Unit 5-1 Reaction rates

1. Unit 5-1 Reaction rates

2. Kinetics • Studies the rate at which a chemical process occurs. • Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

3. Factors That Affect Reaction Rates • Physical State of the Reactants • In order to react, molecules must come in contact with each other. • The more homogeneous the mixture of reactants, the faster the molecules can react.

4. Factors That Affect Reaction Rates • Concentration of Reactants • As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

5. Factors That Affect Reaction Rates • Temperature • At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. • Presence of a Catalyst • Catalysts speed up reactions by changing the mechanism of the reaction. • Catalysts are not consumed during the course of the reaction.

6. Reaction Rates Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.

7. Reaction Rate The change in concentration of a reactant or product per unit of time (M/s if in aq)

8. Average rate = ― [C4H9Cl] t Reaction Rates based on reactants C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) The average rate of the reaction over each interval is the change in concentration divided by the change in time:

9. Reaction Rates based on reactants C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • Note that the average rate decreases as the reaction proceeds. • This is because as the reaction goes forward, there are fewer collisions between reactant molecules.

10. Reaction Rates based on reactants C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • A plot of concentration vs. time for this reaction yields a curve like this. • The slope of a line tangent to the curve at any point is the instantaneous rate at that time.

11. Reaction Rates based on reactants C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • All reactions slow down over time. • Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning.

12. ―[C4H9Cl] t Rate = = [C4H9OH] t Reaction Rates based on products C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • In this reaction, the ratio of C4H9Cl to C4H9OH is 1:1. • Thus, the rate of disappearance of C4H9Cl is the same as the rate of appearance of C4H9OH.

13. 1 2 [HI] t Rate = − = [I2] t Reaction Rates and Stoichiometry • What if the ratio is not 1:1? 2 HI(g) H2(g) + I2(g) • Therefore,

14. aA + bB cC + dD = = Rate = − = ― 1 a 1 b 1 c 1 d [C] t [D] t [A] t [B] t Reaction Rates and Stoichiometry • To generalize, then, for the reaction

15. 2NO2(G)  2NO(G) + O2(G) Reaction Rates: 1. Can measure disappearance of reactants 2. Can measure appearance of products 3. Are proportional stoichiometrically

16. 2NO2(G)  2NO(G) + O2(G) Reaction Rates: 4. Are equal to the slope tangent to that point 5. Change as the reaction proceeds, if the rate is dependent upon concentration [NO2] t

17. Page 549 14.3, 14.9, 14.11 Homework

18. Rate Laws

19. Concentration and Rate One can gain information about the rate of a reaction by seeing how the rate changes with changes in initial concentrations of the reactants.

20. N2(g) + 2 H2O(l) NH4+(aq) + NO2−(aq) Concentration and Rate • Comparing Experiments 1 and 2, when [NH4+] doubles, the initial rate doubles. • Likewise, comparing Experiments 5 and 6, when [NO2−] doubles, the initial rate doubles.

21. Concentration and Rate • This means Rate  [NH4+] Rate  [NO2−] Rate  [NH+] [NO2−] or Rate = k [NH4+] [NO2−] • This equation is called the (differential)rate law, and k is the rate constant.

22. Rate Laws Differential rate lawsexpress (reveal) the relationship between the concentration of reactants and the rate of the reaction. The differential rate law is usually just called “the rate law.” Integrated rate lawsexpress (reveal) the relationship between concentration of reactants and time

23. Differential Rate Laws • A rate law shows the relationship between the reaction rate and the concentrations of reactants. • The exponents tell the order of the reaction with respect to each reactant. • This reaction (Rate = k [NH4+] [NO2−]) is First-order in [NH4+] First-order in [NO2−]

24. Differential Rate Laws • The overall reaction order can be found by adding the exponents on the reactants in the rate law. • This reaction is second-order overall. • The rxn order starts at zero order • O order: rate = k, the rate doesn’t affected by [A].

25. Writing a (differential) Rate Law Problem- Write the rate law, determine the value of the rate constant, k, and the overall order for the following reaction: 2 NO(g) + Cl2(g)  2 NOCl(g)

26. Writing a Rate Law Part 1– Determine the values for the exponents in the rate law: R = k[NO]x[Cl2]y In experiment 1 and 2, [Cl2] is constant while [NO] doubles. The rate quadruples, so the reaction is second order with respect to [NO]  R = k[NO]2[Cl2]y

27. Writing a Rate Law Part 1– Determine the values for the exponents in the rate law: R = k[NO]2[Cl2]y In experiment 2 and 4, [NO] is constant while [Cl2] doubles. The rate doubles, so the reaction is first order with respect to [Cl2]  R = k[NO]2[Cl2]

28. Writing a Rate Law Part 2– Determine the value for k, the rate constant, by using any set of experimental data: R = k[NO]2[Cl2]

29. Writing a Rate Law Part 3– Determine the overall order for the reaction. R = k[NO]2[Cl2] + 2 = 3 1  The reaction is 3rd order Overall order is the sum of the exponents, or orders, of the reactants

30. Page 550 14.15, 14.17, 14.19, 14.21, Homework