THERMODYNAMICS

1. THERMODYNAMICS • SPECIFIC LEARNING OBJECTIVE • At the end of the session the student should be able to explain: • Energy • The first law of thermodynamics • Entropy • Free energy

2. Thermodynamic In Chemistry System In Living System

3. THERMODYNAMICS Thermodynamic is the law that formulated from observation on conversion of energy from one form to the other. i.e. transduction

4. What is the energy? 1. Energy is a much used term, but it represents a rather abstract concept. 2. Energy is usually defined as the capacity to do work. 3. Chemist define work as directed energy change resulting from a process

5. The type of Energy 1. Kinetic energy 2. Radiant energy 3. Thermal energy 4. Chemical energy 5. Potential energy

6. Definition of type energy 1. Kinetic energy – the energy produced by a moving object 2. Radiant energy : comes from the sun (solar energy) and is Earth’s primary energy source. Solar energy heats the atmosphere and Earth’s surface, stimulates the growth of vegetation through the process known as photosynthesis, and influences global climate patterns. continued

7. The Activated Complex form the reaction: A + B AB# P

8. Definition of type energy 3.Thermal energy is the energy associated with the random motion of atoms and molecules. 4. Chemical energy is stored within the structural units of chemical substances; its quantity is determined by the type and arrangement of atoms in the substance being considered. 5. Potential energy is energy that is also available by virtue of an object’s position.

9. Conclusion of Energy • All forms of energy can be interconverted (at least in principle) from one form to another

10. Scientists have concluded that energy can be neither destroyed nor created. • Thermodynamic Law

11. THERMODYNAMICS Work and heat are not state functions Thermodynamic I is the law of conservation of energy Thermodynamic I Work Heat Thermodynamic II Thermodynamic III

12. The relationship between chemical energy and other forms of energy, with examples.

13. Energy change in chemical reactions Almost all chemical reactions absorb or produce (release) energy, generally In the form of heat. • Heat is the transfer of thermal energy between two bodies that are at different temperatures. Although ”Heat” itself implies the transfer of energy, we customarily talk of ”heat absorbed” or ”heat released” when describing the energy changes that occur during a process.

14. Energy changes associated with chemical reactions System Surroundings

15. SYSTEM AND SURROUNDING Systemwe mean that the part of the world we are investigating. Surrounding we mean everything else Three type of systems be Surrounding Open Close Isolated system

16. System be open • Two of this examples are the examples of open system: • e.g. in the living organism, which takes up nutrients, releases the waste products, and generates work and heat. • An example in body, the body takes up nutrient, and then release urine which contains toxin, carbon dioxide, and so on.

17. System be closed Example of close system: • An example of close system is living of an microorganism, it was sealed inside a perfectly insulated box, it will, together with the box, constitute a closed system.

18. HEAT q, to be the manner of energy transfer that results from a temperature difference between the system and its surrounding Positive and negative sign of heat Heat input to a system is considered a positive quantity Heat evolved by a system is considered a negative quantity.

19. W O R K w, to be the transfer of energy between the system of interest and its surroundings as a result of existence of unbalanced forces between two. Positive and negative sign of work if the energy of the system is decreased by the work, or the system does work on the surroundings, or that work is done by the system, and we take it to be a negative quantity If the energy of the system is increased by the work, we say that work is done on the system by surroundings, and we take it to be a positive quantity

20. The effect of work is equivalent to the raising or lowering of mass in the surroundings. • Work is done by the system because the mass is raised work is done on the system because the mass is lowered.

21. E N E R G Y Energy is a state function It is a property that depends only upon the state of the system, and not upon how the system was brought to that state, or upon the history of the system. Thermodynamic I study of conservation of energy

22. The first law of thermodynamic • ∆ U = q + wwhich is essentially a statement of the law of conservation of energy. Where : 1. The term ∆ U represents the change of internal energy of the system, 2. q is the thermal energy (heat) added to the system, and w is the work done on the system.

23. The chemical reactions that need energy 1. The Photoelectric Effect  this is mystery in physics. Experiments had already demonstrated that electrons were ejected from the surface of certain metals exposed to light of at least a certain minimum frequency. Einstein suggested that a beam of light is a stream of particles. These particles of light are called photons. Using Planck’s quantum theory of radiation as a starting point, Einstein deduced that each photon must possess energy E, given by the equation : E = hv In which v is the frequentcy of light and h is Planck’s constant

24. The equation of E = hv E = hv E = KE + BE  hv = KE + BE in which : * KE is the kinetic energy of the ejected electron and * BE is the binding energy of the electron in the metal

25. The energies that the electron in the hydrogen atom • En = – RH • In which RH, the Rydberg constant, has the value 2.18 x 10-18 J • The number n is an integer called the principal quantum number; it has the value n = 1, 2, 3,…. 1 n2

26. Strength of Covalent Bond The Strength of Covalent Bond is defined by the amount of energy needed to break it. A quantitative measure of stability of a molecule is its bond dissociation energy (or bond energy). For example: H2(g)  H(g) + H(g) H = 436.4 kJ HCl(g)  H9g) + Cl(g) H = 431.9 kJ

27. Covalent Bond in Organic Compounds STRUCTURE SATURATED:  Bonding : are formed by overlap of two atomic orbitals, each of which contains one electron UNSATURATED : and  bonding.  bonding   (pi) bond

28. ENTHALPY • The enthalpy of a system, which has the symbol H, is that of Heatcontent (heat of reaction) and is measure of the change in total bonding energy during a reaction. It is defined mathematically as :H = U + PV; H is a function of state The standard enthalpy change for any reaction (∆H0rxn)can determine by using standard enthalpies of formation (∆H0f) and Hess’s Law

29. CONSTANT PRESSURE PROCESSES • Most processes occur in the open at one atmosphere pressure.In these cases, P1 = P2 = P, say, and • ∆H = ∆U + P ∆V

30. Positive and negative sign of Enthalpy • ∆ H has a negative sign for an exothermic change (heat isreleased), is mean the bonds in the products are stronger (more stable) than the bonds in the reactants. • ∆ H has a positive sign for an endothermic change (heat is absorbed), is mean the bonds in the products are weaker (less stable) than the bonds in the reactants.

