Basic Concepts of Thermodynamics

1. Course: Diploma
Subject: Applied science(Chemistry)
Unit: V

2. •It is the study of change in any form of heat or energy
during any process.
•The study of the flow of heat or any other form of
energy into or out of a system as it undergoes a
physical or chemical transformation, is called
Thermodynamics.

3. System:
A quantity of matter or a region in universe which
is chosen for thermodynamic study is called system.
Surroundings:
The mass or region outside the system or the rest is
Surroundings.
Boundary:
The real or imaginary surface that separates the
system from its surroundings.

4. SYSTEM
SURROUNDINGS
BOUNDARY

5. A phase is defined as a homogeneous,
physically distinct and mechanically separable portion of a
system.
Homogeneous System :
When a system is uniform throughout,
it is called a Homogeneous System.
•A homogeneous system is made of one phase only.
•Examples are :
•A pure single solid, liquid or gas, mixtures of gases, and
true solution of a solid in a liquid

6. Heterogeneous system
A heterogeneous system is one
which consists of two or more phases.
•In other words it is not uniform throughout.
•Examples of heterogeneous systems are :
•ice in contact with water,
•ice in contact with vapour
•Here ice, water and vapour constitute separate phases.

8. In such a system the boundary is open and
un-insulated.
An open system is one
which can transfer both energy and
matter to and from its surroundings.
Hot water contained in a beaker placed on
laboratory table is an open system.
The water vapour (matter) and also heat
(energy) is transferred to the surroundings
through the imaginary boundary.
OPEN
system
Mass YES
Energy YES

9. Here the boundary is sealed but not insulated.
A closed system is one which
cannot transfer matter but can transfer
energy in the form of heat, work and
radiation to and from its surroundings.
A specific quantity of hot water contained in a
sealed tube, is an example of a closed system.
While no water vapour can escape from this
system, it can transfer heat through the walls of
the tube to the surroundings.
Mass YES
Energy NO

10. An isolated system is one that can transfer neither
matter nor energy to and from its surroundings.
Let us consider a system 100 ml of water in contact with its
vapour in a closed vessel which is insulated. Since the vessel is
sealed, no water vapour (matter) and heat can escape from it.
A substance, say boiling water, contained in a thermos flask, is
another example of an isolated
system.

13. A property which does not depend on the
quantity of matter present in the system, is
known as Intensive Property.
intensive properties are
pressure
temperature
Density
concentration
 If the overall temperature of a glass of water (our
system) is 20ºC, then any drop of water in that glass
has a temperature of 20ºC.

14. •A property that does depend on the quantity of matter
present in the system, is called an Extensive Property.
•extensive properties are
•volume
•number of moles
•enthalpy
•entropy
•consider the system ‘a glass of water’. If we double the mass of
water, the volume is doubled and so is the number of moles and
the internal energy of the system.

15. •Law of conservation of energy
• energy cannot be created nor destroyed.
•Therefore, the total energy of the universe is a
constant.
•Energy can, however, be converted from one
form to another or transferred from a system to
the surroundings or vice versa.

16. The total energy of an isolated system remains
constant though it may change from one form to
another.
When a system is changed from state A to state B, it undergoes a
change in the internal energy from EA to EB.
Thus, we can write
ΔE = EB – EA
This energy change is brought about by the evolution or
absorption of heat and/or by work being done by the system.

17. As the total energy of the system must remain constant,
we can write the mathematical statement of the
First Law as :
ΔE = q – w ………....(1)
where q = the amount of heat supplied to the system
w = work done by the system
the net energy change of a closed system
is equal to the heat transferred to the system
minus the work done by the system.

18. •(1) Whenever energy of a particular type disappears
equivalent amount of another type must be produced.
•(2) Total energy of a system and surroundings remains
constant (or conserved)
•(3) It is impossible to construct a perpetual motion
machine that can produce work without spending
energy on it.

