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# Chemical Kinetics

Chemical Kinetics. Rates of chemical reaction － definition of reaction rate － integrated and differential rate law － determination of rate law Mechanism of chemical reaction － activated complex theory － model for chemical kinetics － Arrhenius equation. Rate Law. Rate=k[NO 2 ] n

## Chemical Kinetics

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### Presentation Transcript

1. Chemical Kinetics • Rates of chemical reaction －definition of reaction rate －integrated and differential rate law － determination of rate law • Mechanism of chemical reaction － activated complex theory －model for chemical kinetics －Arrhenius equation

2. Rate Law • Rate=k[NO2]n • The concentration of the products do not appear in the rate law. • The value of the exponent n must be determined by experiment; it cannot be written from the balanced equation.

3. 2N2O5→4NO2+O2

4. Types of Rate Laws • Integrated Rate Law: how the concentration of species in the reaction depend on time • Differential Rate Law: how the rate of a reaction depends on concentrations • Determine the differential rate law for a given reaction, the form of integrated rate law can be automatically known, and vice versa

5. Initial-Rate method • To determine the instantaneous rate before the initial concentration of reactants have changed significantly. • Several experiments are carried out using different initial concentrations. • The initial rate is determined for each run.

6. Integrated Rate Law－first order

7. Plot of N2O5 vs. time

8. Half-Life of a First Order Reaction • The time required for a reaction to reach half of its original concentration is called half-life of a reaction and id designated by t1/2.

9. Plot of N2O5 vs. time

10. Integrated Rate Law－second order

11. Plot of C4H6

12. Integrated Rate Law－zero order

13. Pseudo-Order Reaction Law

14. Arrhenius Postulations • Collisions and Rate －the rate of reaction is much smaller than calculated collision frequency. • A threshold energy (activation energy) －This kinetic energy is changed into potential energy as the molecules are distorted during a collision, breaking bonds and rearranging the atoms into product molecules.

15. Collisions Frequency and Molecular orientations • Experiments show that the observed reaction rate is considerably smaller than the rate of collisions with enough energy to surmount the barrier. • The collision must involve enough energy to produce the reaction. • The relative orientation of the reactants must allow formation of any new bonds necessary to products.

16. BrNO collision

17. Potential energy graph for 2BrNO→2NO+Br2

18. Temperature Dependence of Rate Constants • The order of each reactant depends on the detailed reaction mechanism. • Chemical reaction speed up when the temperature is increased. －molecules must collide to react －an increase in temperature increases the frequency of intermolecular collisions.

19. T1/T2 graph

20. T(K) and k

21. Arrhenius Equation Ea: activation energy A: pre-exponential factor

22. Plot ln(k) vs. 1/T

23. Reaction Mechanism • Most chemical reactions occur by a series of steps called the reaction mechanism. • The sum of the elementary steps must give the overall balanced equation for the reaction. • The mechanism must agree with the experimentally determined rate law.

24. Step (1): rate-determining-step

25. k1 Treatments1. Rate-Determining Step Approximation 2. Steady-State Approximation k-1 k2 2 NO N2O2 fast N2O2+H2 N2O+H2O slow Overall: 2NO+H2 N2O+H2O R=[NO]2[H2]

26. Rate-Determining Step Approximation

28. Overall reaction H++HNO2+C6H5NH2→C6H5N2++2H2O Proposed mechanism H++HNO2 H2NO2+ rapid equilib. H2NO2++Br-→ONBr+H2O slow ONBr+C6H5NH2 →C6H5N2++H2O+ Br-fast k1 k-1 k2 k3

29. Activated Complex Theory • The arrangement of atoms found at the top of potential energy hill or barrier is called the activated complex or transition state. • △E has no effect on the rate of reaction. • The rate depends on the size of the activation energy Ea

30. Catalysis • A substance can speed up a reaction without being consumed itself. • The catalyst is to provide a new pathway for the reaction and to decrease activation energy.

31. Effect of a catalyst

32. Heterogeneous Catalysis • Adsorption and activation of the reactants • Migration of the adsorbed reactants on the surface • Reaction among the adsorbed substances • Escape, or desorption, of the products.

33. Hydrogenation of ethylene

34. Exhaust gases

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