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Thermodynamics

Thermodynamics. Modern Methods in Heterogeneous Catalysis F.C. Jentoft, November 1, 2002. Outline. Part I: Reaction + Catalyst Thermodynamics of the target reaction Thermodynamics of catalyst: bulk (see classes on solids and defects) and surface

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Thermodynamics

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  1. Thermodynamics Modern Methods in Heterogeneous Catalysis F.C. Jentoft, November 1, 2002

  2. Outline Part I: Reaction + Catalyst • Thermodynamics of the target reaction • Thermodynamics of catalyst: bulk (see classes on solids and defects) and surface • Thermodynamics of interaction between reactant and catalyst (see class on adsorption) Part II: Practical Matters • Vapor pressure

  3. What Thermodynamics Will Deliver… Gives “big picture”, essence, useful for estimates

  4. Target Reaction - Motivation • Why look at TD? …can’t change it anyway by catalysis E E without catalyst with catalyst Must look at TD because we can’t change it! EA EA Reactants Reactants Products Products Reaction coordinate Reaction coordinate

  5. Target Reaction – Quantities to Look at • Enthalpy of reaction ΔrHexothermic / endothermicΔrH of side reactions • Free Enthalpy (Gibbs Energy) ΔrGexergonic / endergonic • Equilibrium Constant K: Equilibrium Limitations • Change of Temperature and Pressure (variables)

  6. Enthalpy of Reaction • Determines reactor setup (see classes on catalyst testing and reaction engineering)catalyst formulation / dilution“hot spots” / heating powerisothermal operation in the lab • Enthalpy of side reactionsparallel / secondary reactions

  7. Enthalpy of Reaction, ΔrH • Reaction enthalpy needs a reaction equation!!! • Calculate from enthalpies of formation of products and reactants ΔrH°: standard enthalpy of reaction ΔfH°: standard enthalpies of formation vi: stoichiometric factors, positive for products, negative for reactants

  8. Things to Watch in Calculations….. • Stoichiometric factors • Standard conditions • State of the matter (solid, liquid, gaseous) • Which data are available (sometimes only enthalpy of combustion, ΔcH° )

  9. Standard Conditions (IUPAC) • International Union of Pure and Applied Chemistry (IUPAC) Größen, Einheiten und Symbole in der Physikalischen Chemie VCH , Weinheim 1996 FHI library 50 E 49 (English version: 50 E 48) • Standard state indicated by superscript ,° www.iupac.org

  10. Standard Conditions (IUPAC) • „Standard state pressure“(IUPAC 1982) p° = 105 Pa „Standard atmosphere“ (before 1982)p° = 101 325 Pa = 1 atm • „Standard concentration“ c° = 1 mol dm-3 • „Standard molality“ m° = 1 mol kg-1 • „Standard temperature“ ??

  11. Standard Conditions (Textbooks) • AtkinsSTP „Standard temperature and pressure““p = 101 325 Pa = 1 atm, T° = 273,15 K SATP „Standard ambient temperature and pressure“p° = 105 Pa = 1 bar, T° = 298,15 K • Wedler „Standarddruck“p = 1.013 bar = 1 atm = 101.325 kPa„Standardtemperatur“T° = 298,15 K

  12. Standard Conditions (Other) • Catalysis Literature NTP „Normal temperature and pressure““20°C and 760 torr70 degrees F and 14.7 psia (1 atmosphere) ALWAYS CHECK / SPECIFY THE CONDITIONS !!

  13. Sources for Thermodynamic Data • CRC Handbook of Thermophysical and Thermochemical DataEds. David R. Lide, Henry V. Kehiaian CRC Press Boca Raton New York 1994 FHI library 50 E 55 • D'Ans Lax Taschenbuch für Chemiker und Physiker Ed. C. Synowietz Springer Verlag 1983 FHI library 50 E 54

  14. Some Examples: Combustion • Combustion of hydrogen (Knallgasreaktion) ΔcH° = -286 kJ mol-1 • Combustion of carbonΔcH° = -394 kJ mol-1 Reactions with CO2, H2O or other very stable molecules as products are usually strongly exothermic, however….

