The document outlines the principles and concepts of thermodynamics, including various forms of energy and their interconversions, system classifications, and thermodynamic properties. It describes types of systems (open, closed, isolated), state functions, and equilibrium, as well as processes like isothermal, adiabatic, and reversible. The text also explains the first law of thermodynamics, work and heat exchange, and enthalpy, concluding with relationships related to chemical reactions and energy changes.

1. Thermodynamic

2. Energy
Different form of energy- (interconvert
able)
1. K.E
2. P.E
3. Heat energy (thermal energy)
4. Radiant energy (electromagnetic or
light radication)
5. Electrical energy
6. Chemical energy

3. Thermodynamics
• Thermo- heat and dynamic –
motion.
• Branch of science which deals with
the study of interconversion of
different forms of energy and the
quantitative relationship between
them taking place in physical and
chemical process.

4. Limitations
• Does not give info. Rate of physical or
chemical process.
• Doesn’t describe status, mechanism,
history of the process.
• Only deals with microscopic systems.
• Thermodynamic is the only dynamics
which does not consider time factor.

5. Macroscopic Science
Deals with large no. of atom or molecule not individual atom or molecule.

6. Chemical thermodynamics
Which gives a quantitative
information of the energy change
accompanying chemical process and
explains chemical behavior,
Eg. Heat of reaction, effect of
temperature on chemical reactions,
etc.

7. System, Surrounding and Boundry
• System- the portion of the universe under
thermodynamic consideration to study
thermodynamic properties is called a system.
• Under universe portion is system.
• Here thermodynamics means P, V, T, n, E etc.
• System may be very large or very small.
• System is confined by a real or an imaginary
boundry.
• Human, boil water, animal etc.

8. Surroundings
• The remaining portion of the universe.
• It represent large stock of mass and energy.
• Can exchange energy with system when
allowed.
• Eg. Universe, environment, earth etc.

9. Boundary
• The wall separating the system from its
surrounding.
• Boundary may be real or imaginary.
• Boundary exchange heat, matter, between
system and surrounding.
• Everything outside the boundary is surrounding.
• Eg. Hot water beaker wall of the beaker = real
boundary. While open portion show imaginary
boundry.

10. Types of system
• Open system
• Closed system
• Isolated system
• Homogeneous system
• Heterogeneous system

11. Open system
• System which exchange both matter
and energy with its surroundings.
Eg. Beaker containing water.
• Exchange- Water continuosly
absorbs energy from its surroundings
and from vapour relase it.

12. Closed system
• Exchange energy not matter.
Example- closed vessel containing
hot water so that only heat is lost
not matter.

13. Isolated system
• Can neither exchange energy nor matter.
Example- hot water filled in thermally
insulated closed vessel like thermos flask.
• in actual perfect isolated system is not
possible.
• Universe is an eg. Of isolated system.
• Universe has no boundary, surrounding.

14. Homogeneous system
• Only one phase system.
• Single component system – Zn, O,
Water.
• Solution miscible liquid- water and
alcohol or NaCl and water etc.
• Mixture of gases- H, N, O etc.

15. Heterogeneous system
• Two separated phase by boundary.
• Mixture of immiscible liquid- Water
and benzen.
• Solid identical with liquid- ice and
water.
• Liquid identical with vapour- water
and vapour.

16. Thermodynamic Properties of system
• Variable on which property of the
system depend. For eg. P, T, V, D, E
etc
• The property of the system
classified as
1. Extensive Property
2. Intensive Property

17. Extensive Property
• Whose magnitude depend on the amount of
the matter present in the system.
• When the amount of matter change its
magnitude also change.
• Additive.
Example- enthalpy, mass, volume, energy,
weight.

18. Intensive property
• Magnitude independent of the amount of
matter.
• The ration of extensive property represent an
intensive property. Eg. Density= M/V.
Example- B.P- take either 1ml or 1L B.P of water
is 10 ˚C.
0̊
M.P, F.P, Surface tension, specific heat, molar
heat capacity, T, P, D, viscosity.

19. State and state function
State variable-
• Measurable property of a system like P, T, V
etc.
• State describe value of these variable.
• When one or more variable change system
change to new state.
• Macroscopic properties of a system depend
on these state variable.

20. State function
• Which depend on initial and final state but
independent of the path followed by the
system during the process.
• Eg. Mass (initial to final product), P, T, V etc.
• It depend on state.
• Diagram shown by board.
• It shows changes are independent of all three
path but depend on initial and final state.

21. Thermodynamic Equilibrium
• No change in any thermodynamic function or
state function like energy, pressure etc with
time.
• Type-
1. Thermal equilibrium
2. Chemical equilibrium
3. Mechanical equilibrium

22. Thermal equilibrium
• System and surrounding at same
temperature and no exchange of heat.
• Total energy remain const. eg. Water
with its vapour at constant temperature.

