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1. Thermodynamic

2. Energy Different form of energy- (interconvert able) K.E P.E Heat energy (thermal energy) Radiant energy (electromagnetic or light radication ) Electrical energy Chemical energy

3. Thermodynamics Thermo- heat and dynamic – motion . Branch of science which deals with the study of interconversion of different forms of energy and the quantitative relationship between them taking place in physical and chemical process .

4. Limitations Does not give info. Rate of physical or chemical process. Doesn’t describe status, mechanism, history of the process. Only deals with microscopic systems. Thermodynamic is the only dynamics which does not consider time factor.

5. Macroscopic Science Deals with large no. of atom or molecule not individual atom or molecule.

6. Chemical thermodynamics Which gives a quantitative information of the energy change accompanying chemical process and explains chemical behavior , Eg . Heat of reaction, effect of temperature on chemical reactions, etc.

7. System, Surrounding and Boundry System- the portion of the universe under thermodynamic consideration to study thermodynamic properties is called a system. Under universe portion is system. Here thermodynamics means P, V, T, n, E etc. System may be very large or very small. System is confined by a real or an imaginary boundry . Human, boil water, animal etc.

8. Surroundings The remaining portion of the universe. It represent large stock of mass and energy. Can exchange energy with system when allowed. Eg . Universe, environment, earth etc.

9. Boundary The wall separating the system from its surrounding. Boundary may be real or imaginary. Boundary exchange heat, matter, between system and surrounding. Everything outside the boundary is surrounding. Eg . Hot water beaker wall of the beaker = real boundary. While open portion show imaginary boundry .

10. Types of system Open system Closed system Isolated system Homogeneous system Heterogeneous system

11. Open system System which exchange both matter and energy with its surroundings. Eg . Beaker containing water. Exchange- Water continuosly absorbs energy from its surroundings and from vapour relase it.

12. Closed system Exchange energy not matter. Example- closed vessel containing hot water so that only heat is lost not matter.

13. Isolated system Can neither exchange energy nor matter. Example- hot water filled in thermally insulated closed vessel like thermos flask . in actual perfect isolated system is not possible. Universe is an eg . Of isolated system. Universe has no boundary, surrounding.

14. Homogeneous system Only one phase system. Single component system – Zn, O, Water. Solution miscible liquid- water and alcohol or NaCl and water etc. Mixture of gases- H, N, O etc.

15. Heterogeneous system Two separated phase by boundary. Mixture of immiscible liquid- Water and benzen . Solid identical with liquid- ice and water. Liquid identical with vapour - water and vapour .

16. Thermodynamic Properties of system Variable on which property of the system depend. For eg . P, T, V, D, E etc The property of the system classified as Extensive Property Intensive Property

17. Extensive Property Whose magnitude depend on the amount of the matter present in the system. When the amount of matter change its magnitude also change . Additive. Example- enthalpy, mass, volume, energy, weight.

18. Intensive property Magnitude independent of the amount of matter. The ration of extensive property represent an intensive property. Eg . Density= M/V. Example- B.P- take either 1ml or 1L B.P of water is 100̊˚C. M.P, F.P, Surface tension, specific heat, molar heat capacity, T, P, D, viscosity.

19. State and state function State variable- Measurable property of a system like P, T, V etc. State describe value of these variable. When one or more variable change system change to new state. Macroscopic properties of a system depend on these state variable.

20. State function Which depend on initial and final state but independent of the path followed by the system during the process. Eg . Mass (initial to final product), P, T, V etc. It depend on state. Diagram shown by board. It shows changes are independent of all three path but depend on initial and final state.

21. Thermodynamic Equilibrium No change in any thermodynamic function or state function like energy, pressure etc with time. Type- Thermal equilibrium Chemical equilibrium Mechanical equilibrium

22. Thermal equilibrium System and surrounding at same temperature and no exchange of heat. Total energy remain const. eg . Water with its vapour at constant temperature.

23. Chemical equilibrium Chemical composition doesnot change with time. Eg . N + 3H2 → 2NH3 Composition of Reactant and product does not change with time.

24. Mechanical equilibrium No moment of matter in system with respect to its surrounding. Mechanical property remain constant.

