This note contains basic concepts of thermodynamics.

1. THERMODYNAMICS
Presentation by
Dr. Satyajit Mandal

2. What is thermodynamics ?
 The science of energy, that
concerned with the ways in which
energy is stored within a body.
 Energy transformations -mostly
involve heat and work
movements.
 The Fundamental law is the
conservation of energy principle:
energy cannot be created or
destroyed, but can only be
transformed from one form to
another.
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3. System, surroundings and boundary
 System: A quantity of matter or a
region in space chosen for study.
 Surroundings: The mass or region
outside the system
 Boundary: The real or imaginary
surface that separates the system
from its surroundings.
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4. Type of system
Open system
 Open system – both
mass and energy can
cross the selected
boundary.
 Example: an open cup
of coffee
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5. Type of system
Closed system
 Closed system – only
energy can cross the
selected boundary.
 Examples: a tightly
capped cup of coffee.
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6. Type of system
Isolated system
 Isolated system –
neither mass nor
energy can cross the
selected boundary.
 Example: coffee in a
closed, well-insulated
thermos-bottle
Thermally insulated wall.
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7. Properties of a system
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8. Properties of a system
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9. State, Equilibrium and Process
State – a set of properties that
describes the conditions of a system.
Eg. Mass m, Temperature T, volume V.
Thermodynamic equilibrium
system that maintains thermal,
mechanical, phase and chemical
equilibriums.
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10. State, Equilibrium and Process
Process – change from one state to another state.
Process Property held constant
1. isobaric pressure
2. isothermal temperature
3. isochoric volume
4.isentropic entropy
5.Polytropic heat capacity
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11. State, Equilibrium and Process
The prefix “iso”- is often used to designate a process for which a particular property
remains constant.
Isobaric process: A process during which the pressure P remains constant. Pressure is
Constant (ΔP = 0)
Isobaric process:
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12. State, Equilibrium and Process
Isochoric (or isometric) process: A process during which the specific volume V
remains constant.
Isochoric process:
Isothermal process: A process during
which the temperature T remains
constant.
Isothermal process:
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13. State, Equilibrium and Process
Cyclic process - when a system in a given initial
state goes through various processes and finally
return to its initial state, the system has undergone
a cyclic process or cycle.
Reversible process – system can restored to it’s initial state
from final state without the aid of any external agency.
Irreversible process – once the final state reached, system
can not reverse back to its initial state without the aid of
external agency.
Adiabatic process - a process that has no heat transfer
into or out of the system. It can be considered to be
perfectly insulated.
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14. Explanation of reversible process
Reversible process
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Key features of reversible process:
• Driving and opposing force differ by infinitesimal
amount.
• Process occurs at very slowly & take infinite time to
complete.
• At every step, system attains equilibrium.
• Most of the natural process are irreversible except
phase changes ( melting/freezing of ice/water at 0°C
& 1 atm pressure.

15. Path function and state function
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16. Zeroth Law of Thermodynamics
“ If two bodies are in thermal equilibrium with a third body, there are also in thermal
equilibrium with each other.”
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17. WORK
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18. WORK
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Work done in reversible process for an ideal gas:

19. First law of thermodynamics and internal energy
The first law of thermodynamics states that the amount of heat given to a
system is equal to the sum increase in the internal energy of the system and the
external work done.
Statement of the first law of thermodynamics
 Conservation of energy.
 Energy can neither be created nor destroyed but it can change forms.
 Total amount of energy in a closed system remains constant.
Description with formula
Heat/work formula:
ΔU=Q-W
dU= ΔQ-PdV
(W=PdV)
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20. Significance and limitation
Significance:
 The total amount of energy in the world is constant.
 It is impossible to get more energy out of a system than is put into it.
 You can’t get energy from nothing.
Limitation:
First law does not help to predict whether the certain process is possible or not.
Reason:
It does not specify that heat cannot flow from low temperature body to a high temperature
body.
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21. Enthalpy(H) and Heat capacity (Cp & Cv)
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Enthalpy:
A thermodynamic quantity equivalent to the total heat content of a system is called enthalpy.
Denoted by ‘H’.
H=U+PV; where U is internal energy, P & V are pressure and temperature.
Change in enthalpy;
ΔH=ΔU+ΔPV
ΔH = ΔU+PΔV+VΔP; at constant pressure, ΔH= ΔU+PΔV
Classification of reactions based on ΔH
Exothermic reactions (ΔH = (-) ve) (HCl+NaOH NaCl +H2O+heat)
An exothermic reaction is a reaction that releases energy to the surroundings.
2. Endothermic reactions (ΔH = (+)ve) ;(N2(g)+O2(g) +heat 2NO (g))
An endothermic reaction is a reaction that absorbs energy from the surroundings.

22. Enthalpy(H) and Heat capacity (Cp & Cv)
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Heat Capacity:
It is the amount of heat required to increase the temperature by 1° C.
q = cΔT; the coefficient c is called heat capacity.
Specific Heat Capacity:
The amount of energy needed to increase the temperature of 1 gram of a material by 1 °C is
known as its specific heat.
q = m csΔT; q is heat, m is mass, ΔT is change in temperature, the coefficient cs is called
specific heat capacity.
CP and CV Specific heat capacity at constant pressure and constant volume.
Relation between Cp and Cv:
qP = nCP∆T = ∆H (1); qV = nCV∆T = ∆U (2)
∆H = ∆U + ∆(pV ) = ∆U + ∆(RT) = ∆U + R ∆T (for 1 mole ideal gas; PV=RT)
∆H = ∆U + R ∆T (3)
Using the equn. 1&2 in 3,
we get ; CP∆T = CV∆T + R ∆T
CP –CV =R
CP/CV =γ; γ is 1.6, 1.4 and 1.3 for mono, di and polyatomic gases.

23. Standard Enthalpy(H⦵) and Thermochemistry
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Various enthalpy changes:
1. Standard enthalpy change of reaction, ΔH° r (in general)
2. Standard enthalpy change of formation, ΔH°f
3. Standard enthalpy change of combustion, ΔH°c
4. Standard enthalpy change of neutralisation, ΔH°n
5. Standard enthalpy change of atomisation, ΔH°at
6. Standard enthalpy change of solution, ΔH°sol
7. Standard enthalpy change of hydration, ΔH°hyd
Standard condition: Temparature298K and Pressure1 atm (~100 Kpa)
Common interpretation:
Enthalpy change of 1 mole substance  due to some physical process.