19. ENERGY
19
Macroscopic energy Microscopic energy
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.
Those related to motion
and the influence of some
external effects such as
gravity, magnetism,
electricity and surface
tension.
Internal energy, U: The
sum of all the microscopic
forms of energy.
Total energy
of a system

20. •A thermodynamic system containing some quantity of
matter has within itself a definite quantity of energy.
•This energy includes not only the translation kinetic
energy of the molecules but also other molecular energies
such as rotational, vibrational energies.
• The kinetic and potential energy of the nuclei and
electrons within the individual molecules also contribute
to the energy of the system.

21. The total of all the possible kinds of energy of
a system, is called its Internal Energy.
The internal energy of a system, like temperature, pressure,
volume, etc., is determined by the state of a system and is
independent of the path by which it is obtained.
Hence internal energy of a system is a state
function.
The SI unit for internal energy of a system is the joule (J).
Another unit of energy which is not an SI unit is the calorie.
1 cal = 4.184 J.

22. •For example
•we consider the heating of one mole of liquid water
from 0º to 100º C.
•The change in energy is 1.8 kcal and is the same
regardless of the form in which this energy is
transferred to the water by heating.
•Since the value of internal energy of a system depends
on the mass of the matter contained in a system, it is
classed as an extensive property.

23. •At constant volume (say in a sealed tube), the heat content of a
system is the same as internal energy (E), as no PV work is done.
• But in a constant-pressure process, the system (a gas) also
expends energy in doing PV work.
• The total heat content of a system at
constant pressure is equivalent to the internal energy E
plus the PV energy. This is called the Enthalpy (Greek en
= in; thalpos = heat) of the system and is represented by
the symbol H.
•Thus
enthalpy is defined by the equation :
H = E + PV ...(1)

24. In the equation (1) above, E, P, V are all state functions.
Thus H, the value of which depends on the values of E, P, V must also
be a function of state.
Change in Enthalpy
If Δ H be the difference of enthalpy of a system in the final state (H2)
and that in the initial state(H1),
ΔH = H2 – H1 ...(2)
Substituting the values of H2 and H1, as from (1) and (2), we have
ΔH = (E2 + P2V2) – (E1 + P1V1)
= (E2 – E1) + (P2V2 – P1V1)
= ΔE + ΔPV
If P is constant while the gas is expanding, we can write
ΔH = ΔE + PΔV
or ΔH = ΔE + w (w = work) ...(3)

25. According to the First Law,
ΔE = q – w ...(4)
where q = heat transferred
From equations (3) and (4)
ΔH = q
when change in state occurs at constant pressure
This relationship is usually written as
ΔH = qp
where subscript p means constant pressure.
Thus Δ H can be measured by measuring the heat of a
process occurring at constant pressure.

26. Entropy is a measure of molecular disorder, or molecular
randomness.
As a system becomes more disordered, the positions of
the molecules becomes less predictable and the entropy
increases.
Definition:
Entropy is a thermodynamic state quantity that is a
measure of the randomness or disorder of
the molecules of the system.

27. •The symbol of entropy is S
•while the change is represented by ΔS.
•The entropy of a system is a state function and
depends only on the initial and final states of the
system.
•The change in entropy, ΔS, for any process is given by
the equation,
ΔS = Sfinal – Sinitial
• When Sfinal > Sinitial, ΔS is positive.

28. •In 1850 Clausius introduced a numerical definition of
entropy. According to him entropy of a system is a
constant quantity when there is no communication of
heat.
•When heat (q) flows into a system, the entropy increases
by q/T . Heat flowing out of a system produces a
corresponding decrease.

