1. Thermodynamics
Thermodynamics is the branch of Chemistry that is built upon the fundamental laws that heat and work obey.
Thermodynamics is the study of the effects of work, heat, and energy on a system.
Thermodynamics is only concerned with macroscopic (large-scale) changes and observations.

2. SYSTEM AND SURROUNDING
A system is the part of the universe in which we have a special interest. The surroundings are where we make our observations on the systems.
There are three types of systems:
An open system is a system that can exchange both energy and matter with its surroundings.
A closed system is a system that can exchange energy but not matter with its surroundings.
An isolated system is a system that can exchange neither matter nor energy with its surroundings

3. SYSTEM AND SURROUNDING
An example of an open system is a flask that is not stoppered and to which various substances can be added. A biochemical cell is an open system because nutrients and waste can pass through the cell wall. An example of a closed system is a stoppered flask – energy can be exchanged with the contents of the flask because the walls may be able to conduct heat. An example of an isolated system is a sealed flask that is thermally, mechanically and electrically insulated from its surroundings.

4. INTERNAL ENERGY(U)
The internal energy, U, of the system is a measure of the “energy reserves” of the system. It is the sum of all the kinetic and potential contributions to the energy of all the atoms, ions, and molecules in the system. When energy is transferred into the system by heating it or doing work on it, the increased energy is stored in the increased kinetic and potential energies of the molecules.

5. THE ENTHALPY (H) H = U + PV △H=△U + P△V
Enthalpy may be defined as the heat content of a system at constant pressure.
△H= - ve, exothermic (energy is evolved)
△H= + ve, endothermic (energy is absorbed)

6. SOME DEFINITIONS
Standard Enthalpy of Combustion, △Hco, is the change in enthalpy when one mole of a substance is completely burnt in oxygen under standard conditions. It is always negative, indicating that combustion is an exothermic process.
Standard Enthalpy of Formation, △Hfo , is the change in enthalpy when one mole of a substance is formed from its elements in their reference (most stable) states. It may be positive (endothermic) or negative (exothermic). For example, the standard enthalpy of formation of liquid water (at 25oC, 1 bar) is the change in enthalpy when one mole of liquid water is form gaseous hydrogen and gaseous oxygen.

7. SOME DEFINITIONS
Standard Enthalpy of Dissociation, △Hco , is the enthalpy change when 1 mol of a particular bond of a substance is cleaved under standard states. It is always positive, i.e. endothermic. Thus bond breaking is an endothermic process whereas bond formation is an exothermic process.
Molar Enthalpy of Fusion, △Hfus, is the change in enthalpy change when 1 mole of a substance is melted at its melting point. Enthalpy of fusion is negative, i.e. endothermic process.
Enthalpy of Vaporization, △Hvap , is the enthalpy change when 1 mole of a liquid is vaporized completely to gas at 1 atm pressure. It is endothermic process.
Enthalpy of Sublimation, △Hsub, is the change in enthalpy when 1 mole of a solid is converted into gas at 1 atm pressure. It is an endothermic process.

8. LAWS OF THERMODYNAMICS
The Zeroth Law of Thermodynamics
Thermal equilibrium: Two systems are said to be in thermal equilibrium if there is no net flow of heat between them when they are brought into thermal contact.
Zeroth Law of Thermodynamics:
“When two systems individually are in thermal equilibrium with a third system, then all the three systems are in thermal equilibrium with each other.”

9. THE FIRST LAW OF THERMODYNAMICS
The internal energy of a system changes from an initial value Ui to a final value of Uf due to heat Q and work.
Q is positive when the system gains heat and negative when it loses heat. W is positive when work is done by the system and negative when work is done on the system.
Q is

10. THE FIRST LAW OF THERMODYNAMICS
“It states that energy can neither be created nor destroyed, but it can be converted from one form to another.”
Most chemical reactions release energy or absorb it as they occur. So, according to the conservation of energy all such changes must involve only the conversion of energy from one form to another, or its transfer from place to place, but not created or destroyed

11. THERMODYNAMIC PROCESSES
isothermal process, one that takes place at constant temperature. (when the system is an ideal gas.)
There is adiabatic process, one that occurs without the transfer of heat . Since there is no heat transfer, Q equals zero, and the first law indicates that U = Q – W = –W. Thus, when work is done by a system adiabatically, W is positive and the internal energy of the system decreases by exactly the amount of the work done. When work is done on a system adiabatically, W is negative and the internal energy increases correspondingly.

12. THERMODYNAMIC PROCESSES
An isobaric process is one that occurs at constant pressure.
An isochoric process is one that occurs at constant volume.

13. ISOTHERMAL EXPANSION OR COMPRESSION
P = nRT/V
W = P ∆V = P(Vf – Vi)

14. Isothermal Expansion of an Ideal Gas
Two moles of the monatomic gas argon expand isothermally at 298 K, from an initial volume of Vi = m3 to a final volume of Vf = m3. Assuming that argon is an ideal gas, find (a) the work done by the gas, (b) the change in the internal energy of the gas, and (c) the heat supplied to the gas.

