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Chemical Thermodynamics Class 11 Notes Chemistry Chapter 5

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? 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.
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FAQs on Chemical Thermodynamics Class 11 Notes Chemistry Chapter 5

1. What is chemical thermodynamics?
Ans. Chemical thermodynamics is a branch of physical chemistry that deals with the study of energy changes and transformations during chemical reactions and processes. It focuses on understanding the principles and laws governing the behavior of energy in chemical systems.
2. What are the main components of chemical thermodynamics?
Ans. Chemical thermodynamics consists of three main components: energy, entropy, and spontaneity. Energy refers to the capacity of a system to do work, while entropy represents the measure of disorder or randomness in a system. Spontaneity refers to the tendency of a reaction to occur without external intervention.
3. How is chemical thermodynamics applied in real-life scenarios?
Ans. Chemical thermodynamics finds various applications in real-life scenarios. It is used to determine the feasibility and efficiency of industrial processes, such as determining the energy required for a chemical reaction or optimizing reaction conditions. It also helps in understanding and predicting the behavior of substances under different conditions, such as the stability of pharmaceutical drugs or the efficiency of energy storage devices.
4. What are the laws of thermodynamics and their significance in chemical thermodynamics?
Ans. The laws of thermodynamics are fundamental principles governing energy and its transformations. The three laws are: 1. The first law of thermodynamics (law of energy conservation) states that energy cannot be created or destroyed, only transferred or converted from one form to another. It is significant in chemical thermodynamics as it helps in understanding the energy changes during chemical reactions. 2. The second law of thermodynamics states that the entropy of an isolated system tends to increase over time, leading to a decrease in the availability of energy for useful work. This law is crucial in understanding the spontaneity of chemical reactions and the direction in which they proceed. 3. The third law of thermodynamics states that as the temperature approaches absolute zero, the entropy of a pure crystalline substance approaches zero. This law provides a reference point for measuring and calculating absolute entropy values.
5. How is Gibbs free energy related to chemical thermodynamics?
Ans. Gibbs free energy is a thermodynamic potential that combines the effects of both enthalpy (heat) and entropy (disorder) in a system. It is used to determine the spontaneity and equilibrium of a chemical reaction. If the Gibbs free energy change (ΔG) is negative, the reaction is spontaneous and can proceed in the forward direction. If ΔG is positive, the reaction is non-spontaneous and requires energy input to proceed.
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