6. Thermodynamics.pptx cytftftvutftuyvuyuyuy

THERMODYNAMICS
The study of the energy transformations from one state to
other state
is thermodynamics
The laws of thermodynamics deal with energy changes of
macroscopic systems which has large number of molecules
It is based on initial and final states of a system undergoing
the change
Laws of thermodynamics applicable only for a system
which is in equilibrium state or moves from one
equilibrium state to another equilibrium state.

The System and the Surroundings
The part of universe in which observations are taken is known
as system
And remaining part is known as surroundings
Universe = System + Surroundings
The wall which separates the system
from the surroundings is called
Boundary.
The region of space in the
neighborhood of the system
constitutes its surroundings

Types of the System
Open System
Exchange of energy and matter
between system and surroundings
Eg: Reactants in an open beaker
Closed System
No exchange of matter, but exchange of
energy is possible between system and the
surroundings
Eg: Reactants in a closed vessel made of
conducting material

Isolated System
No exchange of energy or matter
between the system and the
surroundings
Eg: Reactants in a thermos flask or any
other closed insulated vessel

The State of the System
The state of the system is represented by state functions or
state variables.
The state of a thermodynamic system is described by its
measurable or macroscopic properties.
The state of a gas is described by its pressure (p), volume
(V), temperature (T ), amount (n) etc where p, V, T, n are the
state variables or state functions.

The Internal Energy as a State Function
The internal energy(U) of the system consists of chemical,
electrical, mechanical or any other type of energy.
This internal energy may change when:
Heat passes into or out of the system
Work is done on or by the system
Matter enters or leaves the system

A system which will not allow the exchange of heat between the
system and surroundings is called Adiabatic System
Adiabatic Process is a process in which there is no transfer of
heat between the system and surroundings.
Work:

The state of the system can be changed in two ways:
One way: Mechanical work of 1 kJ is done by rotating a set of
small paddles and by churning water
Let the new state be B state and its temperature is TB with
internal energy is UB. TB >TA
Change in temperature, T=T
∆ B - TA
Change in internal energy, U= U
∆ B - UA
The initial state of the system as state A at temperature TA with
internal energy UA

Second way: An equal amount (i.e., 1kJ) electrical work with the
help of an immersion rod .
Change in temperature, T=T
∆ B - TA
The difference between the value of U in two states is U
∆
which is the adiabatic work required to change the state of
system
∆U= U2 – U1 = wad
wad is positive then work is done on the system
wad is negative then work is done by the system

Heat:
The change in the internal energy of a system by transfer of
heat from the surroundings to the system which results in the
temperature difference is known as heat(q)
Water is in a container having thermally conducting walls at
temperature TA. It is enclosed in a heat reservoir at temperature
TB
The heat absorbed by the water i.e. system is measured in
terms of temperature difference .
q = TB - TA
Internal Energy is U= q ,
∆ when no work is done at constant
volume.
q is positive, when heat is transferred from the surroundings to the
system
q is negative when heat is transferred from system to the surroundings

General Case:
In this ,the state can be changed by doing work and by transfer
of heat.
Change in internal energy is, U = q + w
∆
If there is no transfer of energy as heat or as work (isolated
system) i.e., w = 0 and q= 0, then U= 0
∆
The energy of an isolated system is constant.
This is the first law of thermodynamics
This law states that the energy of an isolated system is constant
i.e., energy can neither be created nor be destroyed.

Consider pressure-volume work(mechanical work).
A cylinder which contains one mole of an ideal gas fitted
with a frictionless piston
Applications
Work:

Volume of the gas is ‘Vi ‘and Pressure of the gas inside is
‘p’
If external pressure pex is applied which is greater than
p, then piston is moved inward till the pressure p
becomes equal to pex
Then final volume will be Vf
Volume change =l A = V =
∆ Vf
-Vi
l is the distance that piston moves
A is cross-sectional area of the piston

Pressure = force/area
Then force on piston = pex . A
Work done on system by piston is w = force ×distance
w = pex . A. l= pex ( - V)
∆
If the pressure is not constant at
every stage of compression, but
changes in number of finite steps,
then work done on the gas is
w = p V
−∑ ∆