25. Thermodynamic process An operation or transition by which a state of a system changes from initial state to final state. Type of process- Isothermal process(∆T=0) Isobaric process(∆P=0) Isochoric process(∆V=0) Adiabatic process(q=0) Reversible process Irreversible process

26. Isothermal process Temperature of the system constant. ∆T =0 In this process temperature at initial state and final state is constant. In this process system exchange heat energy with its surrounding to maintain constant temperature. Occur in close system. Internal energy of the system remains constant, hence ∆U = 0. In this process, gaseous system P, V of a change.

27. Reversible process A process carried out in such a manner that every stage, the driving force is only infinitesimally greater then the opposing force and it can be reversed by an infinitesimal increase in opposing force and the system exists in equilibrium with its surrounding throughout, is called a reversible process. Slow, infinite number of step.

28. Irreversible process Unidirectional process which proceeds in a definite direction and cannot be reversed at any stage and in which driving force and opposing force differ in a large magnitude. Also called spontaneous process. They are real process not hypothetical. Eg . Flow of heat from high T to lower T.

29. Work W = F.s Work is one of the ways by which a system can exchange energy with its surrounding by changing the state of the system. Example- Object move by applying force( object energy)

30. Kind of work in chemical thermodynamics The type of work is mechanical work i.e. pressure volume work. W = - P ex (V 2 – V 1 ) external pressure apply change in volume Work is also obtain due to chemical process or reaction.

31. Expression for pressure- volume work Ideal gas Massless, frictionless Piston. As the gas Expand it pushes Piston upward through Distance d against external force. d P ex P W = - F x d a Area of cross section

32. If ‘a’ is the cross section area of the cylinder or piston, then W = x d x a Now the pressure is P ex = F/a and ∆V = d x a. W = - P ex x ∆V   Expression for pressure- volume work

33. Sign convention of work during expansion and compression Expansion of a gas: P ex changing the volume from V 1 to V 2 . then ∆V = V 2 – V 1 . W = - P ex x ∆V. During expansion V2 > V1. work perform by surrounding. This result decrease energy of the system. Hence work is – ve i.e W is – ve .

34. Sign convention of work during expansion and compression Compression of a gas: P ex changing the volume from V 1 to V 2 . then ∆V = V 2 – V 1 . W = + P ex x ∆V. During expansion V2 < V1. work perform by system. This result increase energy of the system. Hence work is + ve i.e W is + ve .

35. Concept of Maximum work the process carried out at a constant temperature in the reversible manner by changing the state of the system through infinitesimally small steps in which driving force is infinitesimally greater then opposing force give maximum work. Is called an isothermal reversible process.

36. Process carried out at constant temperature. During the complete process, driving force is infinitesimally greater then opposing force. The work obtained is maximum. This is given W max = -2.303 nRT log 10 or log 10 . ∆U = 0. The heat absorbed irreversible manner q rev , is completely converted into work.   Concept of Maximum work

37. Condition of Maximum work In a thermodynamic process, maximum work is obtained from a system when All the changes taking places in it are thermodynamically reversible . Change in the state of the system take place in infinite no. of step . During change, driving force is infinitesimally greater then the opposing force.

38. Expression for Maximum work Ideal gas Massless, frictionless Piston. As the gas Expand it pushes Piston upward through Distance d against external force. P – dP W = - F x d V V+dV P dV V + dV – V = dV

39. Expression for Maximum work dW = -(P- dP ) dV dW = - PdV – dPdV dPdV = negligible. dW = - PdV W max = - 2.303 nRT log e    

40. Path dependence nature of work Work is not the property of the system. Not state funtion . W = - P (V 2 -V 1 ) A (V 1 ) B (V 2 ) W1 Expand Path 1 W2 Pex 2 W3 exapnsion reversibly 3

41. Concept of Heat Another way of exchanging energy system and surrounding. Not property of system. Not state funtion . Heat exchange only possible by path. Heat is path dependence. Eg . Rod heat transfer

42. Sign Convention of W and q Work and Heat are the form of energy. Due to work and exchange of heat, the energy of the system changes. +q = +W eg surrounding to system. -q = -W eg Gym. +q = heat absorbed and –q = heat released. +W = compressed. -W = expansion.

