1 / 59

Chemistry 102(01) Spring 2014

Chemistry 102(01) Spring 2014. Instructor: Dr. Upali Siriwardane e-mail : upali@latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F 9:30 - 11:30 am. or by appointment.  Test Dates : 9:30-10:45 am., CTH 328.

makan
Télécharger la présentation

Chemistry 102(01) Spring 2014

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. Chemistry 102(01) Spring 2014 Instructor: Dr. UpaliSiriwardane e-mail: upali@latech.edu Office: CTH 311 Phone257-4941 Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F9:30 - 11:30 am. or by appointment.  Test Dates:9:30-10:45 am., CTH 328 March 31, 2014 (Test 1): Chapter 13 April 23, 2014 (Test 2): Chapter 14 &15 May 19, 2014 (Test 3) Chapter 16 &17 May 21, 2014 (Make-up test) comprehensive: Chapters 13-17

  2. Chapter 6. Thermochemistry 6.1 Chemical Hand Warmers 231 6.2 The Nature of Energy: Key Definitions 232 6.3 The First Law of Thermodynamics: There Is No Free Lunch 234 6.4 Quantifying Heat and Work 240 6.5 Measuring for Chemical Reactions: Constant-Volume Calorimetry246 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure 249 6.7 Constant-Pressure Calorimetry: Measuring 253 6.8 Relationships Involving 255 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation 257 6.1 0 Energy Use and the Environment 263

  3. Chapter 17. Free Energy and Thermodynamics 17.1 Nature’s Heat Tax: You Can’t Win and You Can’t Break Even 769 17.2 Spontaneous and Nonspontaneous Processes 771 17.3 Entropy and the Second Law of Thermodynamics 773 17.4 Heat Transfer and Changes in the Entropy of the Surroundings 780 17.5 Gibbs Free Energy 784 17.6 Entropy Changes in Chemical Reactions: Calculating 788 17.7 Free Energy Changes in Chemical Reactions: Calculating 792 17.8 Free Energy Changes for Nonstandard States: The Relationship between and 798 17.9 Free Energy and Equilibrium: Relating to the Equilibrium Constant (K)

  4. What forms of energy are found in the Universe? mechanical thermal electrical nuclear mass: E = mc2 others yet to discover

  5. What is 1st Law of Thermodynamics Eenergy is conserved in the Universe All forms of energy are inter-convertible and conserved Energy is neither created nor destroyed.

  6. What exactly is DH? Heat measured at constant pressure qp Chemical reactions exposed to atmosphere and are held at a constant pressure. Volume of materials or gases produced can change.

  7. What is the internal energy change (DU) of a system? DU is part of energy associated with changes in atoms, molecules and subatomic particles Etotal= Eke + E pe + DU DU = heat (q) + w (work) DU = q + w DU = q -P DV; w =- P DV

  8. How is Internal Energy, DU measured? Heat measured at constant volume qv Chemical reactions take place inside a closed chamber like a bomb calorimeter. Volume of materials or gases produced can not change. ie: work = -PDV= 0 DU = qv + w qv = DU + o; w = 0 DU = qv = DU(internal energy )

  9. Enthalpy Heat changes at constant pressure during chemical reactions Thermochemical equation. eg. H2 (g) + O2 (g) ---> 2H2O(l) DH =- 256 kJ; DH is called the enthalpy of reaction. if DH is + reaction is called endothermic if DH is - reaction is called exothermic

  10. Entropy, S The thermodynamic property related to randomness is ENTROPY, S. Product-favored processes: final state is more DISORDERED or RANDOM than the original. Spontaneity is related to an increase in randomness. Reaction of K with water

  11. Physical Process” S[H2O(l)] > S[H2O(s)] at 0° C.

  12. Gibbs Free Energy, G DGo = DHo - T DSo 1. If DHis negative it helps product to be favored • 2. If DSis positiveit helps product to be favored • 3. If DGis negative reaction is product favored Gibbs free energy change = difference between the enthalpy of a system and the product of its absolute temperature and entropy predictor of spontaneity Total energy change of the system - energy lost in disordering the system

