1. Summary comments on mechanism For a reaction mechanism to be viable, two main conditions apply. 1. The sum of the elementary steps must lead to the overall balanced equation of the observed chemical reaction. 361

2. Summary comments on mechanism For a reaction mechanism to be viable, two main conditions apply. 1. The sum of the elementary steps must lead to the overall balanced equation of the observed chemical reaction. 2. The slow step in the reaction mechanism should yield a rate law that corresponds with the rate law for the overall chemical reaction. 362

3. Summary comments on mechanism For a reaction mechanism to be viable, two main conditions apply. 1. The sum of the elementary steps must lead to the overall balanced equation of the observed chemical reaction. 2. The slow step in the reaction mechanism should yield a rate law that corresponds with the rate law for the overall chemical reaction. For complicated situations, some manipulation may be necessary to prove this correspondence. 363

4. A third condition also applies: 3. Each step in the mechanism should be chemically reasonable. 364

5. A third condition also applies: 3. Each step in the mechanism should be chemically reasonable. This can be rather difficult to apply for a person with limited experience. 365

6. A third condition also applies: 3. Each step in the mechanism should be chemically reasonable. This can be rather difficult to apply for a person with limited experience. The principal reason is that reaction mechanisms will often have highly reactive intermediate species present. And these would be difficult to conjecture for a person with limited background. 366

7. An elementary step for a gas phase reaction that required three bodies to simultaneously strike each other would usually not be a good choice for an elementary step, particularly if the reaction took place under low concentration conditions. 367

8. Catalysis 368

9. Catalysis Catalyst: A substance that increases the rate of a reaction and can be recovered at the end of the reaction. 369

10. Catalysis Catalyst: A substance that increases the rate of a reaction and can be recovered at the end of the reaction. Note that this definition does not exclude the possibility that the catalyst is directly involved in the chemistry. 370

11. Catalysis Catalyst: A substance that increases the rate of a reaction and can be recovered at the end of the reaction. Note that this definition does not exclude the possibility that the catalyst is directly involved in the chemistry. If the catalyst is transformed in one step in the sequence, it must be regenerated in a subsequent step in the sequence. 371

12. One way to speed up a reaction is to raise the temperature. This method can produce undesirable side effects. For example, at elevated temperatures the products formed may undergo other reactions. 372

13. One way to speed up a reaction is to raise the temperature. This method can produce undesirable side effects. For example, at elevated temperatures the products formed may undergo other reactions. A catalyst accelerates a reaction without any need to change the temperature. 373

14. One way to speed up a reaction is to raise the temperature. This method can produce undesirable side effects. For example, at elevated temperatures the products formed may undergo other reactions. A catalyst accelerates a reaction without any need to change the temperature. Regardless of their nature, all catalysts act in the same general manner. 374

15. Consider the following two mechanisms: 1. No catalyst: A + B product (slow) (suppose it is a one step mechanism) 375

16. Consider the following two mechanisms: 1. No catalyst: A + B product (slow) (suppose it is a one step mechanism) 2.Catalyst present: A + catalyst C (faster) B + C product + catalyst (faster) (suppose it is a two step mechanism) 376

17. Impact of a catalyst on the energy profile 377

18. Impact of a catalyst on the energy profile 378

19. 379

20. A catalyst lowers the overall activation energy required for a reaction by providing a completely different pathway for its progress. 380

21. A catalyst lowers the overall activation energy required for a reaction by providing a completely different pathway for its progress. Consider A + B C + D (no catalyst) 381

22. A catalyst lowers the overall activation energy required for a reaction by providing a completely different pathway for its progress. Consider A + B C + D (no catalyst) In the presence of a catalyst, A + B C + D 382

23. A catalyst lowers the overall activation energy required for a reaction by providing a completely different pathway for its progress. Consider A + B C + D (no catalyst) In the presence of a catalyst, A + B C + D By definition: 383

24. Note that the total energies of the reactants (A and B) and those of the products (C and D) are unchanged. 384

25. Note that the total energies of the reactants (A and B) and those of the products (C and D) are unchanged. The important change is the lowering of the activation energy from E a to E a,c. 385

26. Note that the total energies of the reactants (A and B) and those of the products (C and D) are unchanged. The important change is the lowering of the activation energy from E a to E a,c. The increase in the rate constant from k to k c can be understood by using the Arrhenius equation. 386

27. Note that the total energies of the reactants (A and B) and those of the products (C and D) are unchanged. The important change is the lowering of the activation energy from E a to E a,c. The increase in the rate constant from k to k c can be understood by using the Arrhenius equation. 387

28. Now take the ratio, so that: 388

29. Now take the ratio, so that: hence 389

30. Suppose the catalyst lowers the activation energy by 20. kJmol -1, that is, E a,c = E a - 20. (units kJ mol -1 ). In this case, assuming the reaction takes place at 25 o C, then 390

31. Suppose the catalyst lowers the activation energy by 20. kJmol -1, that is, E a,c = E a - 20. (units kJ mol -1 ). In this case, assuming the reaction takes place at 25 o C, then 391

