1. How fast chemical reactions proceed How chemical reactions occur

4. Rate: Change of any physical property with time.

5. A B rate = -  [A] tt rate = [B][B] tt time

6. The Decomposition of Nitrogen Dioxide

8. aA + bB cC + dD rate = -  [A] tt 1 a = -  [B] tt 1 b =  [C] tt 1 c =  [D] tt 1 d

9. Rate Law Equation that expresses how the rate depends on the concentrations of reactants (sometimes also the products). 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 reaction with respect to NO 2 The value of the exponent n must be determined by experiment; cannot be written from the stoichiometric coefficient.

10. Reaction Order For the general equation: aA + bB  pP The rate equation is: Rate = k [A] m [B] n m and n are experimentally determined and are usually integers (0, 1, 2, 3, …). They may be fractions. The reaction is: m th order with respect to A n th order with respect to B. Overall Reaction Order = m + n

11. 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.

12. The rate determined in the linear range just after the reaction starts. The initial concentration don’t change significantly.

13. m=1 n=1 k=2.74×10 -4 L/mol.s

14. n=1m=1p=2k=8 L 3 /mol 3.s

15. y = a + b x

17. =-k

18. time needed for the reactant to reach half its initial concentration For first-order reaction: half-life doesn’t depend on initial concentration!

20. 75% complete means: 75% of the reactant A has been converted into product 25% of A still unreacted [A]=0.25*[A] o Two half-lives!

21. Half-Life Graph 21

22. y = a + b x

23. 2 nd order k=slope

24. Half-life of 2 nd order reaction 2 nd half-life is 3260 s!!

25. y = a + b x Half-life

26. Pseudo-first order reaction Initial conc M Complete reaction [BrO 3 - ] o 1.0×10 -3 0 [Br - ] o 1.00.995constant [H + ] o 1.00.994constant k’

27. Series of steps by which chemical reactions occur. experiment intermediate

28. Elementary steps Elementary: Can not be simplified into a series of steps Represents a simple collision process between reacting molecules. Molecularity: number of molecules colliding in an elementary step. Rate law can be written from molecularity of elementary reactions.

29. Reaction mechanism: A series of elementary steps that satisfy the following requirements: 1. Sum of elementary steps = overall reaction 2. Agrees with experimentally determined rate law.

31. Rate is that of slowest step.

32. Collision Theory Reacting molecules must collide with each other for the reaction to take place. The collision must have the right orientation NO 2 Cl + Cl  NO 2 + Cl 2

33. The collision must have enough energy to overcome an energy barrier called activation energy. This activation energy is needed to loosen the bond to be broken.

34. NO 2 Cl + Cl  NO 2 + Cl 2 As NO 2 Cl and Cl come together –Start to form Cl····Cl bond –Start to break N····Cl bond Requires E, as must bring 2 things together In TS –N····Cl bond ½ broken –Cl····Cl bond ½ formed After TS –Cl—Cl bond forms –N····Cl breaks Releases E as products more stable

36. Effect of temperature EaEa

37. z : number of collisions p : steric factor Arrhenius Equation y = a + b x

38. E a determination Measurement of k at different temperatures

42. Catalyst Substance that changes rate of chemical reaction without itself being used up Speeds up reaction, but not consumed by reaction Appears in mechanism, but not in overall reaction Does not undergo permanent chemical change Regenerated at end of reaction mechanism May appear in rate law May be heterogeneous or homogeneous

44. Heterogeneous Catalyst H 2 & N 2 approach Fe catalyst H 2 & N 2 bind to Fe & bonds break N—H bonds forming NH 3 formation complete NH 3 dissociates

45. Enzymes Acid rain