23. Chemical equilibrium
• Chemical composition doesnot
change with time.
• Eg.
N + 3H2 → 2NH3
Composition of Reactant and product
does not change with time.

24. Mechanical equilibrium
• No moment of matter in system
with respect to its surrounding.
• Mechanical property remain
constant.

25. Thermodynamic process
• An operation or transition by which a state of a
system changes from initial state to final state.
Type of process-
1. Isothermal process(∆T=0)
2. Isobaric process(∆P=0)
3. Isochoric process(∆V=0)
4. Adiabatic process(q=0)
5. Reversible process
6. Irreversible process

26. Isothermal process
Temperature of the system constant. ∆T =0
1. In this process temperature at initial state and
final state is constant.
2. In this process system exchange heat energy
with its surrounding to maintain constant
temperature.
3. Occur in close system.
4. Internal energy of the system remains constant,
hence ∆U = 0.
5. In this process, gaseous system P, V of a change.

27. Reversible process
• A process carried out in such a manner that
every stage, the driving force is only
infinitesimally greater then the opposing
force and it can be reversed by an
infinitesimal increase in opposing force and
the system exists in equilibrium with its
surrounding throughout, is called a
reversible process.
• Slow, infinite number of step.

28. Irreversible process
• Unidirectional process which proceeds
in a definite direction and cannot be
reversed at any stage and in which
driving force and opposing force differ in
a large magnitude.
• Also called spontaneous process.
• They are real process not hypothetical.
• Eg. Flow of heat from high T to lower T.

29. Work
• W = F.s
• Work is one of the ways by which a
system can exchange energy with its
surrounding by changing the state of the
system.
Example- Object move by applying
force( object energy)

30. Kind of work in chemical thermodynamics
• The type of work is mechanical work i.e.
pressure volume work.
• W = - Pex (V2 – V1)
external pressure apply change in volume
• Work is also obtain due to chemical process
or reaction.

31. Expression for pressure- volume work
Ideal gas
Massless,
frictionless
Piston.
As the gas
Expand it pushes
Piston upward through
Distance d against external force.
d
Pex
P
W = - F x d
a
Area of cross section

32. • If ‘a’ is the cross section area of the
cylinder or piston, then
W = x d x a
Now the pressure is Pex = F/a and ∆V = d x
a.
W = - Pex x ∆V
Expression for pressure- volume work

33. Sign convention of work during expansion
and compression
A. Expansion of a gas:
Pex changing the volume from V1 to V2.
then ∆V = V2 – V1.
W = - Pex x ∆V.
• During expansion V2 > V1. work perform by
surrounding. This result decrease energy of
the system.
• Hence work is –ve i.e W is –ve.

34. Sign convention of work during expansion
and compression
A. Compression of a gas:
Pex changing the volume from V1 to V2.
then ∆V = V2 – V1.
W = + Pex x ∆V.
• During expansion V2 < V1. work perform by
system. This result increase energy of the
system.
• Hence work is +ve i.e W is +ve.

35. Concept of Maximum work
the process carried out at a constant
temperature in the reversible manner by
changing the state of the system through
infinitesimally small steps in which driving
force is infinitesimally greater then
opposing force give maximum work. Is
called an isothermal reversible process.

36. • Process carried out at constant temperature.
• During the complete process, driving force is
infinitesimally greater then opposing force.
• The work obtained is maximum. This is given
Wmax = -2.303 nRT log10 or log10.
• ∆U = 0.
• The heat absorbed irreversible manner qrev , is
completely converted into work.
Concept of Maximum work

37. Condition of Maximum work
In a thermodynamic process, maximum work is
obtained from a system when
• All the changes taking places in it are
thermodynamically reversible.
• Change in the state of the system take place in
infinite no. of step.
• During change, driving force is infinitesimally
greater then the opposing force.

38. Expression for Maximum work
Ideal gas
Massless,
frictionless
Piston.
As the gas
Expand it pushes
Piston upward through
Distance d against external force.
P – dP
W = - F x d
V
V+dV
P
dV
V + dV – V = dV

39. Expression for Maximum work
dW = -(P- dP) dV
dW = -PdV – dPdV
dPdV = negligible.
dW = -PdV
Wmax = - 2.303 nRT loge 𝐎𝐫
𝐏𝟏
𝐏𝟐

40. Path dependence nature of work
• Work is not the property of the system.
• Not state funtion.
W = - P (V2-V1)
A
(V1) B
(V2)
W1 Expand Path
1
W2 Pex
2
W3 exapnsion reversibly 3

41. Concept of Heat
• Another way of exchanging energy
system and surrounding.
• Not property of system. Not state
funtion.
• Heat exchange only possible by path.
• Heat is path dependence.
• Eg. Rod heat transfer

42. Sign Convention of W and q
• Work and Heat are the form of energy.
• Due to work and exchange of heat, the
energy of the system changes.
• +q = +W eg surrounding to system.
• -q = -W eg Gym.
• +q = heat absorbed and –q = heat
released.
• +W = compressed.
• -W = expansion.

