do useful work.

4) Third law of thermodynamics:
• The entropy of a perfect crystal of an element in its most
stable form tends to zero as the temperature
approaches absolute zero.
• This allows an absolute scale for entropy to be established
that, from a statistical point of view, determines the
degree of randomness or disorder in a system.

• State in which no spontaneous change (macroscopic changes)
are observed with respect of time.
• It may be three types:
• Thermal equilibrium- equal temperature
• Mechenical equilibrium- equal force/pressure
• Chemical equilibrium- equal chemical force
EQUILIBRIUM

“Systems that are in thermal equilibrium exist at the same
temperature”.
• When two or more bodies at different temperatures are brought
into contact then after some time they attain a common
temperature and they are said to exist in thermal equilibrium.
Thermal equilibrium

• When a body ‘A’ is in thermal equilibrium with another body ‘b’,
and also separately in thermal equilibrium with a body ‘C’, then
body ‘B’ and ‘C’ will also be in thermal equilibrium with each
other. This statement defines the zeroth law of thermodynamics.
The law is based on temperature measurement.
• Zeroth law of
thermodynamics states
that, if the systems are in
thermal equilibrium, no
heat flow will take place.
• This law is mostly used
to compare temperatures
of different objects.
ZEROTH LAW OF THERMODYNAMICS

• The most common application of the zeroth law of
thermodynamics can be seen in thermometers.
• Another example of the zeroth law of thermodynamics is
when you have two glasses of water. One glass will have
hot water and the other will contain cold water. Now if we
leave them in the table for a few hours they will attain
thermal equilibrium with the temperature of the room.
Application of Zeroth law

First Law of Thermodynamics
• The First Law of Thermodynamics states that heat is a form of
energy, and thermodynamic processes are of conservation of
energy.
• This means that heat energy cannot be created or destroyed. It
can be transferred from one location to another and converted to
other forms of energy.
• The First Law says that the internal energy of a system has to be
equal to the work that is being done on the system, plus or minus
the heat that flows in or out of the system. "So, it’s a restatement
of conservation of energy."
• This is expressed mathematically as:
ΔU = Q – W
where ΔU is the change in the internal energy, Q is the heat
added to the system, and W is the work done by the system.

First law of thermodynamics:
ΔEuniv = ΔEsys + ΔEsurr = 0
The changes of in the internal energy of a system is equal to
the sum of the heat gained or lost by the system and the work
done by or on the system.
ΔEsys = q - w (closed system)
Example: Beaker of water on hot plate
When hot plate is on, temp and internal energy of
the system increase, ΔE is positive. When hot plate is off, temp
and internal energy decrease, ΔE is negative.

Application of First law
• Light bulbs transform electrical energy into light energy
(radiant energy).
• One pool ball hits another, transferring kinetic energy and
making the second ball move.
• Plants convert the energy of sunlight (radiant energy) into
chemical energy stored in organic molecules.

Limitations of First Law of
Thermodynamics
• The limitation of the first law of thermodynamics is that
it does not say anything about the direction of flow of
heat.
• It does not say anything whether the process is a
spontaneous process or not.
• The reverse process is not possible. In actual practice, the
heat doesn’t convert completely into work.

HEAT
• Heat is energy transferred between substances or systems due to a
temperature difference between them.
• Heat is conserved as a form of energy.
• Heat is the transfer of thermal energy between systems,
while work is the transfer of mechanical energy between two
systems.
• Heat is transfered via solid material (conduction), liquids and gases
(convection), and electromagnetical waves (radiation).
• Heat describes the transfer of thermal energy between molecules
within a system and is measured in Joules. Heat measures how
energy moves or flows.
• Temperature describes the average kinetic energy of molecules
within a material or system and is measured in Celsius (°C),
Kelvin(K), Fahrenheit (°F).

INTERNAL ENERGY
• Internal energy, in thermodynamics, the property or state
function that defines the energy of a substance.
• The value of the energy depends upon the state of the substance
and not upon the nature of the processes by which it attained that
state.
• According to first law of thermodynamics, system undergoes a
change of state only due to work is involved, the work is equal to
the change in internal energy.
• The law also implies that if both heat and work are involved in
the change of state of a system, then the change in internal
energy is equal to the heat supplied to the system minus the work
done by the system.
• The internal energy as a sum of terms that can be interpreted
as kinetic energy, potential energy, and chemical energy.

