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 The branch of science which deals with
the study of different forms of energy
and their interconversion is called
thermodynamics.
System & Surroundings
System
Surroundings
 A specified part of the
 The remaining portion of
universe which is under
observation is called
system.
surroundings
System
the universe which is not
a part of the system is
called surroundings.
Types of System:
Open, closed and isolated system
All physical and chemical processes taking place in open in our daily life
are open systems because these are continuously exchanging mass and
energy with the surroundings.
Surroundings
Matter
Matter
System
Energy
Matter
Energy
Energy
The state of the system & State
Variable
 State function: Those properties which
determines a particular system is called
state function.
 Path function: Path function depends on
the path of change from initial state to
final state.
 State function: The measureable
properties required to describe the state
of the system are called state variables.
Internal energy and internal energy
change
Internal energy (U)
• The energy associated with the system at a
particular conditions of temperature and pressure.
Internal energy change (∆U)
• It is measure of heat change occurring during the
process at constant temperature & volume.
•
qv = ∆U
Internal energy as a state function
In thermodynamics
internal energy(U)
of the system may
change , when
Heat passes into or
out of the system
Work is done on or
by the system
Matter enters or
leave the system
 1. Heat(Q): energy is exchanged
between the system and the
surroundings as heat if they are at
different temperatures. ∆U = q
 2. Work(w): It is also a mode of
transference between system &
surroundings. Work done by the
system on the surroundings is given
by p∆V
Sign conventions
 Work :
a) positive , if work is done on the
system.
b) negative, if work is done by the
system.
 Heat :
a) positive , if heat absorbed by the
system.
b) negative , if heat released by the
system.
Types of changes
(a) Isothermal change:
A change in which the
remains constant.(∆T= 0)
(a) Adiabatic change:
A change in which the
constant.(∆q = 0)
(a) Isobaric change:
A change in which the
remains constant.(∆P = 0)
(a) Isochoric change:
A change in which the
constant.(∆V = 0)
temperature of the system
heat of the system remains
pressure of the system
volume of the system remains
Law of Conservation of Energy : First Law
of Thermodynamics
It states that the
energy of the
universe remains
constant during
chemical & physical
changes.
Or, the energy of an
isolated system
constant.
 Mathematical expression for first
law of thermodynamics:
initial internal energy = U1,
heat added = q + w
final internal energy (U2)= U1+q+w
Now,
U2 = U1 + q + w
U2 - U1 = q + w
∆U
Change
in
internal
energy
Heat
added to
the
system
q
w
Work
done on
the
system
Work
Irreversible work
Reversible work
 Change should take place
 Change should take place
is fast.
 In each step the system
&surroundings are not
always in eqilibrium.

At any stage the system
cannot be infinitesimal
change.
 W
irrev
= -p∆V
in infinitesimal change in
small steps.
 In each step the system &
surroundings are always in
near equilibrium with each
other.
 At any stage the system
can be infinitesimal
change.
 W rev = 2.303nRT log V2/V1
Work
 Pressure – Volume work
(a) Work done on the system
If, pex > pin
piston moves downwards until pex becomes pin.
pex
F =p/A
F = pex × A
Length(l)
W = (pex × A) × l = pex ×(A ×l)
W = Pex × ∆V
pin
Also, Vi = initial volume, Vf = final volume
∆V = Vf – Vi
w = -pex × ∆V
w= -p ∆ V
Work
 (b) Work done by the system
If, pin > pex
then, Piston will moves upwards.
F = pin × A × l
= pin ∆V
w=F×l
= (pin × A) × l
= pin ∆ V = pin(-∆ V)
w = -pin∆V
pex
Length(l)
pin
Isothermal & free expansion of an
ideal gas
∆U = q +w can be expressed for isothermal irreversible
& reversible change as follows;
 For isothermal irreversible change
q = -w = Pex ( Vf – Vi)
 For isothermal reversible change
q = -w = nRT ln Vf/Vi
= 2.303 nRT log Vf/Vi
 For adiabatic change, q = 0
∆U = wad
Enthalpy & Enthalpy Change
(1) Enthalpy is sum of
internal energy &
pressure-volume
energy at a particular
temperature &
pressure. It is also
called heat content.
H = E + pV
H denotes enthalpy.
