Thermodynamic Aspects of Evaporation Process

1. Thermodynamic Aspects
of the Evaporation Process

2. • The word Thermodynamics means the study of “Heat in Motion”.
• The study of the flow of heat or any other form of energy into or out
of a system as it undergoes a physical or chemical transformation
(thermodynamical process), is called Thermodynamics.
• A system is that part of the universe which is under thermodynamic
study and the rest of the universe is surroundings.
• The real or imaginary surface separating the system from the
surroundings is called the boundary.
• These boundaries may be either movable or fixed.

3. Types of system
• Open system – a pot of boiling water, a jet engine, or a human body.
• Closed system – a piston-cylinder device or a refrigerator.
• Isolated system –a thermos flask with its contents sealed inside.

4. Types of system
• If the system were insulated from the rest of the universe so that no
heat could pass into or out of the system, it would be called an
adiabatic system and any process that it undergoes would be called
an adiabatic process.
• A large system containing many atoms or molecules is called a
macroscopic system.
• A system consisting of a single atom or molecule is called a
microscopic system.

5. Types of system

6. Properties of systems
• Macroscopic properties refer to physical properties of matter that
can be observed and measured directly with the naked eye or with
simple laboratory equipment e.g., Mass, Volume, Density,
Temperature, Pressure, Conductivity, Elasticity, Viscosity etc.
• Microscopic properties refer to the physical properties of matter at
the atomic or molecular level, which cannot be directly observed
with the naked eye or with simple laboratory equipment and are
determined by the fundamental laws of physics and chemistry e.g.,
Kinetic energy, Momentum, Atomic and molecular structure, Electron
configuration, Energy levels, Thermal motion, Magnetic properties,
Optical properties, Electronic properties etc.

7. Properties of systems
• The macroscopic or bulk properties of a system can be divided into two
classes.
• A property which does not depend on the quantity of matter present in the
system, is known as Intensive property.
• A property that does depend on the quantity of matter present in the
system, is called an Extensive property.

8. Thermodynamic parameters
• Since a change in the magnitude of macroscopic properties alters the
state of the system, these are referred to as State variables or State
functions or Thermodynamic parameters.
• A system in which the state variables have constant values
throughout the system is said to be in a state of thermodynamic
equilibrium.
• If a system is left standing for a sufficiently long time, it will reach a
state of thermodynamic equilibrium.

9. Thermodynamic equilibrium
• At thermodynamic equilibrium, it is sufficient to know a small
number of quantities to fully characterize a system.
• A thermodynamic variable is additive if the value associated with a
system composed of several parts is equal to the sum of the values
associated with each individual part.

10. The criteria for equilibrium
• The temperature of the system must be uniform and must be the
same as the temperature of the surroundings (thermal equilibrium).
• The mechanical properties must be uniform throughout the system
(mechanical equilibrium).
• The chemical composition of the system must be uniform with no net
chemical change (chemical equilibrium).
• If the system is heterogeneous, the state variables of each phase
remain constant in each phase.
• The system which consists of more than one phase is called
heterogeneous system.

11. Thermodynamic Processes
• A system undergoes a thermodynamic process when there is some sort of
energetic change within the system, generally associated with changes in
pressure, volume, internal energy/ temperature, or any sort of heat
transfer.
• Every macroscopic process has a driving force that causes it to proceed
e.g., a temperature difference is the driving force that causes a flow of
heat.
• A reversible process is one that can at any time be reversed in direction by
an infinitesimal change in the driving force.
• A reversible process is sometimes called a quasi-equilibrium process or a
quasi-static process.

12. Thermodynamic Processes
• A quasi-equilibrium process, occurs so slowly and smoothly that the
system is always close to being in thermodynamic equilibrium, with
its state changing infinitesimally slowly in response to small changes
in its environment.
• Irreversible process is a process that cannot be reversed by any small
change in external conditions.
• When a system in a given state goes through a number of different
processes and finally returns to its initial state, the overall process is
called a cycle or Cyclic process.

13. Thermodynamic Processes
• Adiabatic process – a process with no heat transfer into or out of the system.
• Isochoric process – a process with no change in volume, in which case the
system does no work.
• Isobaric process - a process with no change in pressure.
• Isothermal process - a process with no change in temperature.

14. Transformations
• In thermodynamics, one is interested in the transformations between
two equilibrium states.
• In particular, quasistatic transformations.

15. Work and Heat
• In thermodynamics, work and heat is the two ways in which energy
can be transferred to or from a system.
• The amount of work done on an object equals the force exerted on it
times the distance it is moved in the direction of the force.
• The relationship between work, heat, and internal energy is described
by the first law of thermodynamics, which states that the change in
internal energy of a system is equal to the sum of the work done on
the system and the heat added to the system.

16. Work and Heat
• Joseph Black was the first to distinguish between the quantity of heat
and the “intensity” of heat (temperature) and to recognize latent heat
absorbed or given off in phase transitions.
• The symbol of heat is q. If the heat flows from the surroundings into
the system to raise then energy of the system, it is taken to be
positive, +q.

