1 / 19

Chapter 19 Principles of Reactivity: Entropy & Free Energy

Chapter 19 Principles of Reactivity: Entropy & Free Energy. Spontaneity Spontaneous processes Factors affecting spontaneity Entropy Entropy changes The third law of thermodynamics Calculating entropy changes in chemical reactions The second law of thermodynamics Gibbs Free Energy

gzifa
Download Presentation

Chapter 19 Principles of Reactivity: Entropy & Free Energy

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 19 Principles of Reactivity: Entropy & Free Energy

  2. Spontaneity • Spontaneous processes • Factors affecting spontaneity • Entropy • Entropy changes • The third law of thermodynamics • Calculating entropy changes in chemical reactions • The second law of thermodynamics • Gibbs Free Energy • Definition • Calculating free energy changes in chemical reactions • Standard free energy of formation • Free energy and temperature • Free energy and equilibrium

  3. I. Spontaneity • A. Spontaneous process • = one that, once started, continues on its own without input of energy • nonspontaneous = needs continual input of energy • Chemical Thermodynamics • Chapter 6: Thermochemistry • heat of reaction, DH • first law: conservation of energy • Thermodynamics: energy changes • entropy (DS), free energy (DG) • second and third laws •  spontaneity of reaction, position of equilibirum

  4. I. Spontaneity B. Factors affecting spontaneity • 1. decrease in potential energy (DH < 0) • apple falling • water flowing downhill • 2H2 + O2 2H2O • HCl + NaOH  NaCl + H2O • thermite • 2. increase in disorder (tendency toward higher probability) • deck of cards (6 x 109 combinations) • gas molecules in a container • salt dissolving in water • ice melting, water evaporating

  5. A. Entropy changes Entropy change, DS = Sf - Si • 1. examples • expansion of a gas DS • dissolution of a solid DS • H2O(s)  H2O(l)  H2O(g) DS • H2O(g)  H2O(l)  H2O(s) DS • H2CO3  CO2+ H2O DS • 2H2 + O2  2H2O DS • CaO + SO2  CaSO3DS remove partition H2O NaCl(s) NaCl(aq) II. Entropy, S = degree of disorder (probability)

  6. q DS DS= 1 T q T 10 K  20 K DS1 DS1 > DS2 q 300 K  310 K DS2 II. Entropy A. Entropy changes 2. calculating DS for fusion and vaporization DS q increase KE  increase in disorder e.g., For H2O(s)  H2O(l), DHfus = +6.02 kJ/mol at 0ºC and for H2O(l)  H2O(g), DHvap = +40.7 kJ/mol at 100ºC. What are DSfus and DSvap?

  7. increase in S II. Entropy B. The third law of thermodynamics • “For a pure, crystalline substance at 0 K, S = 0” • Increase T increase in: translational motion • vibrational motion • rotational motion • Standard state entropy, Sº • = entropy of a substance at 25ºC, 1 atm (Table 19.1, Appendix L) (Sº  0 for elements!)

  8. (note units) II. Entropy C. Calculating entropy changes in chemical reactions • Standard entropy change, DSº • DSºrxn = Sºproducts - Sºreactants • e.g., 2H2 + O2 2H2O • DSºrxn = [2Sº(H2O)] - [2Sº(H2) + Sº(O2)] • = [2(188.8)] - [2(130.6) + (205.0)] • = -88.6 J/K

  9. II. Entropy C. Calculating entropy changes in chemical reactions • e.g., Calculate the standard entropy change for the reaction, • 2Al(s) + 3Cl2(g)  2AlCl3(s) Sº(Al) = 28.3 J/mol·K • Sº(Cl2) = 223.1 J/mol·K • Sº(AlCl3) = 110.7 J/mol·K

  10. universe system surroundings II. Entropy D. The second law of thermodynamics “In any spontaneous process, the entropy of the universe must increase (DSuniverse > 0). DSuniverse = (DSsystem + DSsurroundings) > 0 e.g., erecting a building 6CO2 + 6H2O  C6H12O6 + 6O2 CFCs