29. •Thus entropy could be precisely defined
as : for a reversible change taking place at a fixed temperature
(T), the change in entropy (ΔS) is equal to heat energy absorbed
or evolved divided by the temperature (T).
If heat is absorbed, then ΔS is positive and there will be increase in
entropy. If heat is evolved, ΔS is negative and there is a decrease in
entropy.
entropy is equal to heat energy divided by absolute temperature.
Therefore, it is measured in entropy units (‘eu’) which are calories per
degree per mole
cal mol–1 K–1.
In the SI system, the units are joules per mole per degree i.e., J mol–1 K–
1. These are represented by eu.
1eu = 4.184

30. Implications:
• more particles
-> more states -> more entropy
• higher T
-> more energy states -> more entropy
• less structure (gas vs solid)
-> more states -> more entropy

31. The heat capacity of a defined system is the amount
of heat (usually expressed in calories, kilocalories, or joules)
needed to raise the system's temperature by one degree
(usually expressed in Celsius or Kelvin).
 It is expressed in units of thermal energy per degree
temperature.
To aid in the analysis of systems having certain specific
dimensions, molar heat capacity and specific heat
capacity can be used.
To measure the heat capacity of a reaction, a calorimeter
must be used. Bomb calorimeters are used for constant
volume heat capacities.

32. The amount of heat needed to increase the
temperature of one mole of a substance by one
degree(K OR °C) is the Molar heat capacity.
It is expressed in joules per moles per degrees
Celsius (or Kelvin), Joules/Moles/∘C.
For example, the molar heat capacity of lead is
26.65 Joules/Moles/K, which means that it takes
26.65 Joules of heat to raise 1 mole of lead by °K.

33. The amount of heat needed to increase the
temperature of one gram of a substance by one
degree is the specific heat capacity.
It is expressed in joules per gram per degree
Celsius, Joules/Grams/∘C.
As the specific heat of lead is 0.128
Joules/Grams/C, it takes 0.128 Joules to raise
one gram of lead by one °C.

34. If q calories is the heat absorbed by mass m and the
temperature rises from T1 to T2, the heat capacity (c) is
given by the expression.
Thus heat capacity of a system is the heat absorbed by unit
mass in raising the temperature by one degree (K or º C) at
a specified temperature.
When mass considered is 1 mole, the expression (1) can be
written as
Molar Heat Capacities

35. The molar heat capacity of a system is defined as
the amount of heat required to raise temperature of
one mole of the substance (system) by 1 K.
Since the heat capacity (C) varies with temperature; its true
value will be given as,
where dq is a small quantity of heat absorbed by the system,
producing a small temperature rise dT.

36. Molar Heat Capacity at Constant Volume
According to the first law of thermodynamics
dq = dE + PdV ..(1)
Dividing both sides by dT, we have
At constant volume dV = 0, the equation reduces to
…..(2)

37. Molar Heat Capacity at Constant Pressure
Equation (2) above may be written as
We know H = E + PV
Differentiating this equation w.r.t T at constant
pressure, we get
comparing it with equation (3) we have
….(3)
…(4)

38. During constant Pressure, part of energy is used to increase the
internal energy of the system ,while during constant volume no
work done.
Therefore Cp is greater than Cv.
By definition H = E + PV for 1 mole of an ideal gas
H = E + RT ( PV = RT)
Differentiating w.r.t. temperature, T, we get
…(1) …(2)

39. Cp = Cv + R
or Cp – Cv = R
[By using equations (1) and (2)]

40. When a thermodynamic system changes from
one state to another, the operation is called a Process.
These processes involve the change of conditions
(temperature, pressure and volume).
Adiabatic
Process
Isochoric
process
Isothermal
process
Isobaric
process

41. Those processes in which no heat can flow into
or out of the system, are called adiabatic
processes.
Adiabatic conditions can be approached by
carrying the process in an insulated container.
For an example, ‘thermos’ bottle. High vacuum and
highly polished surfaces help to achieve thermal
insulation.
For an adiabatic process dq = 0

42. Those processes which take place at constant
pressure are called isobaric processes.
For an example, heating of water to its boiling point and
its vaporisation take place at the same atmospheric
pressure.
These changes are, therefore, designated as isobaric
processes and are said to take place isobarically.
For an isobaric process dp = 0