17. Work and Heat
Specific Heat:
• The Specific Heat of a substance is the amount of Energy it requires to
raise the temperature of 1 kg, 1 degree Celsius.
Sensible Heat:
• The energy required to change the temperature of a substance with
no phase change.

18. Enthalpy (H)
• The enthalpy (H) of a thermodynamic system is defined as the sum of
its internal energy and the product of its pressure and volume.

19. Enthalpy (H)
• The pressure-volume term expresses the work required to establish
the system's physical dimensions, i.e. to make room for it by
displacing its surroundings.
• As a state function, enthalpy depends only on the final configuration
of internal energy, pressure, and volume, not on the path taken to
achieve it.

20. The laws of thermodynamics
• The first law states that energy cannot be created or destroyed in an
isolated system.
• The second law of thermodynamics states that the entropy(S) of any
isolated system always increases.
• The third law of thermodynamics states that the entropy of a system
approaches a constant value as the temperature approaches absolute
zero.

21. The first law of thermodynamics
• All energies into the system are equal to all energies leaving the
system plus the change in storage of energies within the system.
• Although all forms of energy are interconvertible, and all can be used
to do work, it is not always possible, even in principle, to convert the
entire available energy into work.

22. The first law of thermodynamics
• The energy of a system can be divided into two types.
• The macroscopic energy includes the overall motion of a system e.g.,
solid in rotation, flow in a fluid etc.
• The internal energy represents the rest of the energy of the system
e.g., molecular motion, energy of interaction between particles, etc.
• In thermodynamics, the macroscopic energy is always constant/ the
system is immobile.
• The energy variations of the system are then equal to the variations
of the internal energy in most cases.

23. The first law of thermodynamics
• For an ideal gas substance, the internal energy depends upon the
kinetic energy and potential energy.
• For an ideal gas, the internal energy depends only on its temperature
and the number of molecules present, and it is independent of the
pressure and volume of the gas.

24. The second law of thermodynamics
• The second law of thermodynamics deals with the direction taken by
spontaneous processes.
• Many processes occur spontaneously in one direction only—that is,
they are irreversible, under a given set of conditions.
• More precisely, an irreversible process is one that depends on path.
• If the process can go in only one direction, then the reverse path
differs fundamentally and the process cannot be reversible.

25. The second law of thermodynamics
• It appears that greater, more uniform dispersal of matter and energy
can be the driving force of a spontaneous process.
• This factor is Entropy. It is another state function.
• Spontaneity is favored by an increase in entropy (S).
• k is the Boltzmann constant (1.38 x 10–23 J/K)
• W is the number of microstates.
• Microstate(W) – A specific configuration of the locations and energies
of the particles in a system.

26. The second law of thermodynamics
• The number of microstates possible is given by:
• n is the number of boxes
• N is the number of particles
• Microstates with equivalent particle arrangements are grouped
together and called distributions.
• The most probable distribution has the largest number of
microstates.
• The most probable distribution is therefore the one of greatest
entropy.

27. The second law of thermodynamics
Consider four particles
distributed among two
boxes.
• There are five possible
distributions for this
system.
• Which distribution is most
probable?

28. The second law of thermodynamics
• Entropy quantifies the energy of a substance that is no longer
available to perform useful work.
• Because entropy tells so much about the usefulness of an amount of
heat transferred in performing work, the steam tables include values
of specific entropy (s = S/m) as part of the information tabulated.
• Entropy is sometimes referred to as a measure of the inability to do
work for a given heat transferred.

29. The second law of thermodynamics
Factors influencing Entropy:
• The phase of the substance.
• The temperature of the substance. Temperature is proportional to the
average kinetic energy of the particles. With higher temperature, the
particles have greater freedom to move around.
• The type and number of particles that make up the substance.
• Variations in the type of particles.

30. The second law of thermodynamics
• The effect of temperature on entropy is due mostly to phase
changes.

31. The second law of thermodynamics
• The change in entropy of the surroundings (ΔSsurr) is directly
proportional to the change in enthalpy of the system.
• ΔSsurr is also inversely proportional to temperature.

32. The third law of thermodynamics
• The entropy of a pure perfect crystalline substance (perfectly ordered
system) at zero Kelvin is zero.
• Zero Kelvin is called absolute zero.
• There is no lower temperature than zero Kelvin.
• At zero Kelvin all molecular movement completely stops.
• There is only one possible way to arrange the molecules.
• The second law of thermodynamics can be used to predict spontaneity.
• But, measurements on the surroundings are seldom made.
• This limits the use of the second law of thermodynamics.

33. The third law of thermodynamics
• It is convenient to have a thermodynamic function that focuses on
just the system and predicts spontaneity.
• The changes in Gibbs free energy (ΔG) or simply change in free
energy allows us to predict spontaneity by focusing on the system
only.

34. The third law of thermodynamics
• The Standard Free Energy of Formation (∆Gf°) for a compound is
defined as the free energy change for the formation of one mole of a
substance from its elements in their standard state at 1 bar and 25 °C
(298K).

