1. MISSION
CHRIST is a nurturing ground for an individual’s
holistic development to make effective contribution to
the society in a dynamic environment
VISION
Excellence and Service
CORE VALUES
Faith in God | Moral Uprightness
Love of Fellow Beings
Social Responsibility | Pursuit of Excellence
Thermodynamics
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Session Guidelines
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Course Outcome
• Students will be able to explain the basic principles of IR spectroscopy
and UV Visible spectroscopy
• Students will be able to outline the oxidation and reduction reactions
which are relevant to study the concepts of corrosion science and
electrochemistry.
• Students will be able to analyze the various types of corrosion occurring
on metal surfaces by knowing electrochemical theory of corrosion
• Students will be able to explain the basic concepts of
thermodynamics, 1st law and 2nd law of thermodynamics
• Students will be able to illustrate the fundamentals of characterization
techniques and waste water treatment
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Syllabus
Definition of thermodynamic terms: system, surrounding etc. Types of
systems, intensive and extensive properties.
First law of thermodynamics, internal energy, enthalpy, relation between
internal energy & enthalpy, heat capacity, free energy.
Second law of thermodynamics, Spontaneous & non-spontaneous reactions,
Gibbs-Helmholtz equation & related problems. Clausius-Clapeyron
equation, Lavoisier & Laplace law.
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What is Thermodynamics?
The science of energy, that concerned with the ways in which energy is
stored within a body.
Energy transformations – mostly involve heat and work movements.
Energy cannot be created or destroyed, but can only be transformed from
one form to another.
If PE of products is less than reactants, the difference must be released
as KE.
All of thermodynamics depends on the law of :
CONSERVATION OF ENERGY
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• System: A quantity of matter or a region in space chosen for study.
• Surroundings: The mass or region outside the system
• Boundary: The real or imaginary surface that separates the system
from its surroundings.
System, surroundings and boundary
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• Isolated system – neither mass nor energy can cross the selected
boundary
• Example (approximate): coffee in a closed, well-insulated thermos
bottle
Type of system (isolated system)
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• Closed system – only energy can cross the selected boundary
• Examples: a tightly capped cup of coffee
Type of system (Closed system)
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• Open system – both mass and energy can cross the selected
boundary
• Example: an open cup of coffee
Type of system (Open system)
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• Intensive properties are those properties whose value is
independent upon the amount of substance present in the system.
Eg:- temperature, boiling point
• Extensive properties are those properties whose value does depends
upon the amount of substance present in the system.
• Eg:- Volume, mass, size
Intensive and extensive properties
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• In a reversible process the system
changes in such a way that the
system and surroundings can be put
back in their original states by
exactly reversing the process.
• Changes are infinitesimally small in a
reversible process.
Reversible Processes
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• Irreversible processes cannot be undone by exactly reversing the
change to the system.
• All Spontaneous processes are irreversible.
• All Real processes are irreversible.
Irreversible Processes
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• Energy cannot be created nor destroyed.
• Therefore, the total energy of the universe is a constant.
• Energy can, however, be converted from one form to another or
transferred from a system to the surroundings or vice versa.
FIRST LAW OF THERMODYNAMICS
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• It is total energy of the system
• Specific Enthalpy,
• Unit of Enthalpy (H) is kJ
• Unit of Specific Enthalpy (h) is kJ/kg
• Most chemical reactions occur at constant P, so
ENTHALPY
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∆H = Hfinal - Hinitial
If Hfinal > Hinitial then ∆ H is positive Process is ENDOTHERMIC
If Hfinal > Hinitial then ∆ H is negative Process is EXOTHERMIC
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What is Internal Energy?
• An energy form inherent in every system is the internal energy, which
arises from the molecular state of motion of matter.
• The symbol E is used for the internal energy and the unit of
measurement is the joules (J).
• Internal energy increases with rising temperature and with changes of
state or phase from solid to liquid and liquid to gas.
• The heat reservoirs store internal energy, and the heat engines convert
some of this thermal energy into various types of mechanical, electrical
and chemical energies.
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The relationship between the enthalpy of the system and the internal
energy of the system is
H = E + PV
The change in the enthalpy of the system during a chemical reaction is equal
to the change in its internal energy plus the change in the product of the
pressure times the volume of the system.
ΔH = Δ E + Δ(PV)
If the reaction is run at constant pressure, the change in the enthalpy that
occurs during the reaction is equal to the change in the internal energy of the
system plus the product of the constant pressure times the change in the
volume of the system.
