Chemical Thermodynamics Class 11 Notes Chemistry Chapter 5

 Page 1


? System : Specific part of universe in which thermodynamic observations
are made.
? Surroundings : Everything which surrounds the system.
? Types of the System :
(i) Open System : Exchange both matter and energy with the
surroundings. For example : Reactants in an open test tube.
(ii) Closed System : Exchange energy but no matter with the surroundings.
For example : Reactants in a closed vessel.
(iii) Isolated System : Neither exchange energy nor matter with the
surroundings. For example : Reactants in a thermos flask. No system is
perfectly isolated.
? Thermodynamic Processes :
(i) Isothermal process : ?T = 0
(ii) Adiabatic process : ?q = 0
(iii) Isobaric process : ?P = 0
(iv) Isochoric process : ?V = 0
(v) Cyclic process : ?U = 0
(vi) Reversible process : Process which proceeds infinitely slowly by a
series of equilibrium steps.
(vii) Irreversible process : Process which proceeds rapidly and the system
does not have chance to achieve equilibrium.
? Extensive Properties : Properties which depend upon the quantity or
size of matter present in the system. For example : mass, volume, internal
energy, enthalpy, heat capacity, work etc.
Page 2


? System : Specific part of universe in which thermodynamic observations
are made.
? Surroundings : Everything which surrounds the system.
? Types of the System :
(i) Open System : Exchange both matter and energy with the
surroundings. For example : Reactants in an open test tube.
(ii) Closed System : Exchange energy but no matter with the surroundings.
For example : Reactants in a closed vessel.
(iii) Isolated System : Neither exchange energy nor matter with the
surroundings. For example : Reactants in a thermos flask. No system is
perfectly isolated.
? Thermodynamic Processes :
(i) Isothermal process : ?T = 0
(ii) Adiabatic process : ?q = 0
(iii) Isobaric process : ?P = 0
(iv) Isochoric process : ?V = 0
(v) Cyclic process : ?U = 0
(vi) Reversible process : Process which proceeds infinitely slowly by a
series of equilibrium steps.
(vii) Irreversible process : Process which proceeds rapidly and the system
does not have chance to achieve equilibrium.
? Extensive Properties : Properties which depend upon the quantity or
size of matter present in the system. For example : mass, volume, internal
energy, enthalpy, heat capacity, work etc.
? Intensive Properties : Properties which do not depend upon the quantity
or size of matter present in the system. For example : temperature, density,
pressure, surface tension, viscosity, refractive index, boiling point, melting
point etc.
? State Functions : The variables of functions whose value depend only on
the state of a system or they are path independent. For example : pressure
(P), volume (V), temperature (T), enthalpy (H), free energy (G), internal
energy (U), entropy (S), etc.
? Internal Energy (U) : It is the sum of all kind of energies possessed by
the system.
? First Law of Thermodynamics : “The energy of an isolated system is
constant.”
Mathematical Form : ?U = q + w
? Sign Conventions for Heat (q) and Work (w) :
(i) W = + ve, if work is done on system
(ii) W = – ve, if work is done by system
(iii) q = + ve, if heat is absorbed by the system
(iv) q = – ve, if heat is evolved by the system
? Work of Expansion/compression : w = – P
ext 
(V
f
 — V
i
)
? Work done in Isothermal Reversible Expansion of an Ideal Gas :
 w
rev
 = – 2.303 nRT log
Or, w
rev 
= – 2.303 nRT log
? Significance of ?H and ?U : ?H = q
p
 and ?U = q
v
? Relation between ?H and ?U: ?H = ?U + (n
p
 – n
r
)RT for gaseous
reaction.
(i) ?H = ?U if (n
p
 – n
r
) is zero; e.g.,  H
2
(g) + I
2
(g) ? 2HI(g)
(ii)  ?H > ?U if (n
p
 – n
r
) is positive; e.g., PC1
5
(g) ? PCl
3
(g) + C1
2
(g)
(iii)  ?H < ?U if (n
p
 – n
r
) is negative; e.g., N
2
(g) + 3H
2
(g) ? 2NH
3
(g)
? Heat capacity (C) : Amount of heat required to raise the temperature of
a substance by 1°C or 1 K.
q     = C ?T
Page 3