(2) Enthalpy change is
the measure of heat
change taking during
constant temperature
and constant pressure
. It is denoted by ∆H.
∆U = q + w irrev
=q – p ∆V
…(1)
at constant
volume(∆V=0), change in internal
energy is equal to change in heat at
constant volume.
From 1st law of thermodynamics,
∆U = q – p ∆V
U2-U1 = qp - p(V2-V1)
(U2+pV2) – (U1+pV1) = qp
(final state)
…(2)
(initial state)
Let,
U + pV = H
enthalpy)
So,
U2 + pV= H2
Hence, from (2)
H2-H1 = qp
(H =
&
U1+ pV = U1
∆H =qp
So, enthalpy is heat change at pressure
Relationship between ∆H & ∆U (1)
We know from 1st law of thermodynamics,
∆U = qp – p∆V
qp = ∆U + p∆V
∆H = ∆U+ p∆V
…(1)
(∆H = qp)
If the total volume of gaseous reactant is VA and
total no. of moles of gaseous reactant is nA.
After reaction the total no. of moles and
volume of gaseous product is nB and Vb
respectively.
aA + bB
cC + dD
Then , by the ideal gas equation;
Relationship between ∆H & ∆U (2)
Total vol. of gaseous reactant =VA
Total no. of moles of gaseous reactant = nA
Total vol. of gaseous product =VB
Total no. of moles of gaseous product = nB
For reactants (initial state) ;
pV = nART
…(2)
For products (final state) ;
pV= nBRT
…(3)
Relationship between ∆H & ∆U (3)
Subtracting (3) from (2)
P(VB-VA )= (nB-nA) RT
p∆V = ∆ngRT
…(4)
Substitute (4) in (1)
∆H=∆U+∆ngRT
where,
∆H = change in enthalpy
∆U = change in internal energy
∆ng = change in no. of moles of gaseous
reactants & products.
Macroscopic Properties of the System
 The properties of the system which arise from the bulk
behavior of matter are called macroscopic properties.
 In thermodynamics the macroscopic properties can be
divided into two types :
Extensive
properties
• The properties which depends
upon quantity of matter present in
system.
• e.g. mass, volume etc.
Intensive
properties
• The properties which does not
depend upon the quantity of
matter present in the system.
• e.g. temperature, viscosity etc.
 Heat capacity:
The amount of heat required to raise the temperature of
the substance by 1C is known as its heat capacity.
q = C∆T
(q=heat change, C=heat capacity)
 Specific Heat capacity:
The amount of heat required to raise the temperature of unit of
substance by 1C is known as specific heat capacity.
q = mc∆T
(mc=specific heat capacity)
 Molar Heat capacity:
The amount of heat required to raise the temperature of substance
by 1C is known as molar heat capacity.
cm =(C / n)
(cm=molar heat capacity)
Relationship between Cp and Cv for
an ideal gas (1)
Cp = specific heat capacity at constant pressure
Cv = specific heat capacity at constant volume
R = universal gas constant
We know,
q = C∆T
At constant pressure
qp = Cp∆T
∆H = Cp∆T
At constant volume
qv = Cv∆T
∆U = Cv∆T
…(1)
…(2)
Relationship between Cp and Cv for
an ideal gas (2)
We know,
∆H = ∆U + p∆V
…(3)
By substitute (1)&(2) in (3)
Cp∆T = Cv∆T + p∆V
From ideal gas equation,
pV = RT
p∆V = R∆T
…(4)
(for one mole gas)
…(5)
Now , substitute (5) from (4)
Cp∆T = Cv∆T + R∆T
Enthalpy of reaction & standard
enthalpy of reaction
 The enthalpy change taking place during a chemical
reaction is known as enthalpy of reaction.
rH = enthalpy of reaction.
 The enthalpy change taking place during a chemical
reaction when all reactants and products are at
standard state at that temperature is known as
standard enthalpy of reaction.
∆rH° = ∑νP HP - ∑νR HR
Sum of
enthalpy of
products
Sum of
enthalpy of
reactants
Standard state
 Standard state of any substance at any
temperature is its most stable state at
that temperature.
∆H = H2-H1
∆rH = ∑νP HP - ∑νR HR
νP = stoichiometric coefficient of products
νR = stoichiometric coefficient of reactant s
Enthalpy changes during phase
transformation
 Enthalpy of fusion (∆fusH°):
It is enthalpy change taking place during
the fusion of 1 mol of solid at its melting
point.
