Unit 14: Chemical Thermodynamics

1. CHM 1046: General Chemistry and Qualitative Analysis Unit 14:Chemical Thermodynamics Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL • Textbook Reference: • Chapter # 15 (sec.12-17) • Module # 3 (sec. VIII-XII)

2. First Law of Thermodynamics • The law of conservation of energy: energy cannot be created nor destroyed. (James Joule in 1843 ) • Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa. E = q + w • Therefore, the total energy of the universe is a constant. w = PV Esys + Esurr = 0 E = q + PV Esys = -Esurr E = q + RT

3. Spontaneity Second Law of Thermodynamics First Law of Thermodynamics E1 Stone Do all processes that loose energy occur spontaneously (by themselves, without external influence)?????? + Work E2 - (work +heat) E = E2 – E1

4. Spontaneous Processes {Spontaneity} • can proceed without any outside intervention. • Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

5. Spontaneous Processes • 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. Spontaneous @ T < 0ºC Is the spontaneity of melting ice dependent on anything? Spontaneous @ T > 0ºC

6. Spontaneity vs. Speed C diamond C graphite Thermodynamics vs. Kinetics

7. Irreversible Processes • Irreversible processes cannot be undone by exactly reversing the change to the system. E1 Stone • Heat energy is lost to dissipation and that energy will not be recoverable if the process is reversed. + Work • Spontaneous processes are irreversible. Reversible Processes • 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. E2 - (work + heat)

8. q T Entropy (S) Entropy is a measure of the energy that becomes dissipated and unavailable (friction, molecular motion = heat). • Entropy (S) is a term coined by Rudolph Clausius in the 1850’s. Clausius chose "S" in honor of Sadi Carnot(whogave the first successful theoretical account of heat engines, now known as the Carnot cycle, thereby laying the foundations of the second law of thermodynamics). • Clausius was convinced of the significance of the ratio of heat delivered and the temperature at which it is delivered, Entropy (S) =

9. Entropy (S) Gas ENTROPY • Entropy can be thought of as a measure of the randomness (disorder) of a system. • It is related to the various modes of motion in molecules. Liquid • Like total energy, E, and enthalpy, H, entropy is a state function. • Therefore, • S = SfinalSinitial Solid {Entropy.WaterBoiling}

10. Second Law of Thermodynamics • the entropy of the universe increases for spontaneous (irreversible) processes. Suniv = Ssystem + Ssurroundings > 0 Suniv = Ssystem + Ssurroundings = 0 • the entropy of the universe does not change for reversible processes.

11. Second Law of Thermodynamics All spontaneous processes cause the entropy of the universe to increase. So what is our fate as a result of the second law operating in our Universe? ENTROPIC DOOM!

12. Entropy on the Molecular Scale • Molecules exhibit several types of motion (Kinetic energies): • Translational: Movement of a molecule from one place to another. • Vibrational: Periodic motion of atoms within a molecule. • Rotational: Rotation of the molecule on about an axis or rotation about  bonds. • Boltzmann envisioned the motions of a sample of molecules at a particular instant in time. • This would be akin to taking a snapshot of all the molecules. • He referred to this sampling as a microstate (W) of the thermodynamic system • Entropy is ……. • S = klnW • …..where k is the Boltzmann constant, 1.38  1023 J/K.

13. Entropy on the Molecular Scale S = klnW …..where k is the Boltzmann constant, 1.38  1023 J/K. • The number of microstates (W) and, therefore, the entropy (S) tends to increase with increases in which variables…. • Temperature (T). • Volume (V). • The number of independently moving molecules ().

14. Entropy Changes • Entropy increases with the freedom of motion of molecules. • S(g) > S(l) > S(s) • In which of the following does Entropy increase & WHY?……. • Gases are formed from liquids and solids. • Liquids or solutions are formed from solids. • The number of gas molecules (or moles) increases. Heat H2O (l) H2O(g) {*Entropy&PhaseOfMatter} H2O CaCl2 (s) Ca 2+(aq) + 2Cl-(aq) {*EntropySolutions.KMnO4(aq)} Electricity 2 H2O (l) 2 H2 (g) + O2(g) 2 C8H18 (l) + 25 O2 (g) 16 CO2(g) + 18 H2O(g)  gas= 34-25 = +9/2 = 4.5 C8H18

15. Third Law of Thermodynamics The entropy (S) of a pure crystalline substance at absolute zero (-273°C) is 0.

16. Standard Entropies • Standard entropies tend to increase with increasing molar mass. • Larger and more complex molecules have greater entropies (greater ways to execute molecular motions) {*Entropy&MolecuarSize} {Entropy&Temp}C7H15 @ 500 K S=921J/nK vs, @ 200 K

17. Standard Entropies (298 K) from Absolute Entropies (0K) Absolute Entropy (S) @ - 237°C (0 K), S = 0 Standard Entropy (S˚) @ 25°C (298 K), S = ???? Liquid Gas Solid q = mcT H°vap S q = mcT S° H°fus Calculate the sum of all the infinitesimally small changes in entropy as T varies from T=0  T= 298, by taking its Integral. q = mcT Temp (K) 298

18. Entropy Changesin the System S°syst = S°rxn @ T Entropy changes for a reaction (=system) can be estimated in a manner analogous to that by which H is estimated: where n and m are the coefficients in the balanced chemical equation.