ENTHALPY
• Enthalpy, the sum of the internal energy and the product of the
pressure and volume of a thermodynamic system.
• Enthalpy is an energy-like property or state function—it has the
dimensions of energy (and is thus measured in units
of joules or ergs).
• Its value is determined entirely by the temperature, pressure,
and composition of the system.
• The enthalpy (H) equals the sum of the internal energy (E) and
the product of the pressure (P) and volume (V) of the system:
H = E + PV
• The change in internal energy is equal to the heat transferred to,
less the work done by the system. If the only work done is a
change of volume at constant pressure, the enthalpy change is
exactly equal to the heat transferred to the system.

ENTROPY
• Entropy, the measure of a system’s thermal energy per
unit temperature.
• Entropy is a function of the state of the system, so the change in
entropy of a system is determined by its initial and final states.
• The amount of entropy is also a measure of the molecular
disorder, or randomness, of a system.
• The entropy of a substance can be obtained by measuring the heat
required to raise the temperature a given amount, using a
reversible process.
• Enthalpy is the measure of total heat present in the
thermodynamic system where the pressure is constant
while Entropy is the measure of disorder in a thermodynamic
system.
where Q is the heat content and T is the temperature.
ΔS = ΔQ /T

• Zero entropy means perfect knowledge of a state ; no motion,
no temperature, no uncertainty.
• Entropy is not conserved because it can be created. But once it
is created and dissipated in surroundings the action cannot be
reversed, so therefore entropy always keep on increasing.
• As we know laws of conservation of mass and energy say that
both the quantities are uncreatable and undestroyable.

• This is irreversible process.
• The second law of thermodynamics states that the heat energy
cannot transfer from a body at a lower temperature to a body at a
higher temperature without the addition of energy.
• The second law of thermodynamics states that the entropy of an
isolated system is always increasing
• The entropy of the system is measured in terms of the changes
the system has undergone from the previous state to the final
state.
• Thus the entropy is always measured as the change in entropy of
the system denoted by ∆S.
• The universe is an isolated system. This means the entropy of the
universe will always increase. The entropy change of the universe
is the sum of the entropy change of the system and the entropy
change of the surroundings.
Second Law of Thermodynamics

Causes of increase in entropy of the closed system are:
• In a closed system, the mass of the system remains constant
but it can exchange the heat with surroundings.
• Any change in the heat content of the system leads to
disturbance in the system, which tends to increase the entropy
of the system.
• Due to internal changes in the movements of the molecules of
the system, there is disturbance inside the system.
• This causes irreversibilities inside the system and an increase
in its entropy.

• Exergonic Reaction: It is spontaneous reaction that
release energy. The entire reaction is catabolic. The release
of energy (Gibbs free energy) is negative.
• Endergonic Reaction: It is anabolic reaction that
consume energy. It has a positive ΔG because energy is
required to break bonds.
• The free energy gained or lost in a reaction can be
calculated:
ΔG = ΔH – TΔS
Where, G= Gibbs free energy
H= enthalpy
S= entropy
T= temperature

Application of Second law
• This law is applicable to all types of heat engine cycles
including Otto, Diesel, etc. for all types of working fluids used
in the engines. This law has led to the progress of present day
vehicles.
• Another application is refrigerators and heat pumps based on
the Reversed Carnot Cycle. If you want to move heat from a
body at a lower temperature to a body at a higher temperature,
then you have to supply external work.
• Removing heat from the food items in the refrigerator and
throwing it away to the higher temperature atmosphere doesn’t
happen automatically. We need to supply external work via the
compressor to make this happen in the refrigerator.
• Air conditioner is another example.

• The third law of thermodynamics states that the entropy of a
perfect crystal at a temperature of zero Kelvin (absolute zero) is
equal to zero.
• At zero kelvin the system must be in a state with the minimum
energy.
• If the crystal has only one minimum energy state, third law
holds true.
• Entropy is related to the number of microstates, and with only
one microstates available at zero kelvin, the entropy exactly
zero.
Third Law of Thermodynamics

• It helps in the calculation of the absolute entropy of a
substance at any temperature ‘T’.
• These determinations are based on the heat
capacity measurements of the substance.
Application of the third law

Thermodynamics

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