  11. DHsystem T DHsystem T qsystem T DSsurroundings = – = – DSuniverse = DSsystem – DGsystem = DHsystem – TDSsystem • Gibbs free energy • state function • relates spontaneity to heat of reaction (DH) and entropy (DS) • for a process to be spontaneous, DSuniverse > 0 or DGsystem < 0 III. Gibbs Free Energy A. Definition DSuniverse = DSsystem + DSsurroundings TDSuniverse = DHsystem – TDSsystem

  12. Reactants Products G Equilibrium III. Gibbs Free Energy A. Definition DG < 0: exergonic, spontaneous DG > 0: endergonic, nonspontaneous DG = 0: system is at equilibrium (no free energy remains) DGrxn < 0 “spontaneous” R  E DG < 0 spontaneous P  E DG < 0 spontaneous E  R DG > 0 nonspontaneous E  P DG > 0 nonspontaneous

  13. III. Gibbs Free Energy B. Calculating free energy changes in chemical reactions e.g., C(s) + ½O2(g) + 2H2(g)  CH3OH(l) at 25º DHºrxn = DHfº(products) – DHfº(reactants) = –238.7 kJ/mol DSºrxn = Sºproducts – Sºreactants = [Sº(CH3OH)] – [Sº(C) + ½Sº(O2) + 2Sº(H2)] = [126.8] – [(5.7) + ½(205.1) + 2(130.7)] = –242.9 J/mol·K DGºrxn = DHºrxn – TDSºrxn = (–238.7 kJ/mol) – (298 K)(–242.9 J/mol·K) = –166.3 kJ/mol = DGfº(CH3OH)

  14. standard free energies of formation • (Table 19.3, Appendix L; note: DGfº(element) = 0) III. Gibbs Free Energy C. Using standard free energies of formation DGºrxn = DGfº(products) - DGfº(reactants) e.g., CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) DGºrxn = [DGfº(CO2) + 2DGfº(H2O)] - [DGfº(CH4)] = [(-394.4) + 2(-237.1)] - [(-50.7)] = -817.9 kJ

  15. III. Gibbs Free Energy C. Using standard free energies of formation • e.g., If DGºrxn for the following reaction is -272.8 kJ, what is DGºf for TiCl2? • TiCl2(s) + Cl2(g)  TiCl4(s) (DGºf(TiCl4) = -737.2 kJ/mol)

  16. H2O • NH4NO3(s)  NH4+(aq) + NO3–(aq) (gets cold) • DH > 0, DS > 0  DG < 0 at higher temperatures III. Gibbs Free Energy D. Free energy and temperature DG = DH – TDS DH DS DG – + – at all temperatures + – + at all temperatures + + – at higher temperatures, + at lower temperatures – – – at lower temperatures, + at higher temperatures • e.g., C12H22O11 + 12O2 12CO2 + 11H2O (produces heat) • DH < 0, DS > 0  DG < 0 at any temperature • (for reverse reaction, DG > 0 at any temperature; requires input of energy)

  17. III. Gibbs Free Energy D. Free energy and temperature • e.g., CH3–CH3 CH2=CH2 + H2DHº = +137 kJ • DSº = + 128 J/mol·K • at 25ºC, DGº = +98.9 kJ nonspontaneous • Above what temperature does the reaction become spontaneous?

  18. DGº = –RTlnK G  R P R P R P DGº > 0 DGº < 0 DGº = 0 III. Gibbs Free Energy E. Free energy and equilibrium DG = DGº + RTlnQ At equilibrium, DG = 0, Q = K 0 = DGº + RTlnK DGº < 0  K > 1 equilibrium lies to the right DGº > 0  K < 1 equilibrium lies to the left DGº = 0  K = 1 equilibrium lies in the middle

  19. e.g., 2SO2+ O2 2SO3 What is K at 25ºC? DGºf(SO2) = -300.2 kJ/mol DGºf(SO3) = -371.1 kJ/mol III. Gibbs Free Energy E. Free energy and equilibrium DGº = –RTlnK or K = e–Gº/RT

More Related