43. Unit of energy and Work
• Litre- atmosphere(L atm OR
lit.atm)
• Erg- W= dyne x cm = 1 erg.
Force & distance.
• Calorie- heat energy
• Joule- amount of work.

44. Interconversion of work and energy
• W = 1 atm x 1 lit
• 1 atm = 1.013 x 105
Nm2
• 1 lit = 10-3
m3
• W = (1.013 x 105
x 10-3
)Nm
• W = 101.3 J
• W = 24.22 cal
• W = 1.013 x 109
erg

45. Internal Energy[U]
• Total energy K.E and P.E present in the
system.
• State function.
• Value depends on the state of a system.
• Change in internal energy, ∆U = U2 – U1.
• Extensive property.
• Same unit as work and energy.

46. Total energy [U] = P.E + K.E.
U total = U = U potential + U Kinetic
U potential = Uintramolecular + Uintermolecular
U Kinetic = Utranslational + Uvibrational +
Urotational + Uelectric
U = Uintra+ Uinter+ Utrans+ Uvib+ Urota + Uelectric
Internal Energy[U]

47. First law of thermodynamics
• Law of conservation of energy.
• 1 kind of energy consumed another kind of
energy disappears.
• It is impossible to construct a perpetual
motion machine.
U = q + W
Total amount of work is converted into heat
energy.

48. Mathematical equation of 1st
Law of
thermodynamic
(V2, U2)Final state Initial state(V1, U1)
Heat absorb
from
surrounding

49. • Due to volume change, the system perform the
work W, hence total energy U2 of the system in the
final state is,
• U2 = U1 + q + W
• U2 – U1 = q + W
• ∆U = q + W
• For infinitesimally small change the mathematical
expression is,
• dU = dq + dW
Mathematical equation of 1st
Law of
thermodynamic

50. First law of thermodynamic for various
processes
• Isothermal process:- ∆T= 0
System depends on the temperature there is no
change in the internal energy U of the system.
Hence ∆U = 0.
∆U = q + W
0 = q + W
+q(expansion) = -W or
W = -q(consumed).

51. First law of thermodynamic for various
processes
• Isobaric process:- ∆P = 0
System performs the work of expansion due to
volume change . W = -Pex x ∆V
qP heat absorb by the system at constant pressure.
∆U = qP + W
∆U = qP – Pex ∆V
Or qp = ∆U + Pex∆V
qp heat absorbed used to increase the internal energy of the
system.

52. Isochoric process
∆V = 0
Hence system doesn’t perform mechanical
work..
• W = - P∆V = 0.
• ∆U = q + W
• ∆U = qv
• qv = heat absorbed at constant volume.
• ∆U and q is state funtion.

53. Adiabatic process
• q= 0.
• ∆U = q + W
• ∆U = Wad
• System Expansion - ∆U decrease internal
energy and temperature of system decrease.
• System Compression - ∆U increase internal
energy and temperature of system increase.

54. Modern form of the first law of
thermodynamic
• According to Einstein's theory, mass can be
converted in to energy.
• Hence mass is also form of energy.
• The sum of mass and energy of an isolated
system remain constant.

55. IUPAC sign convention of q, U and W
• For heat q:-
+q = heat absorb by system.
-q = heat loss by system. Heat energy left.
• For work W:-
+W = work done on the system by compression.
-W = expansion. Internal energy of system lose.
• For internal energy U:-
+U = internal energy of system increase by
absorption of heat. Similarly, –U.

56. Enthalpy
H = U + PV
 Enthalpy represent total heat content of
the system, at constant pressure.
 State function and extensive property.
 Absorption of heat by system increase its
enthalpy.
 Hence enthalpy is called heat content of
the system.

57. Expression of enthalpy change
H1, U1,P1,V1 H2,U2,P2,V2
H1 = U1 + P1V1 & H2 = U2 + P2V2
The enthalpy change ∆H is given by,
∆H = H2 – H1
∆H = U2 + P2V2 – (U1 + P1V1)
∆H = U2 – U1 + P2V2 – P1V1
∆H = ∆U + P∆V

58. Show that the heat absorbed at constant
pressure is equal to the change in enthalpy of
the system.
• By the first law of thermodynamic,
• ∆U = q + W
• q= ∆U – W
• If qp= heat absorbed at const. P.
• W = -P∆V
• qp = ∆U + P∆V. (∆H=∆U + P∆V)
• qp = ∆H
• Enthlpy also called heat content of the system.