43. Unit of energy and Work Litre - atmosphere(L atm OR lit.atm ) Erg- W= dyne x cm = 1 erg. Force & distance. Calorie- heat energy Joule- amount of work.

44. Interconversion of work and energy W = 1 atm x 1 lit 1 atm = 1.013 x 10 5 Nm 2 1 lit = 10 -3 m 3 W = (1.013 x 10 5 x 10 -3 )Nm W = 101.3 J W = 24.22 cal W = 1.013 x 10 9 erg

45. Internal Energy[U] Total energy K.E and P.E present in the system. State function. Value depends on the state of a system. Change in internal energy, ∆U = U 2 – U 1 . Extensive property. Same unit as work and energy.

46. Total energy [U] = P.E + K.E. U total = U = U potential + U Kinetic U potential = U intramolecular + U intermolecular U Kinetic = U translational + U vibrational + U rotational + U electric U = U intra + U inter + U trans + U vib + U rota + U electric Internal Energy[U]

47. First law of thermodynamics Law of conservation of energy. 1 kind of energy consumed another kind of energy disappears. It is impossible to construct a perpetual motion machine. U = q + W Total amount of work is converted into heat energy.

48. Mathematical equation of 1 st Law of thermodynamic (V 2 , U 2 )Final state Initial state(V 1 , U 1 ) Heat absorb from surrounding

49. Due to volume change , the system perform the work W, hence total energy U 2 of the system in the final state is, U2 = U1 + q + W U2 – U1 = q + W ∆U = q + W For infinitesimally small change the mathematical expression is, dU = dq + dW Mathematical equation of 1 st Law of thermodynamic

50. First law of thermodynamic for various processes Isothermal process:- ∆T= 0 System depends on the temperature there is no change in the internal energy U of the system . Hence ∆U = 0. ∆U = q + W 0 = q + W +q(expansion) = -W or W = -q(consumed).

51. First law of thermodynamic for various processes Isobaric process:- ∆P = 0 System performs the work of expansion due to volume change . W = - P ex x ∆V q P heat absorb by the system at constant pressure. ∆U = q P + W ∆U = q P – P ex ∆V Or q p = ∆U + P ex ∆V q p heat absorbed used to increase the internal energy of the system.

52. Isochoric process ∆V = 0 Hence system doesn’t perform mechanical work.. W = - P∆V = 0. ∆U = q + W ∆U = q v q v = heat absorbed at constant volume. ∆U and q is state funtion .

53. Adiabatic process q= 0. ∆U = q + W ∆U = W ad System Expansion - ∆U decrease internal energy and temperature of system decrease. System Compression - ∆U increase internal energy and temperature of system increase.

54. Modern form of the first law of thermodynamic According to Einstein's theory, mass can be converted in to energy. Hence mass is also form of energy. The sum of mass and energy of an isolated system remain constant.

55. IUPAC sign convention of q, U and W For heat q:- +q = heat absorb by system. -q = heat loss by system. Heat energy left. For work W:- +W = work done on the system by compression. -W = expansion. Internal energy of system lose. For internal energy U:- +U = internal energy of system increase by absorption of heat. Similarly, –U.

56. Enthalpy H = U + PV Enthalpy represent total heat content of the system, at constant pressure. State function and extensive property. Absorption of heat by system increase its enthalpy. Hence enthalpy is called heat content of the system.

57. Expression of enthalpy change H 1 , U 1 ,P 1 ,V 1 H 2 ,U 2 ,P 2 ,V 2 H 1 = U 1 + P 1 V 1 & H 2 = U 2 + P 2 V 2 The enthalpy change ∆H is given by, ∆H = H 2 – H 1 ∆H = U 2 + P 2 V 2 – (U 1 + P 1 V 1 ) ∆H = U 2 – U 1 + P 2 V 2 – P 1 V 1 ∆H = ∆U + P∆V

58. Show that the heat absorbed at constant pressure is equal to the change in enthalpy of the system. By the first law of thermodynamic, ∆U = q + W q= ∆U – W If q p = heat absorbed at const. P. W = -P∆V q p = ∆U + P∆V. (∆H=∆U + P∆V) q p = ∆H Enthlpy also called heat content of the system.