 E (nthalpy of vaporization ∆vapH°):
It is enthalpy change taking place during
the vaporization of 1 mol of a liquid at its
boiling point.
Enthalpy changes during phase
transformation
 Enthalpy of sublimation (∆subH°):
It is enthalpy change when one mole of a
solid substance sublimes at a constant
temperature and under standard pressure
(1bar).
Standard enthalpy of formation
(∆fH°)
 The enthalpy change taking place when
one mole of compound is formed from its
constituents elements and all are in their
standard state.
∆rH° =[sum of standard enthalpy of formation of products][ sum of standard enthalpy of formation of reactants]
Thermochemical equation
 The balanced chemical equation equation which also
informs about heat change taking place is known as
thermochemical equation. A thermochemical equation
can be written as follows ;
(a) by writing the heat evolved or absorbed as a term in
the eqation,
C(s)+O2(g)
CO2(g) +393.5 kJ
(b)By using H notation, i.e., writing H = -ve for exothermic
and H = +ve for endothermic reactions, as
C2H5OH(l) + 3O2
2CO2 + 2H2O : ∆H = -1367 kJ
Hess law
 The total enthalpy change remains
constant whether the process takes place
in single step or many steps.
C(s) + O2(g)
∆H2= -110.5
kJ/mol
∆HCO2 = 393.5 kJ/mol
+1/2 O2
CO2(g)
+1/2 O2
CO2(g)
∆H3 = -283 kJ/mol
Enthalpy of different types of
reactions
 Standard enthalpy of combustion(∆cH°):
The enthalpy change taking place when one mol
of compound is undergoes in complete
combution and all reactants and products are in
their standard state.
 Enthalpy of solution(∆solH°):
The enthalpy change taking place when one mol
of solute is dissolved in one mol of solvent .
 Enthalpy of atomization(∆aH°):
The enthalpy change taking place when one mol
of solid or molecule or compound is changed
into constituents atoms.
Born - Haber cycle
 Born – Haber cycle is an indirect method to
construct an enthalpy digaram.
Na(s) + 1/2Cl2(g)
∆subH
Na(g)
∆iH
∆fH°
NaCl(s)
1/2∆aH
Cl(g)
∆egH
Na(g) + Cl(g)
∆Latt H
∆fH° = ∆subH + ∆iH + 1/2∆aH + ∆egH + ∆LattH
Spontaniety
 A process which can take place by itself under the
given sets of conditions once it has been initiated if
necesarrry, is said to be a spontaneous process.
It may be of two types :
(a) Spontaneous process where no initiation is needed.
e.g.
2NO(g)+O2(g)
2 NO2(g)
(b) Spontaneous process where some initiation is
required.
e.g.
CH4(g)+O2(g)
CO2(g)+2H2O(l)
Entropy
 The property of a system which measures
the degree of disorder or randomness in
the system.
it is denoted by S. Entropy is a
state function.
2nd law of thermodynamics
 The entropy of the universe always increase
in the course of every spontaneous change.
 In open system the process will be
spontaneous if ∆STotal is positive.
∆Stotal = ∆Ssys + ∆Ssurr
∆ST > 0, spontaneous
∆ST = 0, equilibrium
∆ST < 0, non - spontaneous
Gibbs energy
The maximum amount of energy available to a system
during a process that can be converted into useful
work is called Gibbs energy.
G = Gibbs free energy
= H – TS
∆STotal = ∆Ssys + ∆Ssurr
= ∆Ssys + (∆H)surr/T
{∆S=qrev/T=∆Hsurr/T }
= ∆Ssys + {(-∆H)sys/T}
{-(∆H)surr=-(∆H)sys}
T∆STotal = T∆Ssys – (∆H)sys
(∆H)sys – T∆STotal = - T∆Stotal
∆Gsys = -T∆STotal
{∆G=∆H-T∆S}
Gibbs energy change & equilibrium
 Gibbs energy change of the system may
be expressed as
∆G = ∆G° + RT lnK/Qc
{ lnK = logeK}
At equilibrium, ∆G = 0
0 = ∆G° + RTlnK
∆G° = -RTlogeK
∆G° = -2.303RT log10K
{Qc = K}