? System : Specific part of universe in which thermodynamic observations
are made.
? Surroundings : Everything which surrounds the system.
? Types of the System :
(i) Open System : Exchange both matter and energy with the
surroundings. For example : Reactants in an open test tube.
(ii) Closed System : Exchange energy but no matter with the surroundings.
For example : Reactants in a closed vessel.
(iii) Isolated System : Neither exchange energy nor matter with the
surroundings. For example : Reactants in a thermos flask. No system is
perfectly isolated.
? Thermodynamic Processes :
(i) Isothermal process : ?T = 0
(ii) Adiabatic process : ?q = 0
(iii) Isobaric process : ?P = 0
(iv) Isochoric process : ?V = 0
(v) Cyclic process : ?U = 0
(vi) Reversible process : Process which proceeds infinitely slowly by a
series of equilibrium steps.
(vii) Irreversible process : Process which proceeds rapidly and the system
does not have chance to achieve equilibrium.
? Extensive Properties : Properties which depend upon the quantity or
size of matter present in the system. For example : mass, volume, internal
energy, enthalpy, heat capacity, work etc.
? Intensive Properties : Properties which do not depend upon the quantity
or size of matter present in the system. For example : temperature, density,
pressure, surface tension, viscosity, refractive index, boiling point, melting
point etc.
? State Functions : The variables of functions whose value depend only on
the state of a system or they are path independent. For example : pressure
(P), volume (V), temperature (T), enthalpy (H), free energy (G), internal
energy (U), entropy (S), etc.
? Internal Energy (U) : It is the sum of all kind of energies possessed by
the system.
? First Law of Thermodynamics : “The energy of an isolated system is
constant.”
Mathematical Form : ?U = q + w
? Sign Conventions for Heat (q) and Work (w) :
(i) W = + ve, if work is done on system
(ii) W = – ve, if work is done by system
(iii) q = + ve, if heat is absorbed by the system
(iv) q = – ve, if heat is evolved by the system
? Work of Expansion/compression : w = – P
ext 
(V
f
 — V
i
)
? Work done in Isothermal Reversible Expansion of an Ideal Gas :
 w
rev
 = – 2.303 nRT log
Or, w
rev 
= – 2.303 nRT log
? Significance of ?H and ?U : ?H = q
p
 and ?U = q
v
? Relation between ?H and ?U: ?H = ?U + (n
p
 – n
r
)RT for gaseous
reaction.
(i) ?H = ?U if (n
p
 – n
r
) is zero; e.g.,  H
2
(g) + I
2
(g) ? 2HI(g)
(ii)  ?H > ?U if (n
p
 – n
r
) is positive; e.g., PC1
5
(g) ? PCl
3
(g) + C1
2
(g)
(iii)  ?H < ?U if (n
p
 – n
r
) is negative; e.g., N
2
(g) + 3H
2
(g) ? 2NH
3
(g)
? Heat capacity (C) : Amount of heat required to raise the temperature of
a substance by 1°C or 1 K.
q     = C ?T
? S p ecifi c 	 h eat	 cap acity 	 (C
s
) : Amount of heat required to raise the
temperature of 1 g of a substance by 1°C or 1K.
q = C
s 
× m × ?T
? Molar Heat Capacity (C
m
) : Amount of heat required to raise the
temperature of 1 mole of a substance by 1°C or 1K.
q = C
m 
× n × ?T
? Standard State of a Substance : The standard state of a substance at a
specified temperature is its, pure form at 1 bar.
? Standard Enthalpy of Formation ( ?
f
H
?
) : Enthalpy change accompanying
the formation of one mole of a substance from its constituent elements
under standard condition of temperature (normally 298 K) and pressure
(1 bar).
Ø ?
f
H
?
 of an element in standard state is taken as zero.
Ø Compounds with – ve value of ?
f
H
? 
are more stable than their
 constituents.
Ø ?
r
H° = S
i
a
i
?
f 
H
?
 (products) – S
i
b
i
?
f 
H
?
 (reactants) : Where ‘a’ and  
 ‘b’ are coefficients of products and reactants in balanced equation.  
? Standard Enthalpy of Combustion ( ?
c
H
?
) : Enthalpy change
accompanying the complete combustion of one mole of a substance under
standard conditions (298 K, 1 bar)
? Hess’s Law of Constant Heat Summation : The total enthalpy change
of a reaction remains same whether it takes place in one step or in several
steps.
? Bond Dissociation Enthalpy : Enthalpy change when one mole of a
gaseous covalent bond is broken to form products in gas phase. For
example : Cl
2
(g) ? 2Cl(g); ?
Cl-Cl
 H
?
 = 242k/mol
–1
.
 (i) For diatomic gaseous molecules; Bond enthalpy = Bond dissociation 
Enthalpy = Atomization Enthalpy.
 (ii) For Polyatomic gaseous molecules; Bond Enthalpy = Average of the bond 
dissociation enthalpies of the bonds of the same type.
? ?
r
H
?
 = S?
bond
H
?
 (Reactants) –– S?
bond
H
?
 (Products).
? Spontaneous Reaction : A reaction which can take place either of its own
or under some initiation.
Page 4