59. Relation between ∆U & ∆H in Isochoric process:- ∆H = ∆U + P∆V ∆V = 0. ∆H = ∆U 2. Isobaric process:- ∆P =0 ∆H = ∆U

60. Derive the expression for the heat of reaction at Constant pressure- q= ∆U – W. ∆ H p = ∆U – W (W = - P∆V) ∆ H p = ∆U + P∆V q p = ∆ H p 2. Constant Volume- q= ∆U

61. Derive the relation ∆H=∆U+∆ nRT Consider a reaction in which n1 moles of gaseous reactant in initial state change to n2 moles of gaseous product in the final state. n 1 A (H 1 U 1 P 1 V 1 ) n 2 B(H 2 U 2 P 2 V 2 ) Enthalpy change, ∆H= H 2 - H 1 . H 1 = U 1 + P 1 V 1 & H 2 = U 2 + P 2 V 2 The enthalpy change ∆H is given by, ∆H = H 2 – H 1 ∆H = U 2 + P 2 V 2 – (U 1 + P 1 V 1 ) ∆H = U 2 – U 1 + P 2 V 2 – P 1 V 1 PV = n 1 RT

62. ∆H = ∆U + ∆ nRT If qp and qv are the heats involved in the reaction at constant pressure and volume, then q p = ∆H and q v = ∆U. q p = q v + ∆ nRT Derive the relation ∆H=∆U+∆ nRT

63. Expression for work done in a chemical reaction Chemical reaction depend upon the change in no. of gaseous moles of product to reactant. Consider, n1 V1 A n2 V2 B. Initial state, PV 1 = n 1 RT Final state, PV 2 = n 2 RT W = -P∆V = -∆ nRT . If n1=n2, W=0. If n2>n1, expansion W- ve If n2<n1, compression W+ve .

64. Enthalpies of physical change Phase Transitions:- change in physical state of matter. Type of phase changes. Fusion- solid to liquid state. Heat absorbed, endothermic (∆H>0). Eg . Ice to water. Vaporisation or evaporation- liquid to gas state. Heat absorbed, endothermic (∆H>0) Sublimation- solid to gas state. ∆H>0. eg Camphor . Note- Temperature and Pressure remain constant.

65. (1) Phase Transition = 40.7kJmol-1 At 0̊˚C = 6.01kJmol-1 At 0̊˚C = 40.7kJmol-1 At 100̊˚C

66. (2) Enthalpy of atomic or molecular changes A. enthalpy of ionization.

67. B. enthalpy of atomization- Dissociation of 1 mole gas substance into free gaseous atom. (2) Enthalpy of atomic or molecular changes

68. When one mole of a substance is dissolved in a large excess of a solvent, so that further dilution will not change the enthalpy at const. T and P. Eg . HCl (g) + aq → HCl ( aq ) ∆H = 4kJmol-1 (C) Enthalpy of Solution

69. (D) Enthalpy if dilution Solution – concentrated solution is diluted to form another concentrated solution.eg HCl + 50 H 2 O → HCl (50H 2 O), ∆H= -73.26kJ

70. Thermochemistry Study of heat change during chemical reaction. Heat of reaction:-(enthalpy of chemical reaction) ∆H = ∑ H product - ∑ H reactant Endothermic reaction- ∑ H product > ∑ H reactant absorption of heat , ∆H + ve . Exothermic reaction- ∑ H product - ∑ H reactant lose of heat, ∆H – ve .

71. Thermochemical equation A chemical reaction which is represent by Physical state of of all reactant and product. Reactant reacting to form product. A balanced equation Enthalpy change(heat given in or out during reaction) A + B → C + D ∆H= (+,- ) kJmol-1

72. IUPAC Guideline for writing thermochemical equations Physical state, balance chemical equation. Heat and enthalpy changes are measured at STP. 298K & 1atm. ∆H written on R.H.S. Proper sign must be indicate +H, -H. Enthalpy of the element in their STD states is taken as zero. STD shown by H˚. Allotropic form must mentioned.eg C graphite . For reverse rxn , ∆H value same but sign change.

73. Standard Enthalpy of reaction The difference between the sum of enthalpy of products and reactant with every substance in its standard state at constant T(298K) and P(1atm). ∆H˚ = ∑ H˚ product - ∑ H˚ reactant Reactant → product

74. Standard enthalpy of formation or standard heat of formation(∆ f H ˚)