  13. Thermodynamics Standard States • The thermodynamic standard stateof a substance is its most stable pure form under • standard pressure (1 atm) and at some specific temperature (25 ºC or 298 K) • superscript circle is used to denote a thermodynamic quantity that is under standard state conditions: • ΔH = ΔH°ΔS = ΔS°ΔG = ΔG°

  14. Standard Thermodynamic Data ΔHof - Standard Enthalpy of Formation ΔGof - Standard Free Energy of Formation So - Standard Free Energy of Formation

  15. Standard Molar Entropy Values

  16. Chemical Thermodynamics spontaneous reaction – reaction which proceed without external assistance once started chemical thermodynamics helps predict which reactions are spontaneous

  17. Thermodynamics Will the rearrangement of a system decrease its energy? If yes, system is favored to react — a product-favoredsystem. Most product-favored reactions are exothermic. Often referred to as spontaneousreactions. “Spontaneous” does not imply anything about time for reaction to occur. Kinetic factors are more important for certain reactions.

  18. 1) Give the definitions of the following: • a) Enthalpy (H): • b) Enthalpy change of a thermo-chemical reaction (DH): • c) Entropy of a substance (S): • d) Entropy change of a chemical reaction(DS): • e) Thermodynamic Standard State(0):

  19. Laws of Thermodynamics Zeroth: Thermal equilibrium and temperature First: The total energy of the universe is constant Second : The total entropy (S) of the universe is always increasing Third : The entropy(S) of a pure, perfectly formed crystalline substance at absolute zero is zero

  20. 2) Give the definitions of the following: • a) ZerothLaw of thermodynamics: • b) First Law of thermodynamics: • c) Second Law of thermodynamics: • d) Third Law of thermodynamics:

  21. universe system surroundings Why is it necessary to divide Universe into System and Surrounding Universe = System + Surrounding Boundary?

  22. Types of Systems • Isolated system • no mass or energy exchange • Closed system • only energy exchange • Open system • both mass and energy exchange

  23. Why is it necessary to divide Universe into System and Surrounding Universe = System + Surrounding

  24. 3) Why we need to divide universe into surroundings and system for thermodynamic calculations? • Give the signs of the DH (heat) and DS (disorder) and DG ( free energy) when system lose or gain them. • Loss Gain • DH (heat) • DS (disorder) • DG ( free energy)

  25. Second Law of Thermodynamics In the universe the ENTROPY cannot decrease for any spontaneous process The entropy of the universe strives for a maximum in any spontaneous process, the entropy of the universe increases for product-favored process DSuniverse = ( Ssys + Ssurr) > 0 DSuniv = entropy of the Universe DSsys = entropy of the System DSsurr = entropy of the Surrounding DSuniv = DSsys + DSsurr

  26. Entropy of the Universe DSuniv = DSsys + DSsurr DsunivDSsysDSsurr + + + + +(DSsys>DSsurr) - + + (DSsurr>DSsys)

  27. 4) Explain the ways that DS of the universe, DSuniv could be +.

  28. Entropy and Dissolving

  29. 5) Assign a sign to the entropy change for the following systems. • a) mixing aqueous solutions of NaCl and KNO3 together: • b) spreading grass seed on a lawn: • c) raking and bagging leaves in the fall: • shuffling a deck of cards: • e) raking and burning leaves in the fall:

  30. Expansion of a Gas The positional probability is higher when particles are dispersed over a larger volume Matter tends to expand unless it is restricted

  31. Gas Expansion and Probability

  32. S (gases) > S (liquids) > S (solids) Entropies of Solid, Liquidand Gas Phases

  33. 6) Taking following examples explain how disorder is related to a measuring positional probability) or dispersion among the allowed energy states? • a) Expansion of gases: Two gas molecules trapped in two vessels with a tube with a stop cock.