If the pressure is not constant but
changes during the process such that
it is infinitesimally greater than the
pressure of the gas, then volume
decreases by an infinitesimal amount
dV. Thus work done is
At each stage pex = (pin + dp) : Compression
At each stage pex = (pin - dp) : Expansion
Thus, in general pex = (pin ± dp) which is called Reversible
Processes

Under reversible conditions the relation of work to internal
pressure of the system is
( neglect dp )
Therefore, at constant temperature
But the gas equation is pV = nRT p = nRT/V

Free Expansion:
Expansion of a gas in vacuum (pex = 0) is called free expansion
We know internal energy is U = q +w
∆
∆U = q - pex V
∆
If a process is carried out at constant volume (ΔV = 0), then
∆U = qv
Isothermal and free expansion of an ideal gas:
For isothermal irreversible change , q= -w = pex V
∆
For isothermal reversible change
For adiabatic change(q= 0) then U= w
∆ ad

Enthalpy (H):
Change in internal energy is, U = q + w
∆
∆U = qp p V ;
− ∆ at constant pressure
U2 – U1 = qp – p(V2 – V1)
qp = (U2 + pV2)-(U1+ pV1)
But Enthalpy H = U+ pV
qp = H2 – H1 = H
∆
Thus H = U + p V
∆ ∆ ∆
∆H is negative for exothermic reactions which evolve heat
during the reaction
∆H is positive for endothermic reactions which absorb heat
from the surroundings

Consider a reaction involving gases. If is the total volume of
the gaseous reactants, is the total volume of the gaseous
products,is the number of moles of gaseous reactants and is
the number of moles of gaseous products, all at constant
pressure and temperature, then using the ideal gas law
=)
RT
is the number of moles of gaseous products minus the number of
moles of gaseous reactants.
Substituting the value of Pv in H = U +p
∆ ∆

Extensive and Intensive Properties:
The property whose value depends on the quantity or size of matter
present in the system is known as extensive property
Eg: mass, volume, internal energy, enthalpy, etc.
The property whose values do not depends on the quantity or size of
matter present in a system is known as intensive property.
Eg: temperature, density, pressure etc
Extensive Property:
Intensive Property:
Here volume is halved but temperature is constant. Thus volume is
an extensive property and temperature is an intensive property

Heat Capacity:
When heat is absorbed by the system then temperature of the
system rises. By this heat transferred to a system can be
measured.
The increase in the temperature is proportional to the heat
transferred,
q = coeff X T
∆
q= C T
∆
Where ‘C’ is called Heat Capacity
‘C’ depends on the size, composition and nature of system
Heat capacity of one mole of the substance is Cm = (C/n)
The heat( q ) required to raise the temperatures of sample is
q = c X m X T = C T
∆ ∆
Where c is specific heat of the substance
m is the mass of the substance
∆T is the temperatures change

Relationship between Cp and Cv for an ideal
gas:
At constant volume, the heat capacity is Cv
∆U = Cv T = q
∆ v
At constant pressure, the heat capacity is Cp
∆H = Cp T = q
∆ p
For a mole of an ideal gas:
∆H = U + (pV)
∆ ∆
∆H = U + (RT)
∆ ∆
∆H = U + R T
∆ ∆
Substituting H and U then C
∆ ∆ p T = C
∆ v T + R T
∆ ∆
Cp = Cv + R
Cp - Cv = R

Measurement of U and H
∆ ∆
Measurement of energy changes associated with chemical
or physical processes by an experimental technique called
Calorimetry.
By knowing heat capacity of the liquid in which calorimeter is
immersed and the heat capacity of calorimeter, the heat
evolved in the process can be determined by temperature
changes .
Measurements are taken under two different conditions:
• At constant pressure, qp
• At constant volume, qv

∆U measurements:
Absorption of heat at constant volume can be measured in a
bomb calorimeter
The energy changes
associated with reactions
are measured with no work
done under constant
volume
Temperature change of the
calorimeter by the completed
reaction is then converted to qv