Δ H = Δ E + P Δ V (at constant pressure)
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Substituting the first law of thermodynamics into this equation gives the
following result.
Δ H = (qp + w) + P Δ V
Assuming that the only work done by the reaction is work of expansion
gives an equation in which the P Δ V terms cancel.
Δ H = (qp - P Δ V) + P Δ V
Thus, the heat given off or absorbed during a chemical reaction at constant
pressure is equal to the change in the enthalpy of the system.
Δ H = qp (at constant pressure)
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The relationship between the change in the internal energy of the system
during a chemical reaction and the enthalpy of reaction can be
summarized as follows.
1. The heat given off or absorbed when a reaction is run at constant
volume is equal to the change in the internal energy of the system.
Δ Esys = qv
2. The heat given off or absorbed when a reaction is run at constant
pressure is equal to the change in the enthalpy of the system.
Δ Hsys = qp
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What is Heat Capacity?
The heat capacity of a substance can be defined as the amount of heat
required to change its temperature by one degree.
Mathematically,
Q=CΔT
Where Q is the heat energy required to bring about a temperature change of
ΔT and C is the heat capacity of the system under study.
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What is specific Heat Capacity?
The specific heat of a substance is the amount of energy required to raise
the temperature of 1 gram of the substance by 1oC .
Mathematically it is given as: Q=msΔT
Here Q is the amount of heat energy required to change the temperature of
m (kg) of a substance by ΔT, s is the specific heat capacity of the system.
Specific Heats of Some Common Substances
Substance Specific Heat (J/go
C)
Water (l) 4.18
Water (s) 2.06
Water (g) 1.87
Ammonia (g) 2.09
Ethanol (l) 2.44
Aluminum (s) 0.897
Carbon, graphite (s) 0.709
Copper (s) 0.385
Gold (s) 0.129
Iron (s) 0.449
Lead (s) 0.129
Mercury (l) 0.140
Silver (s) 0.233
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Gibbs Free Energy (G) –
The energy associated with a chemical reaction that can be used to do
work. The free energy of a system is the sum of its enthalpy (H) plus the
product of the temperature (Kelvin) and the entropy (S) of the system:
G = H - TS
Free energy of reaction (G)
The change in the enthalpy (H) of the system minus the product of the
temperature (Kelvin) and the change in the entropy (S) of the system:
Δ G = Δ H - T Δ S
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Standard-state free energy of reaction (Δ G°)
The free energy of reaction at standard state conditions:
Δ G° = Δ H° - T Δ S°
Standard-state conditions
The partial pressures of any gases involved in the reaction is 0.1 MPa.
The concentrations of all aqueous solutions are 1 M.
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Limitations of the first law of thermodynamics
1. No restriction on the direction of the flow of heat: the first law
establishes definite relationship between the heat absorbed and the work
performed by a system.
2. The first law does not indicate whether heat can flow from a cold end to
a hot end or not.
For example: we cannot extract heat from the ice by cooling it to a
low temperature. Some external work has to be done.
3. Does not specify the feasibility of the reaction: first law does not specify
that process is feasible or not for example: when a rod is heated at one end
then equilibrium has to be obtained which is possible only by some
expenditure of energy.
4. Practically it is not possible to convert the heat energy into an
equivalent amount of work.
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Second Law of Thermodynamics
The entropy of the universe does not change for reversible processes and
increases for spontaneous processes.
Reversible (ideal):
Δ S Universe = Δ S System + Δ S Surroundings = 0
Irreversible (real, spontaneous):
Δ S Universe = Δ S System + Δ S Surroundings > 0
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Spontaneous Processes
• Spontaneity does not imply that the reaction proceeds with great
speed. For example, the decay of diamonds into graphite is a
spontaneous process that occurs very slowly, taking millions of years.
• The rate of a reaction is independent of its spontaneity, and instead
depends on the chemical kinetics of the reaction.
• Every reactant in a spontaneous process has a tendency to form the
corresponding product.
• This tendency is related to stability.
• For spontaneous processes, the change in Gibbs free energy is negative
(ΔG<0).
• Examples of spontaneous processes:
1) Heat flows from hotter body to a colder body.
2) A solid KCl spontaneously dissolves in water.
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• Spontaneous processes are those that
can proceed without any outside
intervention.
• The gas in vessel B will spontaneously
effuse into vessel A, but once the gas is
in both vessels, it will not spontaneously
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• Processes that are spontaneous in one
direction are non spontaneous in the
reverse direction.