  34. 6) Taking following examples explain how disorder is related to a measuring positional probability) or dispersion among the allowed energy states. • b) Distribution of Kinetic energy at 0, 25 and 100°C for O2

  35. Entropy and Molecular Structure

  36. Entropy, S Entropies of ionic solids depend on coulombic attractions. So (J/K•mol) MgO 26.9 NaF 51.5

  37. Qualitative Guidelines for Entropy Changes Entropies of gases higher than liquids higher than solids Entropies are higher for more complex structures than simpler structures Entropies of ionic solids are inversely related to the strength of ionic forces Entropy increases when making solutions of pure solids or pure liquids in a liquid solvent Entropy decrease when making solutions of gases in a liquid

  38. Entropy of a Solution of a Gas

  39. 7) Arrange following in the order of increasing entropy? • a) C(s) (diamond) • b) C(s) (graphite) • c) O2(g) • d) CO2(g) • e) CO(g) • f) Hg(l)

  40. Entropy Change Entropy (DS) normally increase (+) for the following changes: i) Solid ---> liquid (melting) + ii) Liquid ---> gas + iii) Solid ----> gas most + iv) Increase in temperature + v) Increasing in pressure(constant volume, and temperature) + vi) Increase in volume +

  41. Qualitative prediction of DS of Chemical Reactions • Look for (l) or (s) --> (g) • If all are gases: calculate Dn Dn = Sn (gaseous prod.) - S n(gaseous reac.) N2 (g) + 3 H2 (g) --------> 2 NH3 (g) Dn = 2 - 4 = -2 If Dn is -DS is negative (decrease in S) If Dn is +DS is positive (increase in S)

  42. Predict DS! 2 C2H6(g) + 7 O2(g)--> 4 CO2(g) + 6H2O(g) 2 CO(g) + O2(g)-->2 CO2(g) HCl(g) + NH3(g)-->NH4Cl(s) H2(g) + Br2(l) --> 2 HBr(g)

  43. 8) Taking following physical and chemical changes qualitatively predict the sign of DS. • a) 2H2O (g) ------> 2 H2O (l) • b) 2H2O (g) ------> 2 H2 (g) + O2 (g) • c) N2(g) + 3 H2 (g) ------> 2 NH3 (g)

  44. Entropy Changes for Phase Changes For a phase change, DSSYS = qSYS/T (q = heat transferred) Boiling Water H2O (liq)  H2O(g) DH = q = +40,700 J/mol

  45. 9) How is entropy related to the heat and temperature?

  46. Phase Transitions Heat of Fusion energy associated with phase transition solid-to-liquid or liquid-to-solid DGfusion = 0 = DHfusion - T DSfusion 0 = DHfusion - T DSfusion DHfusion = T DSfusion Heat of Vaporization energy associated with phase transition gas-to-liquid or liquid-to-gas DHvaporization = T DSvaporization

  47. 10) The normal boiling point of benzene is 80.1°C and heat of evaporation (∆H°vap)is 30.7 kJ/mol. Calculate the ∆Ssurr (in J/K mol) for the evaporation of benzene.

  48. 2nd Law of Thermodynamics 2 H2(g) + O2(g)  2 H2O(liq) DSosys = -326.9 J/K Entropy Changes in the Surroundings Can calc. that DHorxn = DHosystem = -571.7 kJ = +1917 J/K

  49. 2nd Law of Thermodynamics 2 H2(g) + O2(g)  2 H2O(liq) DSosys = -326.9 J/K DSosurr = +1917 J/K DSouni = +1590. J/K The entropy of the universe is increasing, so the reaction is product-favored.

  50. H -D sys S = + S D D univ sys T Gibbs Free Energy, G DSuniv = DSsurr + DSsys Multiply through by (-T) -TDSuniv = DHsys - TDSsys -TDSuniv = DGsystem Under standard conditions — DGo = DHo - TDSo

More Related