∆H measurements:
The measurement of heat change at constant atmospheric
pressure
At constant pressure H = q
∆ p
The heat absorbed or evolved at
constant pressure is also called the heat
of reaction or enthalpy of reaction(∆rH)
In an exothermic reaction, heat is evolved then qp is
negative and ∆rH is also negative
In an endothermic reaction, heat is absorbed then qp is
positive and ∆rH is also positive

Enthalpy Change( ∆R H)
∆rH = Sum of enthalpies of products - Sum of enthalpies of
reactants
ai and bi are the stoichiometric coefficients of the products and
reactants
Eg: For the reaction: CH4
(g) + 2O2
(g) CO
→ 2
(g) + 2H2
O (l)
∆rH = [Hm(CO2,g) + 2Hm(H2O, l)] - [Hm(CH4, g)+ 2Hm (O2, g)]
Where Hm is the molar enthalpy
Enthalpy change is a very useful quantity used to calculate
temperature dependence of equilibrium constant
The enthalpy change accompanying a reaction is called the
reaction enthalpy.

Standard Enthalpy Of Reactions( H
∆ -
):
It is the enthalpy change for a reaction when all the
participating substances are in their standard states
The standard state of a substance at a specified temperature
is its pure form at 1 bar
Eg: The standard state of liquid ethanol at 298 K is pure liquid
ethanol at 1 bar

Enthalpy Changes During Phase Transformations:
The enthalpy change that accompanies melting of one
mole of a solid substance in standard state is called
standard Enthalpy of Fusion or MolarEnthalpy of Fusion(∆fus
)
H2O(s) H2O(l) ; ∆fus =6.00 kJ mol-1

The heat required to vaporize one mole of a liquid at
constant temperature and under standard pressure
(1bar) is called its standardEnthalpy of Vaporization or
Molar Enthalpy of Vaporization (∆vap )
H2O(l) H2O(g); ∆vap =+40.79kJ mol-1
The change in enthalpy when one mole of a solid
substance sublimes at a constant temperature and
under standard pressure (1bar) is called Standard
Enthalpy of Sublimation(∆vap )
It is the direct conversion of a solid into its vapour

Standard Enthalpy Changes of Fusion and
Vaporisation

Standard Enthalpy Of Formation:
The standard enthalpy change for the formation of one
mole of a compound from its elements in their stable
states is called Standard Molar Enthalpy of Formation(∆f)
H2 (g) + O2(g) H2O(l); ∆f = -285.8 kJ mol-1

Thermochemical Equations:
C2H5OH(l) + 3O2(g) → 2CO2 (g) + 3H2O(l); ∆r = -1367 kJ mol-1
This type of balanced equation with ∆rvalue is called
Thermochemical Equations
The coefficients in thermo-chemical equation refers to
number of moles of reactants and products
The numerical value of ∆rrefers to the number of moles
of substances specified by an equation

Hess’s Law of Constant Heat
Summation:
If a reaction takes place in several steps then its standard
reaction enthalpy is the sum of the standard enthalpies of the
intermediate reactions into which the overall reaction may be
divided at the same temperature
∆r+

Enthalpies for Different Types of
Reactions
Standard Enthalpy of Combustion:(∆c)
The enthalpy change per mole (or per unit amount) of a
substance, when it undergoes combustion is known as
standard enthalpy of combustion(∆c )
C4H10(g) + 13/2 O2(g) 4CO
→ 2(g) + 5H2O(l);
∆c= -2658.0 kJ mol-1

Enthalpy Of Atomization:(∆a)
The enthalpy change on breaking one mole of bonds completely
to obtain atoms in the gas form is known as Enthalpy of
atomization (∆a)
H2(g) 2H(g);
→ ∆a = 435.0 kJ mol-1
The enthalpy of atomization is same as the enthalpy of
sublimation for this reaction
Na(s) →Na(g) ; ∆a = 108.4 kJ mol-1