35. The third law of thermodynamics
• There is a general equation that enables you to calculate dG under
non-standard conditions.
• T is temperature in K
• R = 0.008314 kJ/molK
• Q is the reaction quotient

36. The third law of thermodynamics
• Another measure of reaction spontaneity is the equilibrium constant,
K.
• For a reaction to be spontaneous K must be greater than 1.
• The relationship between ΔG° and K can be found starting with this
general equation.

37. Scope of Thermodynamics
• Most of the important laws of Physical Chemistry, including the Van't
Hoff law of lowering of vapour pressure, Phase Rule and the
Distribution Law, can be derived from the laws of thermodynamics.
• It tells whether a particular physical or chemical change can occur
under a given set of conditions of temperature, pressure and
concentration.
• It also helps in predicting how far a physical or chemical change can
proceed, until the equilibrium conditions are established.

38. Limitations of Thermodynamics
• Thermodynamics is applicable to macroscopic systems consisting of
matter in bulk and not to microscopic systems of individual atoms or
molecules. It ignores the internal structure of atoms and molecules.
• Thermodynamics does not bother about the time factor. That is, it
does not tell anything regarding the rate of a physical change or a
chemical reaction. It is concerned only with the initial and the final
states of the system.

39. Temperature
• Intensity of heat = Temperature
• Temperature is related to the average kinetic energy of atoms and
molecules in a system.
• Absolute zero is the temperature at which there is no molecular
motion.

40. Phase Change
• Heat Energy entering or leaving a system will cause either a
Temperature Change Q = mc∆T or a Phase Change Q = mL.
• It is possible to design systems to take advantage of the phase
changes to enhance the performance of the system.

41. Phase Change
• There are two variables to consider when looking at phase transition,
pressure (P) and temperature (T).
• For the gas state, the relationship between temperature and pressure
is defined by the equations 𝑃𝑉 = 𝑛 𝑅𝑇 (ideal gas law).
• The melting point is the temperature that a solid phase will become a
liquid.
• At different pressures, different temperatures are required to melt a
substance.
• The boiling point is the temperature that a liquid will evaporate into a
gas phase.

42. Phase Change
• Pressure can be used to change the phase of the substance.

43. Phase Change
• The temperature at which vaporization (boiling) starts to occur for a
given pressure is called the saturation temperature or boiling point.
• The pressure at which vaporization (boiling) starts to occur for a
given temperature is called the saturation pressure.
• For a pure substance there is a definite relationship between
saturation pressure and saturation temperature.
• The graphical representation of this relationship between
temperature and pressure at saturated conditions is called the vapor
pressure curve.

44. Phase Change
• The vapor/ liquid mixture is at saturation
when the conditions of pressure and
temperature fall on the curve.
• If a substance exists as a liquid at the
saturation temperature and pressure, it is
called a saturated liquid.
• If a substance exists entirely as vapor at
saturation temperature, it is called saturated
vapor.
• Sometimes the term dry saturated vapor is
used to emphasize that the quality is 100%.

45. Phase Change
• If the temperature of the liquid is lower than the saturation
temperature for the existing pressure, it is called either a subcooled
liquid (implying that the temperature is lower than the saturation
temperature for the given pressure) or a compressed liquid (implying
that the pressure is greater than the saturation pressure for the given
temperature).
• When the vapor is at a temperature greater than the saturation
temperature, it is said to exist as superheated vapor.
• The substances we call gases are highly superheated vapors.

46. Vapor Quality – Dryness Fraction
• The mass fraction of the
vapor in a two-phase liquid-
vapor region is called the
vapor quality/ dryness
fraction (x).

47. Property Diagrams
The phases of a substance and the relationships between its properties
are most commonly shown on property diagrams. There are five basic
properties of a substance that are usually shown on property
diagrams.
• pressure (P)
• temperature (T)
• specific volume (ν)
• specific enthalpy (h)
• and specific entropy (s)

48. Property Diagrams
There are six different types of property diagrams.
• Pressure Temperature (P-T) diagrams
• Pressure-Specific Volume (P-ν) diagrams
• Pressure-Enthalpy (P-h) diagrams
• Enthalpy-Temperature (h-T) diagrams
• Temperature-entropy (T-s) diagrams
• Enthalpy-Entropy (h-s) or Mollier diagrams.

49. Pressure -Temperature Diagram of Water
• The line that separates the solid and
vapor phases is called the
sublimation line.
• The line that separates the solid and
liquid phases is called the fusion
line.
• The line that separates the liquid
and vapor phases is called the
vaporization line.
• A temperature and pressure at
which all three phases exist in
equilibrium is called Triple Point.

50. Pressure -Temperature Diagram of Water
• The point where the vaporization line ends is called the critical point.
• At temperatures and pressures greater than those at the critical
point, no substance can exist as a liquid no matter how great the
pressure is exerted upon it.

51. Enthalpy-Temperature (h-T) Diagram
• The region between the saturated
liquid line and the saturated vapor
line represents the area of two
phases existing at the same time.
• The vertical distance between the
two saturation lines represents the
latent heat of vaporization.
• Operation outside the saturation
lines results in a subcooled liquid
or superheated steam.