59. Relation between ∆U & ∆H in
1. Isochoric process:-
• ∆H = ∆U + P∆V
• ∆V = 0.
• ∆H = ∆U
2. Isobaric process:-
• ∆P =0
• ∆H = ∆U

60. Derive the expression for the heat of
reaction at
1. Constant pressure-
• q= ∆U – W.
• ∆Hp = ∆U – W (W = - P∆V)
• ∆Hp = ∆U + P∆V
• qp= ∆Hp
2. Constant Volume-
• q= ∆U

61. Derive the relation ∆H=∆U+∆nRT
• Consider a reaction in which n1 moles of gaseous reactant
in initial state change to n2 moles of gaseous product in
the final state.
• n1A (H1U1P1V1) n2B(H2U2P2V2)
• Enthalpy change, ∆H= H2 - H1.
• H1 = U1 + P1V1 & H2 = U2 + P2V2
• The enthalpy change ∆H is given by,
• ∆H = H2 – H1
• ∆H = U2 + P2V2 – (U1 + P1V1)
• ∆H = U2 – U1 + P2V2 – P1V1 PV = n1RT

62. • ∆H = ∆U + ∆nRT
• If qp and qv are the heats involved in the
reaction at constant pressure and volume,
then qp = ∆H and qv = ∆U.
• qp= qv + ∆nRT
Derive the relation ∆H=∆U+∆nRT

63. Expression for work done in a chemical
reaction
• Chemical reaction depend upon the change in
no. of gaseous moles of product to reactant.
• Consider, n1
V1A n2
V2B.
• Initial state, PV1 = n1RT
• Final state, PV2 = n2RT
• W = -P∆V = -∆nRT.
• If n1=n2, W=0.
• If n2>n1, expansion W-ve
• If n2<n1, compression W+ve.

64. Enthalpies of physical change
• Phase Transitions:- change in physical state of
matter. Type of phase changes.
a. Fusion- solid to liquid state. Heat absorbed,
endothermic(∆H>0). Eg. Ice to water.
b. Vaporisation or evaporation- liquid to gas state.
Heat absorbed, endothermic(∆H>0)
c. Sublimation- solid to gas state. ∆H>0. eg
Camphor .
Note- Temperature and Pressure remain constant.

65. (1) Phase Transition
= 40.7kJmol-1
At ˚C
0̊
= 6.01kJmol-1
At ˚C
0̊
= 40.7kJmol-1
At 10 ˚C
0̊

66. (2) Enthalpy of atomic or molecular changes
• A. enthalpy of ionization.

67. B. enthalpy of atomization-
• Dissociation of 1 mole gas substance into free
gaseous atom.
(2) Enthalpy of atomic or molecular changes

68. When one mole of a substance is dissolved
in a large excess of a solvent, so that
further dilution will not change the
enthalpy at const. T and P.
Eg. HCl(g) + aq → HCl(aq) ∆H = 4kJmol-1
(C) Enthalpy of Solution

69. (D) Enthalpy if dilution
• Solution – concentrated solution is
diluted to form another concentrated
solution.eg
HCl + 50 H2O → HCl(50H2O), ∆H= -73.26kJ

70. Thermochemistry
• Study of heat change during chemical
reaction.
• Heat of reaction:-(enthalpy of chemical
reaction)
• ∆H = ∑Hproduct - ∑Hreactant
• Endothermic reaction- ∑Hproduct > ∑Hreactant
absorption of heat, ∆H +ve.
• Exothermic reaction- ∑Hproduct - ∑Hreactant
lose of heat, ∆H –ve.

71. Thermochemical equation
• A chemical reaction which is represent by
1. Physical state of of all reactant and product.
2. Reactant reacting to form product.
3. A balanced equation
4. Enthalpy change(heat given in or out during
reaction)
A + B → C + D ∆H= (+,- ) kJmol-1

72. IUPAC Guideline for writing thermochemical
equations
• Physical state, balance chemical equation.
• Heat and enthalpy changes are measured at STP.
298K & 1atm.
• ∆H written on R.H.S.
• Proper sign must be indicate +H, -H.
• Enthalpy of the element in their STD states is
taken as zero. STD shown by H˚.
• Allotropic form must mentioned.eg Cgraphite.
• For reverse rxn, ∆H value same but sign change.

73. Standard Enthalpy of reaction
• The difference between the sum of enthalpy
of products and reactant with every substance
in its standard state at constant T(298K) and
P(1atm).
• ∆H˚ = ∑H˚product - ∑H˚reactant
• Reactant → product

74. Standard enthalpy of formation or standard
heat of formation(∆fH˚)