19. Entropy Changesin the System Problem: Calculate the standard entropy changes for the following reaction at 25oC. N2(g) + 3 H2(g) 2 NH3(g) S° = nS°(prod) - mS°(react) 2(192.5) – [(191.5)+3(130.6)] S° = - 198.3 J/

20. DGrxn = S DGf (products)  S Gf (reactants) Thermodynamic Changes in Systems (Chem. Reactions) Hrxn =  Hf (products) - Hf (reactants) ☺ Appendix 1 (CHM 1046 Module): notice S° is in J not kJ.

21. Hsys T (qsys) T Ssurr= Ssurr= Entropy Changes in the Surroundings What in a chemical reaction causes entropy changes in the surroundings? • Heat (q) that flows into or out of the system changes the entropy of the surroundings: Ssurr∝- (qsys) • For an isothermal process: q q q q System • At constant pressure, qsys is simply H for the system. q q q Surroundings

22. -Hsys T Ssurr = Entropy Change in the Universe Problem: Calculate the Suniv for the synthesis of ammonia @ 25oC. N2(g) + 3 H2(g) 2 NH3(g)H°rxn = - 92.6 kJ/mol Suniv = Ssyst or rxn + Ssurr nS(prod) - mS(react) 2(192.5) – [(191.5)+3(130.6)] S°syst = - 199 J/K·mol Ssurr= 311 J/K·mol Suniv = - 198.3 J/K·mol+ 311 J/K·mol Suniv = 113 J/K·mol

23. Hsystem T -Hsys T Ssurr = Entropy Change in the Universe Suniv = Ssyst or rxn + Ssurr • Since: • Then: Suniv = Ssyst + Multiplying both sides by T, TSuniv = TSsyst + Hsyst TSuniv = Hsyst  TSsyst TDSunivis defined as the Gibbs (free) Energy, G. J. Willard Gibbs USA, 1839-1903

24. Gibbs Free Energy (G) ☺ TSuniv = DGuniv = DHsys  TSsys • WhenSuniv is positive, G is negative. • When G is negative, the process is spontaneous. Gibbs Energy (-TΔS) measures the "useful" or process-initiating work obtainable from an isothermal, isobaricthermodynamic system. Technically, the Gibbs free energy is the maximum amount of non-expansion work which can be extracted from a closed system or this maximum can be attained only in a completely reversible process.

25. Free Energy Changes At temperatures other than 25°C, DG° = DH  TS How does G change with temperature? ☺ • There are two parts to the free energy equation: • H— the enthalpy term • TS — the entropy term • The temperature dependence of free energy, then comes from the entropy term.

26. Spontaneity: Enthalpy & Entropy DG° = DH  TS Spontaneous @ all T NonSpontaneous @ all T Spontaneous @ low T Spontaneous @ high T

27. Spontaneity: Enthalpy & Entropy DG = DH(TS) {Entropy Driven Reactions} Entropy H2O Enthalpy NH4NO3(s) NH4+(aq) + NO3-(aq) S = + H = + (-TS) {Enthalpy Driven Reaction} 2 H2(g) + O2 (g) 2 H2O(g) H = - S = - n = 2-3 = -1 (+TS) {EntropySyst+Surr.FormationOfWater} {Entropy & Enthalpy Driven Reaction} H = - S = + Na2CO3(s) + HCl(aq) NaCl(aq) + CO2 (g) (-TS)

28. Problems DG° = DHT(S) (-763) – (-804) +41 (.3549) – (.2219) +.1330 Product Reactant (-T) TiCl4(l)  TiCl4(g) DG° = DH  TS 0 = (131.3kJ)  T(.133kJ) T = 987

29. DG = SnDG(products)  SmGf (reactants) f Standard Free Energy Changes Analogous to standard enthalpies of formation are standard free energies of formation, G. f where n and m are the stoichiometric coefficients.

30. DGrxn = SnDG(prod)  SmG(react) f Standard Free Energy Changes 2 C6H6 (l) + 15 O2 (g) 12 CO2(g) + 6 H2O(g) Calculate the standard free energy changes for the above reaction @ 25 °C. f [12 CO2(g) + 6 H2O(g)] – [2 C6H6 (l) + 15 O2(g)] [12(-394) + 6(-229)]– [2(125) + 15 (0)] Grxn = - 6352 J/mol · K

31. Fourth Law of Thermodynamics: Emergence Complex emergent systems spontaneously (-G) arise when energy flows through a collection of many interacting particles, resulting in new patterns of complex behaviors that are much more than the sum of the individual parts (+S) The formation of these complex patterns in emergent systems is more efficient in the dissipation of energy (-H), thus speeds up the increase of entropy in the universe. G = H-TS A precise definition of emergence and a useful mathematical formulation of this phenomenon remains elusive. C  f [n, i, E(t)] Examples: cells forming living organisms; stars forming galaxies; neurons forming conscious brain.

38. 2002 B

41. 2003 A

43. 2004 A

46. 2004 B

48. 2005 A

50. 2006 (B)