  15. Steam Reforming of Methanol ΔcH° = 93 kJ mol-1

  16. ΔfH° = 49.0 kJ mol-1 State of the Matter • Formation of benzene at 298.15 K ΔfH° = 82.93 kJ mol-1 • Enthalpy of evaporation of benzene?ΔvapH° = 30.8 kJ mol-1 at 80°C

  17. Partial Oxidation of Propene • Oxidation of propene to acroleinΔrH° = ??? kJ mol-1

  18. Examples for Sources

  19. Examples for Sources

  20. Partial Oxidation • Only enthalpy of combustion, ΔcH°, of acrolein is given ΔcH° = -1633 kJ mol-1 Enthalpies of combustion are easily determined quantities (e.g. from quantitative combustion in a bomb calorimeter)

  21. ΔcH° = -1754 kJ mol-1 ΔcH° = -1633 kJ mol-1 ΔfH° = -121 kJ mol-1 Use Hess’s Law Enthalpy is a State Function

  22. E EA EA Reactants Partial Oxidation Product Total Oxidation Products • Oxidation of acrolein to CO2 ΔcH° = -1633 kJ mol-1 Reaction coordinate Partial vs. Total Oxidation • Oxidation of propene to acroleinΔrH° = -427 kJ mol-1

  23. Oxidative dehydrogenation of isobutane to isobuteneΔrH° = -124 kJ mol-1 Dehydrogenation vs. Oxidative Dehydrogenation • Dehydrogenation of isobutane to isobuteneΔrH° = 117 kJ mol-1

  24. Oxidative Dehydrogenation:Thermodynamic Traps • Combustion of isobuteneΔcH° = - 2525 kJ mol-1 Nevertheless, the oxidative dehydrogenation of isobutene is in commercial operation (CrO3/Al2O3 or supported Pt catalyst)

  25. Dehydrogenation • Dehydrogenation of ethylbenzene to styreneΔrH° = 117 kJ mol-1

  26. Enthalpy Products, T2 Products, T1 ΔrH2 ΔrH1 Reactants, T2 Reactants, T1 Reaction coordinate Change of ΔrH with Temperature • Most of the time, we are not interested in room temperature

  27. How to Calculate ΔrH as Function of T • Each enthalpy in the reaction equation changes according to Kirchhoff’s law • And, if Cp = constant over the temperature range of interest

  28. Heat Capacity as a Function of T, Condensed Phases

  29. Heat Capacity as a Function of T, Gases

  30. How to Calculate ΔrH as Function of T • Cp as a function of temperature is usually a polynomial expression such as • If there is a phase transition within the temperature range, it must be accounted for

  31. Isomerization • Isomerization of butane ΔrH° = - 7 kJ mol-1 ΔrS° = -15 J mol-1 ΔrG°= - 2.3 kJ mol-1 • Consistency check....

  32. Free Enthalpy ΔrG, and Equilibrium Constant K • Relation between ΔrG° and K in equilibrium, ΔrG=0 • Composition dependence of ΔrG • Thermodynamic equilibrium constant (dimensionless)

  33. Different Equilibrium Constants K • Kp • correlation between Kth and Kp [Pai] For low pressures (a few bars and less), the fugacity coefficients are about 1 All pressures, including po should be in the same units.

  34. at 298 K 28 % 72 % Isomerization Equilibrium • Isomerization of butane ΔrG°= - 2.3 kJ mol-1 • With and

  35. Indefinite integration Definite integration Equilibrium Constant Temperature Dependence van’t Hoff’s Equation

  36. Equilibrium Temperature Dependence Start your research by calculating the thermodynamics of your reaction!

  37. Part II: Practical Matters • Vapor pressure and saturators Gas in Gas out Saturator, 100 ml Methanol 79.17 g, is 2.47 mol

  38. Methanol Thermodynamic Data

  39. Heat Consumed by Evaporation • Assumption: saturator is adiabatic, evaporate 20 ml of methanol, all energy for evaporation taken from remaining 80 ml methanol • 20 ml is about 0.5 mol, need about 17.7 kJ for evaporation • 80 ml is about 2 mol, Cp of liquid MeOH is 81.6 J mol-1 K-1 • The temperature of the methanol would theoretically drop by 108 K

  40. For sublimation and evaporation assumes ideal behavior of the gas phase August’s vapor pressure formula assumes enthalpy is constant within given temperature range The Clausius-Clapeyron Equation General differential form of the Clausius-Clapeyron Equation

  41. Vapor Pressure and Temperature • At 64.4°C, the vapor pressure of methanol is 755 torr and the enthalpy of evaporation is 35.4 kJ mol-1 • T1 = 337.6 K, p = 100.66 kPa • The carrier gas will dissolve in the liquid and the vapor pressure will be lowered

  42. Methanol Vapor Pressure Small temperature changes can cause significant changes in vapor pressure

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