15. ADIABATIC EXPANSION OR COMPRESSION

16. ADIABATIC EXPANSION OR COMPRESSION
[Ti = PiVi/(nR)]
[Tf = PfVf/(nR)].

17. First Law of Thermodynamics
Type of Thermal Process 
 Work Done 
 First Law of Thermodynamics
(DU = Q – W) 
 Isobaric (constant pressure) 
 W = P(Vf – Vi) 
                                             
 Isochoric (constant volume) 
 W = 0 J 
                             
 Isothermal (constant temperature) 
                
(for an ideal gas) 
                                                         
 Adiabatic (no heat flow) 
                                      
(for a monatomic ideal gas) 
                                                    

18. Specific Heat Capacities
where the capital letter C refers to the molar specific heat capacity in units of J/(mol·K).

19. Specific Heat Capacities

21. TEMPERATURE DEPENDENCE OF ENTHAPLY (KIRCHHOFF’S EQUATION)

22. The Second Law of Thermodynamics
THE SECOND LAW OF THERMODYNAMICS: THE HEAT FLOW STATEMENT
“Heat flows spontaneously from a substance at a higher temperature to a substance at a lower temperature and does not flow spontaneously in the reverse direction.”

23. Carnot's Principle and the Carnot Engine
A reversible process is one in which both the system and its environment can be returned to exactly the states they were in before the process occurred.
CARNOT’S PRINCIPLE: AN ALTERNATIVE STATEMENT OF THE SECOND LAW OF THERMODYNAMICS
“No irreversible engine operating between two reservoirs at constant temperatures can have a greater efficiency than a reversible engine operating between the same temperatures. Furthermore, all reversible engines operating between the same temperatures have the same efficiency.”

24. Heat Engines Limits on the Efficiency of a Heat Engine
It is the second law that limits the efficiencies of heat engines to values less than 100%.
Efficiencies are often quoted as percentages obtained by multiplying the ratio W/QH by a factor of 100.

25. Limits on the Efficiency of a Heat Engine
An automobile engine has an efficiency of 22.0% and produces 2510 J of work. How much heat is rejected by the engine?

26. Entropy
In general, irreversible processes cause us to lose some, but not necessarily all, of the ability to perform work. This partial loss can be expressed in terms of a concept called entropy.
Reversible processes do not alter the total entropy of the universe.

27. “The Entropy of the Universe Increases”
1200 J of heat flow spontaneously through a copper rod from a hot reservoir at 650 K to a cold reservoir at 350 K. Determine the amount by which this irreversible process changes the entropy of the universe, assuming that no other changes occur.

28. Entropy
A block of ice is an example of an ordered system relative to a puddle of water.

29. Order to Disorder
Find the change in entropy that results when a 2.3-kg block of ice melts slowly (reversibly) at 273 K (0 °C).

30. ENTROPY
That is, the change in entropy of a substance is equal to the energy transferred as heat to it reversibly divided by the temperature at which the transfer takes place.
The unit of entropy is J K-1. It is an extensive property.
However, molar entropy is an intensive property with its unit J K-1 mol-1.
Entropy is a state function., like internal energy.

31. The Variation of Entropy with Volume
Thus, on expansion, Vf>Vi, so , i.e. entropy increases on expansion.

32. The Variation of Entropy with Temperature
The total change in entropy, , when the temperature changes from Ti to Tf is the sum (integral) of all such infinitesimal terms:

33. The Variation of Entropy During Phase Change
∆ Sfus =
∆ Svap =
Trouton noticed that △Hvap/Tb is approximately the same, equal to about 85 J K-1mol-1, for all liquids except when hydrogen bonding or some other kind of specific bonding is present.. This is known as Trouton’s rule
Trouton’s rule is explained by the fact when a liquid vaporizes, the compact condensed phase changes into a widely dispersed gas that occupies approximately the same volume whatever its identity. To a good approximation, therefore, we expect the increase in disorder, and therefore the entropy of vaporization, to be almost the same for all liquids at their boiling temperatures. △Hvap/Tb for water is high (109.1) because of hydrogen bonding.

34. Absolute Entropies and Third Law of thermodynamics
THE THIRD LAW OF THERMODYNAMICS
It is not possible to lower the temperature of any system to absolute zero (T = 0 K) in a finite number of steps.
“The entropies of all perfectly crystalline substances are the same at T=0”.
“The entropy of a perfectly crystalline substance is zero at absolute zero”.

35. Absolute Entropies Gas Liquid Solid
Water vapour Water Ice 45
NH C6H Diamond, C 2.4
CO CH3CH2OH Graphite, C 5.7
He CaO 39.8
H CaCO
Ne Cu 33.2
N Lead, Pb 64.8
O MgCo
C12H22O

36. THE GIBB’S FREE ENERGY (G)
Gibb’s free energy is defined as
G = H – TS
Because H, T and S are state functions, G is a state function too. A change in Gibb’s free energy, △G, at constant temperature arises from changes in enthalpy and entropy, and is given by
At constant temperature, ∆G = ∆ H - T ∆ S
It is the driving force for chemical reactions. It is also considered as measure of available energy compared to the entropy , which is a measure of unavailable energy.
∆G < 0: spontaneous (exergonic)
∆ G = 0: equilibrium
∆ G > 0: no spontaneous reaction possible (endergonic)