1. IC-1/18 Lecture-2 02-09-2004 Kinetics

2. IC-2/18 Lecture-2 02-09-2004 What is Kinetics ? Analysis of reaction mechanisms on the molecular scale Derivation of rate expressions Design and analysis of experiments to test rate equations and derive kinetic parameters Theoretical prediction of rate constants How can we improve it?

3. IC-3/18 Lecture-2 02-09-2004 Elementary steps Discuss examples of elementary steps! A reaction is elementary if it cannot be split up in further steps. An elementary reaction proceeds exactly as expressed by the reaction equation. A reaction mechanism is a sequence of elementary steps.

4. IC-4/18 Lecture-2 02-09-2004 The Rate Equation

5. IC-5/18 Lecture-2 02-09-2004 Free Energy and Entropy Equilibrium:  G = 0 Free energy: minimum Entropy: maximum

6. IC-6/18 Lecture-2 02-09-2004 Reaction:  G :  0 free energy: > minimum entropy: < maximum Free Energy and Entropy Kinetics Thermodynamics

7. IC-7/18 Lecture-2 02-09-2004 Chemical Equilibrium The Chemical Potential

8. IC-8/18 Lecture-2 02-09-2004 Thermodynamics Table 2.2 Thermodynamic Data for Important Catalytic Reactions Reaction  H o (kJ/mol)  G o 298(K) NH 3  ½ N 2 + 3/2 H 2 + 45.9 + 16.4 ½N 2 + 3/2 H 2  NH 3 - 45.9 - 16.4 N 2 + 3 H 2  2 NH 3 * - 91.9* - 32.8 2 NO  N 2 + O 2 *- 182.6 -175.2 CH 4 + H 2 O  CO + 3 H 2 +205.9 +142.0 CH 4 + ½ O 2  CO + 2 H 2 - 35.9 - 86.8 CH 4 + 2 O 2  CO 2 + 2 H 2 O- 802.6 - 801.0 CH 4 + ½ O 2  CH 3 OH- 275.6 - 111.77 CO 2 + 3H 2  CH 3 OH + H 2 O - 49.3 + 3.5 CO + 2H 2  CH 3 OH- 90.5 - 25.1 CO + H 2 O  CO 2 + H 2 - 41.2 - 28.6 * this per two mole of NH 3 or NO. Data taken from the JANAF Thermodynamic Tables.

9. IC-9/18 Lecture-2 02-09-2004 EX: Ammonia Synthesis N 2 + 3 H 2 = 2 NH 3 G° = -32.8 kJ/mol H° = -91.9 kJ/mol

10. IC-10/18 Lecture-2 02-09-2004 inlet outlet A B D C The ammonia reactor N 2 +3 H 2 = 2 NH 3  G° = -32.8 kJ/mol  H° = -91.9 kJ/mol

11. IC-11/18 Lecture-2 02-09-2004 Power Rate Laws Parametrization of the rate: Reaction Order Example:

12. IC-12/18 Lecture-2 02-09-2004 Equilibrium Constant and Free Energy Van 't Hoff equation: `  H° = standard reaction enthalpy  S° = Standard reaction entropy  G° = standard Gibbs energy of reaction=  H° - T  S° K = e -  G° / RT = e - (  H°-T  S°) / RT

13. IC-13/18 Lecture-2 02-09-2004 Temperature dependence of the rate constant - E act / RT E act reaction parameter E products reactants E act = activation energy v = preexponential factor Arrhenius equation: k = v e

14. IC-14/18 Lecture-2 02-09-2004 Potential Energy Representation: dissociation - association

15. IC-15/18 Lecture-2 02-09-2004 Temperature Dependence of the Rate r = = k [A] [B] d [AB] d t 1000 / T ln k ln k = ln v E act / RT k = v e - E act /RT Arrhenius:

16. IC-16/18 Lecture-2 02-09-2004 Apparent Activation Energy Example:

17. IC-17/18 Lecture-2 02-09-2004 Heterogeneous Catalysis Adsorption Reaction Desorption

18. IC-18/18 Lecture-2 02-09-2004 Basic surface interactions Reactions take place on the metal surface Desorption Reaction Molecular adsorption Dissociativ adsorption Diffusion CO CO 2 O