43. Those processes in which the volume
remains constant are known as isochoric
processes.
The heating of a substance in a non-expanding
chamber is an example of isochoric process.
For isochoric processes dV = 0

44. Those processes in which the
temperature remains fixed, are termed
isothermal processes.
This is often achieved by placing the system in a
thermostat (a constant temperature bath).
For an isothermal process dT = 0

45. Reversible process
A thermodynamic reversible process is one that
takes place infinitesimally slowly and its direction
at any point can be reversed by an infinitesimal
change in the state of the system.
Irreversible process
When a process goes from the initial to the final
state in a single step and cannot be carried in
the reverse order, it is said to be an irreversible
process.

46. (a)Reversible process (b)Irreversible process

47. Consider a certain quantity of a gas contained in a cylinder having
a weightless and frictionless piston. The expansion of the gas can
be carried by two methods illustrated.
Reversible process
•Let the pressure applied to the piston be P
•the pressure on the piston is decreased by an infinitesimal
amount dP.
•Thus the external pressure on the piston being P – dP
•If at any point of the process the pressure is increased by dP, the
gas would behave reversibly.
•Irreversible process
•If the pressure on the piston is decreased suddenly.
• It moves upward rapidly in a single operation. The gas is in
equilibrium state in the initial and final stages only.
• The expansion of the gas takes place in an irreversible manner.

48. Reversible Process Irreversible Process
1. All changes are reversed when the
process is carried out in reversible
direction.
1. After this type of process has occurred
all changes do not return to the initial
stage by themselves.
2. It is imaginary as it assumes the
presence of frictionless and weight less
piston.
2. It is real and can be performed
actually.
3. It is in equilibrium state at all stage of
the operation.
It is in equilibrium state only at the
initial and final stage of the operation.
4. It takes place in infinite number of
infinitesimally small steps and it would
take infinite time to occur.
4. It takes place infinite time.
5. It is extremely slow. 5. It proceeds at measureable speed.
6. Work done by a reversible process is
greater than the corresponding
irreversible process.
6. Work done by a irreversible process is
smaller than the corresponding
reversible process.

49. (A) Spontaneous process
A process which proceeds of its own accord, without any
outside assistance, is termed a spontaneous or natural
process. ex. Mountain Climbing, Gas Flow.
(B)Nonspontaneous process
The reverse process which does not proceed on its own, is
referred to as a nonspontaneous or
unnatural process.
Entropy
Entropy is a thermodynamic state quantity that is a
measure of the randomness or disorder of
the molecules of the system.

50. whenever a spontaneous process takes place, it
is accompanied by an increase in the total energy of the
universe.
ΔSuniv = ΔSsyst + Δssurr
when an irreversible spontaneous process occurs,the entropy of the
system and the surroundings increases.
In other words ΔSuniv > 0.
When a reversible process occurs, the entropy of the system remains
constant.
ΔSuniv = 0.
Since the entire universe is undergoing spontaneous change,
stated as : the entropy of the system is constantly increasing

52. At absolute zero temperature, the entropy of
a pure crystal is zero.
. That is, S = 0 at T = 0 K
 The entropy of a substance varies directly with temperature.
The lower the temperature, the lower the entropy.
For example, water above 100ºC at one atmosphere exists as a
gas and has higher entropy(higher disorder).
The water molecules are free to roam about in the entire
container.
When the system is cooled, the water vapour condenses to form
a liquid.

53. If we cool the solid crystal still further, the vibration
of molecules held in the crystal lattice gets slower and
they have very little freedom of movement (very little
disorder) and hence very small entropy.
Finally, at absolute zero all molecular vibration ceases
and water molecules are in perfect order. Now the
entropy of the system will be zero.

54. Molecular states in Solid crystal

55. References
1. Elements of Physical chemistry by Glass ton &
Lewis
2. Essentials of Physical chemistry by Bahl & Tuli