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• Processes that are spontaneous at one temperature may be
nonspontaneous at other temperatures.
• Above 0°C it is spontaneous for ice to melt.
• Below 0°C the reverse process is spontaneous.
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Non-Spontaneous Processes
• An endergonic reaction (also called a nonspontaneous reaction or an
unfavorable reaction) is a chemical reaction in which the standard change in
free energy is positive, and energy is absorbed. (ΔG>0).
• The total amount of energy is a loss (it takes more energy to start the reaction
than what is gotten out of it) so the total energy is a negative net result.
• Endergonic reactions can also be pushed by coupling them to another reaction,
which is strongly exergonic, through a shared intermediate.
• A non-spontaneous processes does not takes place on its own.
• It needs continuous external influence.
• Once started, a non-spontaneous processes will stop, when the continuous
external force is withdrawn.
• Examples of non-spontaneous processes:
1) Flow of heat from inside of refrigerator to the room. Room is at
higher temperature than refrigerator.
2) Boiling of water.
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When a process occurs at constant temperature and pressure, P, we can
rearrange the second law of thermodynamics and define a new quantity
known as Gibbs free energy:
Gibbs free energy
G=H−TS where H, is enthalpy, T, is temperature (in kelvin, S, is the entropy.
When using Gibbs free energy to determine the spontaneity of a process,
we are only concerned with changes in G, rather than its absolute value.
The change in Gibbs free energy for a process is thus written as ΔG, which
is the difference between Gibbs free energy of the products, and the Gibbs
free energy of the reactants.
Gibbs free energy and spontaneity
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Spontaneity in terms of free energy change is given by
∆G = ∆H - T∆S
∆G = (-ve) spontaneous
∆G = 0 Equilibrium
∆G = (+ve) Non-Spontaneous
This equation takes into account both the concepts (a) energy factor
and other (b) entropy factor
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Depending upon the signs of H and T.S. and their relative magnitudes the
following possibilities arise:
When both H and T.S are negative: Energy factor favors and randomness oppose
then:
if H <T.S(non-spontaneous), G=negative
When both H = T.S the process is in equilibrium and G =0
when both H and T.S are positive: Energy factor opposes and randomness favours
then:
If H >T. S process is non spontaneous, G =+ve
If H<T.S process is spontaneous, G= -ve
If H=T.S the process is in equilibrium, G=0
When H = -ve and T.S =+ve, process is spontaneous and G= -ve
When H = +ve and T.S =-ve, process is non spontaneous and G= +ve
Effect of temperature on spontaneity of process
Endothermic process at high t-> spontaneous
Exothermic process at low t-> spontaneous
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• It is used for predicting the spontaneity of a process.
Applications of the Gibbs-Helmholtz equation:
• Used in the calculation of the change in enthalpy using the change in
Gibbs energy when the temperature is varied at constant pressure.
• Used in the calculation of the change in enthalpy for a reaction with
temperature other than 298K.
• Used in the calculation of the effect of temperature change on the
equilibrium constant.
Gibbs Helmholtz equation
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CLAUSIUS–CLAPEYRON EQUATION
relates the latent heat (heat of transformation) of vaporization or
condensation to the rate of change of vapor pressure with temperature
or
in the case of a solid-liquid transformation, it relates the latent heat of
fusion or solidification to the rate of change of melting point with
pressure
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Lavoisier and Laplace’s law (1782)
This law may be stated in the general form as the heat change
accompanying a chemical reaction in one direction is exactly equal in
magnitude, but opposite in sign, to that associated with the same
reaction in the reverse direction.
This law states that; the heat change (or enthalpy change) of a chemical
reaction is exactly equal but opposite in sign for the reverse reaction.
This is evident from the following two fractions:
(a) CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) [here, ∆H0 = -890.3 kJ mol-1]
(b) CO2 (g) + 2H2O (l) → CH4 (g) + 2O2 (g) [here, ∆H0 = +890.3 kJ mol-1]
Thus, it can be concluded that ∆Hforward reaction = ∆Hbackword reaction
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Examples
1.
S(s) + O2(g) → SO2(g) ΔH = -296.9 kJ
SO2(g) → S(s) + O2(g) ΔH = +296.9 kJ
2.
C(s) + O2 → CO2 (g); ΔH = – 94.3 kcal
CO2 → C(s) + O2; ΔH = + 94.3 kcal.
