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ChE 452 Lecture 02

ChE 452 Lecture 02. Basic Concepts. Stoichiometry. Invented by Lavoisier Molecules react in fixed proportions  1 A +  2 B   3 C+  4 D Stoichiometric Coefficent,  n Number of molecules produced when reaction goes once. Example. 2CO + O 2  2CO 2 CO + ½ O 2  CO 2. Answer.

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ChE 452 Lecture 02

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  1. ChE 452 Lecture 02 Basic Concepts

  2. Stoichiometry • Invented by Lavoisier • Molecules react in fixed proportions 1 A + 2B  3C+ 4D • Stoichiometric Coefficent, n • Number of molecules produced when reaction goes once

  3. Example 2CO + O2 2CO2 CO + ½ O2  CO2 Answer Answer

  4. Reaction Rate Priestley Total Moles/hr Van’t Hoff Moles/hr/volume

  5. Priestley’s Experiment Changing the amount of powder did not change the rate of weight loss

  6. Van’t Hoff’s Experiments 2N2O5  4NO2 + O2 Special poisoned walls Reaction rate scaled as the volume of the reactor

  7. Reaction Rate production rate (Moles/hr) reactor volume (liter) • rA called the Rate of production of A • rA has dimensions moles/lit-hr • rA is positive for a product, negative for a reactant

  8. Heterogeneous Reactions Scale As Surface Area production rate (Moles/hr) surface area (cm2) • RA is also called the Rate of production of A • RA has dimensions moles/cm2 –hr • RA is positive for a product, negative for a reactant

  9. The Rate Of A Reaction

  10. Heterogeneous vs. Homogeneous Reactions

  11. Variations In RateWith Conditions Rates vary with: • Concentrations of all species (reactants, products, inerts) (factors of 2-100) • Temperature (factors of 100 or more) • The presence of solvents (factors of 1012 or more) • The presence of catalysts (factors of 1012 or more)

  12. General Effect Of Concentration Rate for homogeneous reactions goes up with concentration Figure 2.3 The rate of CH3NC isomerization to CH3CN as a function of the CH3CN pressure

  13. Rate Of Heterogeneous Reactions Often Shows A Maximum Figure 2.15 The influence of the CO pressure on the rate of CO oxidation on Rh(111). Data of Schwartz, Schmidt, and Fisher.

  14. Rate Equations • Rate as a function of the concentration of all of the species in the reactor

  15. Typical Rate Laws For Simple A  C Reactions: rA = -k (CA) First order rA = -k (CA)2 Second order rA = -k (CA)3 Third order rA = -k (CA)n nth order n is the order k is the rate constant

  16. Rate Equations ForA + B  C rA = -k (CA)n(CB)m nth order in A, mth order in B overall (m+n)th order rA = -k(CA)(CB)2 First order in A, second order in B, third order overall. (2.13)

  17. Notation k1, k2 rate constants K1, K2 equilibrium constants CA=[A] concentration of species A

  18. Discussion Problem 1 What is the order of reaction? Answer

  19. Discussion Problem 2 What is the order of the reaction? Answer

  20. More Complex Rate Equations • Very few real reactions have simple reaction orders over a wide range of conditions: (2.18)

  21. More Complex Rate Equations CH3NC  CH3CN No industrially important reaction is first order over a wide range of conditions. Figure 2.3 The rate of CH3NC isomerization to CH3CN as a function of the CH3CN pressure

  22. Rate Equation Not Simply Related To Stoichiometry (2.21)

  23. ChBE424 Assumed First Or Second Order Reactions Why does this work? • Usually reactors run under a limited range of conditions – fitting rate to a line or parabola often good enough • Rate changes with temperature or catalysts much larger than changes with concentration.

  24. Not All Reactions Have Rate Equations Figure 2.22 2CO+O22 CO2 On ruthenium Figure 2.23Belousov-Zhabotinskii Reaction

  25. Summary For Today Table 2.1 Summary of key definitions

  26. Summary For Today Table 2.4 The key definitions from Section 2.3.

  27. Summary For Today • Real reactions rarely follow simple rate laws over wide ranges of conditions. • Simple rate laws OK if the range of conditions small • Some reactions do not have a rate law.

  28. Class Question • What did you learn new today?

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