32. Suppose the catalyst lowers the activation energy by 20. kJmol -1, that is, E a,c = E a - 20. (units kJ mol -1 ). In this case, assuming the reaction takes place at 25 o C, then 392

33. Suppose the catalyst lowers the activation energy by 20. kJmol -1, that is, E a,c = E a - 20. (units kJ mol -1 ). In this case, assuming the reaction takes place at 25 o C, then that is. 393

34. Suppose the catalyst lowers the activation energy by 20. kJmol -1, that is, E a,c = E a - 20. (units kJ mol -1 ). In this case, assuming the reaction takes place at 25 o C, then that is. The large increase in the rate constant is due to the exponential connection between k and E a. 394

35. Heterogeneous Catalysis 395

36. Heterogeneous Catalysis Heterogeneous catalysis: The catalyst is in a different phase than the reactants and products. 396

37. Heterogeneous Catalysis Heterogeneous catalysis: The catalyst is in a different phase than the reactants and products. Usually, the catalyst is a solid, and the reactants and products are in the gas or liquid phase. 397

38. Example: The synthesis of ammonia. N 2(g) + 3 H 2(g) 2 NH 3(g) 398

39. Example: The synthesis of ammonia. N 2(g) + 3 H 2(g) 2 NH 3(g) The atmosphere provides a very cheap source for N 2. 399

40. Example: The synthesis of ammonia. N 2(g) + 3 H 2(g) 2 NH 3(g) The atmosphere provides a very cheap source for N 2. Dihydrogen can be produced by passing steam over heated coal: H 2 O (g) + C (s) CO (g) + H 2(g) 400

41. Example: The synthesis of ammonia. N 2(g) + 3 H 2(g) 2 NH 3(g) The atmosphere provides a very cheap source for N 2. Dihydrogen can be produced by passing steam over heated coal: H 2 O (g) + C (s) CO (g) + H 2(g) The formation of NH 3 is slow at room temperature. 401

42. Example: The synthesis of ammonia. N 2(g) + 3 H 2(g) 2 NH 3(g) The atmosphere provides a very cheap source for N 2. Dihydrogen can be produced by passing steam over heated coal: H 2 O (g) + C (s) CO (g) + H 2(g) The formation of NH 3 is slow at room temperature. Raising the temperature does accelerate the reaction – but this also promotes the decomposition of NH 3 into N 2 and H 2. 402

43. 403

44. For an industrial process, require – 404

45. For an industrial process, require – 1. An appreciable rate. 405

46. For an industrial process, require – 1. An appreciable rate. 2. High yield. 406

47. For an industrial process, require – 1. An appreciable rate. 2. High yield. The problem of preparing NH 3 using the above reaction with an appropriate catalyst was solved by Fritz Haber (1905). He discovered that iron plus a few per cent of oxides of potassium and aluminum catalyzes the reaction. 407

48. In heterogeneous catalysis, the surface of the catalyst is usually the site of the reaction. 408

49. In heterogeneous catalysis, the surface of the catalyst is usually the site of the reaction. There are two ways in which molecules may be attached to the surface of a solid. 409

50. In heterogeneous catalysis, the surface of the catalyst is usually the site of the reaction. There are two ways in which molecules may be attached to the surface of a solid. Physical adsorption: Relatively weak intermolecular forces are responsible for holding the molecules on the surface. 410

51. In heterogeneous catalysis, the surface of the catalyst is usually the site of the reaction. There are two ways in which molecules may be attached to the surface of a solid. Physical adsorption: Relatively weak intermolecular forces are responsible for holding the molecules on the surface. Physical adsorption usually plays no role or only a very minor role in heterogeneous catalysis. 411

52. Chemical Adsorption: Involves the formation of covalent bonds between the molecules and the solid surface. 412

53. Chemical Adsorption: Involves the formation of covalent bonds between the molecules and the solid surface. A very important consequence of chemical adsorption is that normal covalent bonds of the reactant molecules are weakened. 413

54. One way to form NH 3 from N 2 and H 2 is to directly break the N N and H H bonds: 414

55. One way to form NH 3 from N 2 and H 2 is to directly break the N N and H H bonds: N 2(g) N (g) + N (g) 415

56. One way to form NH 3 from N 2 and H 2 is to directly break the N N and H H bonds: N 2(g) N (g) + N (g) H 2(g) H (g) + H (g) 416

57. One way to form NH 3 from N 2 and H 2 is to directly break the N N and H H bonds: N 2(g) N (g) + N (g) direct bond breaking H 2(g) H (g) + H (g) 417

58. One way to form NH 3 from N 2 and H 2 is to directly break the N N and H H bonds: N 2(g) N (g) + N (g) direct bond breaking H 2(g) H (g) + H (g) In the gas phase, this requires a large amount of energy (in the form of heat). This makes the process expensive on the industrial scale. Also, the NH 3 formed at these high temperatures would not be stable. 418

59. With a catalyst: 419

60. With a catalyst: N 2(g) N 2(chemisorbed) 420