2. Chpt 12 - Chemical Kinetics Reaction Rates Rate Laws Reaction Mechanisms Collision Theory Catalysis HW: Chpt 12 - pg. 580-592, #s Due Fri Jan. 8

3. Reaction Rate ? Change in concentration of a reactant or product per unit time. [A] means concentration of A in mol/L; A is the reactant or product being considered.

4. Decomposition of NO 2 Graph Instantaneous Rate Value of the rate at a particular time. Can be obtained by computing the slope of a line tangent to the curve at that point. 2NO 2 --> 2NO + O 2

5. Calculate ave. rate from data

6. Rate Law Shows how the rate depends on the concentrations of reactants. For the decomposition of nitrogen dioxide: 2NO 2 (g) → 2NO(g) + O 2 (g) Rate = k[NO 2 ] n :  k = rate constant  n = order of the reactant

7. Rate Law (cont) Rate = k[NO 2 ] n The concentrations of the products do not appear in the rate law because the reaction rate is being studied under conditions where the reverse reaction does not contribute to the overall rate. The value of the exponent n must be determined by experiment; it cannot be written from the balanced equation.

8. Types of Rate Laws Differential Rate Law (rate law) – shows how the rate of a reaction depends on concentrations. Integrated Rate Law – shows how the concentrations of species in the reaction depend on time.

9. Rate Laws (cont) typically consider reactions when the reverse reaction is unimportant, so our rate laws involve only [reactants]. differential and integrated rate laws for a given reaction are related in a well–defined way, can use either rate law. Experimental convenience usually dictates which type of rate law is determined experimentally. Knowing the rate law for a reaction is important mainly because we can usually infer the individual steps involved in the reaction from the specific form of the rate law.

10. Determining the form of the Rate Law Determine experimentally the power to which each reactant concentration must be raised in the rate law.

11. Method of Initial Rates The value of the initial rate is determined for each experiment at the same value of t as close to t = 0 as possible. Several experiments are carried out using different initial concentrations of each of the reactants, and the initial rate is determined for each run. The results are then compared to see how the initial rate depends on the initial concentrations of each of the reactants.

13. Initial Rates examples Pg. 549 Table 12.4 discuss data Pg. 550 Table 12.5 discuss data

14. Overall Reaction Order The sum of the exponents in the reaction rate equation. Rate = k[A] n [B] m Overall reaction order = n + m k = rate constant [A] = concentration of reactant A [B] = concentration of reactant B

15. Integrated Rate Laws - zero first & second order Take actual data for each reactant and plot it… [A] vs t or ln[A] vs t or 1/[A] vs t straight line says O o, 1 o, 2 o for calculations at any time t t = time, [A] o = initial conc, [A] = conc at t, k = rate constant

16. Which order?

18. Exercise of Reaction Order Consider the reaction aA  Products. [A] 0 = 5.0 M and k = 1.0 x 10 –2 (assume the units are appropriate for each case). Calculate [A] after 30.0 seconds have passed, assuming the reaction is: a) Zero order b) First order c) Second order 4.7 M 3.7 M 2.0 M

19. Reaction Mechanism Most chemical reactions occur by a series of elementary steps. An intermediate is formed in one step and used up in a subsequent step and thus is never seen as a product in the overall balanced reaction.

20. Mechanism example A Molecular Representation of the Elementary Steps in the Reaction of NO 2 and CO NO 2 (g) + CO(g) → NO(g) + CO 2 (g)

21. Mechanisms Unimolecular – reaction involving one molecule; first order. Bimolecular – reaction involving the collision of two species; second order. Termolecular – reaction involving the collision of three species; third order. Elementary Steps (Molecularity)

23. Heterogeneous Catalysis