? System : Specific part of universe in which thermodynamic observations
are made.
? Surroundings : Everything which surrounds the system.
? Types of the System :
(i) Open System : Exchange both matter and energy with the
surroundings. For example : Reactants in an open test tube.
(ii) Closed System : Exchange energy but no matter with the surroundings.
For example : Reactants in a closed vessel.
(iii) Isolated System : Neither exchange energy nor matter with the
surroundings. For example : Reactants in a thermos flask. No system is
perfectly isolated.
? Thermodynamic Processes :
(i) Isothermal process : ?T = 0
(ii) Adiabatic process : ?q = 0
(iii) Isobaric process : ?P = 0
(iv) Isochoric process : ?V = 0
(v) Cyclic process : ?U = 0
(vi) Reversible process : Process which proceeds infinitely slowly by a
series of equilibrium steps.
(vii) Irreversible process : Process which proceeds rapidly and the system
does not have chance to achieve equilibrium.
? Extensive Properties : Properties which depend upon the quantity or
size of matter present in the system. For example : mass, volume, internal
energy, enthalpy, heat capacity, work etc.
? Intensive Properties : Properties which do not depend upon the quantity
or size of matter present in the system. For example : temperature, density,
pressure, surface tension, viscosity, refractive index, boiling point, melting
point etc.
? State Functions : The variables of functions whose value depend only on
the state of a system or they are path independent. For example : pressure
(P), volume (V), temperature (T), enthalpy (H), free energy (G), internal
energy (U), entropy (S), etc.
? Internal Energy (U) : It is the sum of all kind of energies possessed by
the system.
? First Law of Thermodynamics : “The energy of an isolated system is
constant.”
Mathematical Form : ?U = q + w
? Sign Conventions for Heat (q) and Work (w) :
(i) W = + ve, if work is done on system
(ii) W = – ve, if work is done by system
(iii) q = + ve, if heat is absorbed by the system
(iv) q = – ve, if heat is evolved by the system
? Work of Expansion/compression : w = – P
ext 
(V
f
 — V
i
)
? Work done in Isothermal Reversible Expansion of an Ideal Gas :
 w
rev
 = – 2.303 nRT log
Or, w
rev 
= – 2.303 nRT log
? Significance of ?H and ?U : ?H = q
p
 and ?U = q
v
? Relation between ?H and ?U: ?H = ?U + (n
p
 – n
r
)RT for gaseous
reaction.
(i) ?H = ?U if (n
p
 – n
r
) is zero; e.g.,  H
2
(g) + I
2
(g) ? 2HI(g)
(ii)  ?H > ?U if (n
p
 – n
r
) is positive; e.g., PC1
5
(g) ? PCl
3
(g) + C1
2
(g)
(iii)  ?H < ?U if (n
p
 – n
r
) is negative; e.g., N
2
(g) + 3H
2
(g) ? 2NH
3
(g)
? Heat capacity (C) : Amount of heat required to raise the temperature of
a substance by 1°C or 1 K.
q     = C ?T
? S p ecifi c 	 h eat	 cap acity 	 (C
s
) : Amount of heat required to raise the
temperature of 1 g of a substance by 1°C or 1K.
q = C
s 
× m × ?T
? Molar Heat Capacity (C
m
) : Amount of heat required to raise the
temperature of 1 mole of a substance by 1°C or 1K.
q = C
m 
× n × ?T
? Standard State of a Substance : The standard state of a substance at a
specified temperature is its, pure form at 1 bar.
? Standard Enthalpy of Formation ( ?
f
H
?
) : Enthalpy change accompanying
the formation of one mole of a substance from its constituent elements
under standard condition of temperature (normally 298 K) and pressure
(1 bar).
Ø ?
f
H
?
 of an element in standard state is taken as zero.
Ø Compounds with – ve value of ?
f
H
? 
are more stable than their
 constituents.
Ø ?
r
H° = S
i
a
i
?
f 
H
?
 (products) – S
i
b
i
?
f 
H
?
 (reactants) : Where ‘a’ and  
 ‘b’ are coefficients of products and reactants in balanced equation.  
? Standard Enthalpy of Combustion ( ?
c
H
?
) : Enthalpy change
accompanying the complete combustion of one mole of a substance under
standard conditions (298 K, 1 bar)
? Hess’s Law of Constant Heat Summation : The total enthalpy change
of a reaction remains same whether it takes place in one step or in several
steps.
? Bond Dissociation Enthalpy : Enthalpy change when one mole of a
gaseous covalent bond is broken to form products in gas phase. For
example : Cl
2
(g) ? 2Cl(g); ?
Cl-Cl
 H
?
 = 242k/mol
–1
.
 (i) For diatomic gaseous molecules; Bond enthalpy = Bond dissociation 
Enthalpy = Atomization Enthalpy.
 (ii) For Polyatomic gaseous molecules; Bond Enthalpy = Average of the bond 
dissociation enthalpies of the bonds of the same type.
? ?
r
H
?
 = S?
bond
H
?
 (Reactants) –– S?
bond
H
?
 (Products).
? Spontaneous Reaction : A reaction which can take place either of its own
or under some initiation.
? Entropy (S) : It is measure of degree of randomness or disorder of a
system. ?S
sys
 =  . Unit of Entropy = JK
–1
 mol
–1
? Second Law of Thermodynamics : For all the spontaneous processes
totally entropy change must be positive.
?S
total 
= ?S
sys 
+
 
?S
surr
 > 0
? Gibbs Helmholtz Equation for determination of Spontaneity :
			?G = ?H – T ?S
(i)  If ?G = – ve, the process is spontaneous
  (ii) If ?G = + ve, the process is non-spontaneous
  (iii) If ?G = 0, the process is in equilibrium
? Relation between Gibbs Energy Change and Equilibrium Constant :
?G
?
 = – 2.303 RT log K
c
.
? Third law of thermodynamic : The entropy of a perfectly crystalline
solid at absolute zero (0 K) is taken to be zero.