31. HESS’S LAW The principle of constant heat summation, often known as Hess’s Law, is thus seen to lead directly from the fact that H is a function of state. Hess’s Law be valid for : ∆r H˚ or ∆f H˚, ˚ = all reactants and products are in • their standard states. f = formation standard enthalpies of formation Pº = 1 atmosphere, and temperature 25ºC or 298.15°K

32. This idea is immensely powerful, because it enables Hº298 values to be determined for any reaction, as long as the H of formation are known for each reactant and product. • ∆rH = H prod– H react

33. Example No. 1 : • Consider the following two chemical equations. • 1. C(s) + ½ O2 (g) CO (g) ∆rH (1) = -110.5 kJ 2. CO (g) + ½ O2 (g) CO2 (g) ∆rH (2) = -283.0 kJHow many Joule ∆rH (3) = ….? For below equation C (s) + O2 (g) CO2 (g) ∆rH (3) = ...?

34. Example No. 2 : 2 P(s) + 3Cl2(g) 2 PCl3(l) ∆rH (1)= -640 kJ 2 P(s) + 5 Cl2(g) 2 PCl5(s) ∆rH (2)= -887 kJ Please calculate the value of ∆rH for below equation PCl3(l) + Cl2(g) PCl5(s) ∆r H (3) = .....? • Please you make the application of Hess’s Law, consider the use of

35. solution No. 2: 2 P(s) + 3Cl2(g) 2 PCl3(l) ∆rH (1)= -640 kJ 2 P(s) + 5 Cl2(g) 2 PCl5(s) ∆rH (2)= -887 kJ Please calculate the value of ∆rH for below equation PCl3(l) + Cl2(g) PCl5(s) ∆r H (3) = .....? • Please you make the application of Hess’s Law, consider the use of A. – 247 kJ C. – 124 kJ E. – 1527 kJ. B. + 247 kJ D. + 124 kJ

36. Spontaneous Changes The process tends to occur or not 1. The towards minimization of energy is one such directing influence, but there is also a tendency for material to become more physically disorganized. Two driving forces in nature 2. The tendency for entropy to increase is nature’s second driving force.

37. ENTROPY • The symbol of entropy = Sa thermodynamic function of state

38. NATURAL OR IRREVERSIBLE PROCESS • The entropy of system and surroundings together increases during all natural or irreversible process; • ∆ Ssystem + ∆ Ssurrounding = ∆ Suniverse> 0 REVERSIBLE PROCESS • For reversible process, the total entropy is unchanged; • ∆ Ssys + ∆ Ssur = ∆ Suniverse = 0

39. CYCLIC PROCESSES • For a cyclic process, a process in which the final state is the same as the initial state, ∆S = 0

40. Changes of entropy with temperature • ∆ S = S2 – S1 = CP ln T2/T1 (P constant) • ∆ S = S2 – S1 = CV ln T2/T1 (V constant)

41. Absolute entropy The third Law of Thermodynamics All truly perfect crystals at absolute zero temperature have zero entropy.

42. FREE ENERGY

43. Gibbs Free Energy The Gibbs energy determines the direction of a Spontaneous Process for a System at Constant Pressure and Temperature G is function of state Gibbs free energy, G,. It is a function of state which provides possible or not a change of any kind will tend to occur.

44. The value of ∆ G • For a favorable reaction, ∆G has a negative value, meaning that energy is released to the surroundings  Exergonic • For a unfavorable reaction, ∆G has a positive value, meaning that energy is absorbed from the surroundings  Endergonic

45. REACTION AT CONSTANT TEMPERATURE & PRESSURE dG ≤ 0 (constant T and P) The quantity G is called the Gibbs energy

46. Value of G in a system at constant T and P • The Gibbs energy will decrease as the result of any spontaneous processes until the system reaches equilibrium, where d G = 0.

47. The Gibbs free energy is defined as: • G = H - TS

48. RELATIONSHIP BETWEEN THE PROCESSES WITH GIBB’S FREE ENERGY Spontaneous processes, that is, those with negative ∆ G values, are said to be exergonic; they can be utilized to do work. Processes that are not spontaneous, those with positive ∆ G values, are termed endogonic; they must be driven by the input of free energy. Processes at equilibrium, those in which the forward and backward reactions are exactly balance, are characterized by ∆ G = 0.

49. Thermodynamic In Chemistry System In Living System

50. What was The Thermodynamic Studied in Living System? • Thermodynamic In Living System • 1. ∆ H (heat) • 2. ∆ S (the extent of disorder of the • system) • 3. ∆ G (Gibbs change in free energy that • proportion of the total energy change • in a system, that is available for doing • work)