Bond Enthalpy:
Energy is required to break a bond and energy is released when
a bond is formed
Two types of bond enthalpy are:
Bond dissociation enthalpy
Mean bond enthalpy
Bond dissociation enthalpy is defined as the energy required to
break one mole of gaseous bonds to form gaseous atoms
The mean bond enthalpy is defined as the energy required to
break a covalent bond or bonds in the gaseous state

Enthalpy of Solution: ()
The enthalpy change that takes place when 1 mole of
a solute dissolves in a solvent to form an ‘infinitely’
dilute solution
NaCl(s) + aq NaCl(aq) ()
= +5 kJ mol-1
Where - Enthalpy of Solution.
can be exothermic or endothermic
Enthalpy of Solution is the sum of two imaginary steps:
reverse of the Lattice Enthalpy and the Hydration
Enthalpy
= -+

NaCl(s) + aq
Na+
(g) + Cl-
(g)
-HLE = +776 kJ mol-
1
Na+
(aq) + Cl-
(aq)
Hhyd = -771 kJ mol-1
Hsol = +5 kJ mol-1
Hsol = -HLE + (Hhyd(cation) + Hhyd(anion))
Hsol = - -776 + -771 = + 5 kJ mol-1
Enthalpy Diagram

Lattice Enthalpy:
The enthalpy change that takes place when 1 mole of a
solid ionic lattice forms from its gaseous ions
Na+
(g)+Cl-
(g)  Na+
Cl-
(s) HLE = -776 kJ mol-1
 HLE is defined exothermically.
 The more closely ions pack together in the solid
lattice the more exothermic is HLE. The smaller
and more highly charged the ions (the greater their
charge density) the closer they pack

Spontaneity
A spontaneous process is one that can occur in a
system left to itself; no action from outside the system
is necessary to bring the change.
If a process is spontaneous, the reverse process is
nonspontaneous and vice versa
A spontaneous process is an irreversible process which
can be reversed by some external agency
The exothermic reactions are spontaneous
The endothermic reactions are nonspontaneous

The decrease in enthalpy in passing from reactants to products
is observed in exothermic reaction
The increase in enthalpy in passing from reactants to products
is observed in endothermic reaction

Entropy and Spontaneity
Consider mixing of two gases which occurs spontaneously, and
the gases form a homogeneous mixture
The driving force is a thermodynamic quantity called Entropy
The total energy of a system remains unchanged in the mixing of
the gases but the number of possibilities for the distribution of
that energy increases

The entropy change is inversely proportional to the temperature
DS = qrev /T
The greater the number of configurations of the microscopic
particles (atoms, ions, molecules) among the energy levels in a
particular state of a system, the greater the entropy of the
system
∆Stotal =∆Ssystem +∆Ssurr
For both reversible and irreversible expansion for an ideal gas,
under isothermal conditions, U = 0 but S
∆ ∆ total is not zero

Gibbs Energy and Spontaneity
The free energy change (∆G) for a process at constant
temperature and pressure is given by the Gibbs equation
∆ Gsys = ∆ Hsys – T ∆ Ssys
If ∆G < 0 (negative), a process is
spontaneous.
If ∆G > 0 (positive), a process is
nonspontaneous.
If ∆G = 0, the process is at
equilibrium.

Gibbs Energy Change and Equilibrium
The reversible reactions are the reactions which
proceed in both the directions with a decrease in free
energy. It is possible only if
at equilibrium the free energy of a system is minimum
The criterion for equilibrium A+B C+D is ∆r = 0
∆r is related to the equilibrium constant of the reaction
as follows: 0 = ∆r + RT ln K
∆r= -RT ln K
∆r= - 2.303 RTlog K

For strongly endothermic reactions, the value of ∆rmay
be large and positive. In such a case, value of K will be
much smaller than 1 and the reaction is unlikely to
form much product.
In case of exothermic reactions, ∆r is large and
negative, and ∆ris likely to be large and negative too. In
such cases, K will be much larger than 1. We may
expect strongly exothermic reactions to have a large K,
and hence can go to near completion.
It is possible to obtain an estimate of ∆r from the
measurement of ∆rand ∆r

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