Chemical Kinetics

1. Chemical Kinetics BLB 11th Chapter 14

2. Chemical Reactions • Will the reaction occur? Ch. 5, 19 • How fast will the reaction occur? Ch. 14 • How far will the reaction proceed? Ch. 15

3. Expectations • Review: concentration units, graphing (line, slope) • Work with reaction rates. • Determine a rate law for a given reaction from experimental data. • Work with rate laws. • Understand the factors that affect the rate of a chemical reaction. • Calculate the activation energy.

4. Thermochemistry review… C(s, diamond) C(s, graphite) ΔH °rxn = Is the reaction favorable?

5. Kinetics … • studies the rates at which chemical reactions occur. • gives information about how the reaction occur, that is, the reaction mechanism • Determining the reaction mechanism is the overall goal of kinetic studies. (sect. 6)

6. 14.1 Factors that Affect Reaction Rates • Physical state of the reactants – states that promote contact have faster rates; homogeneous vs. heterogeneous • Concentration of the reactants: conc. ↑, rate ↑ (or pressure for gases) • Temperature: temp. ↑, rate ↑ due to higher molecular energy and speed (section 5) • Catalysts: rate ↑ by changing the mechanism and reaction energy (section 7) • Other physical things like stirring and grinding solid reactants.

7. 14.2 Reaction Rates • Rate – change in some variable per unit time • Reaction rate – change in concentration per unit time; M/s or mol/L·s • Rates are determined by monitoring concentration as a function of time. • Rates are positive quantities; for reactant A:

8. C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.

9. Reaction Rates cont. • Rates change over time: • reactant rates decrease • product rates increase • Instantaneous rate – rate at a specific time • Average rate – Δ[A] over a specific time interval • Initial rate – instantaneous rate at t = 0 • Note: Rates and rate laws are not based on stoichiometry!! They must be determined experimentally.

11. C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) (Fig. 14.4, p. 577)

12. Reaction Rates and Stoichiometry • The molar ratios between reactants and products correspond to the relative rates of the reaction. • Relative rates – relationship between rates of reactant disappearance and product appearance at a given time. 2 HI(g) → H2(g) + I2(g)

13. 2 NO2(g) → 2 NO(g) + O2(g) Rate = 2.4 x 10-5 M/s Rate = 8.6 x 10-5 M/s Rate = 4.3 x 10-5 M/s

14. Relative Rates PracticeAt a given time, the rate of C2H4 reaction is 0.23 M/s. What are the rates of the other reaction components?C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(g)0.23 M/s ? ? ?

15. Types of Rate Laws • Differential rate law or rate law (section 3) • Shows how the reaction rate changes with concentration • Must be specific in how defined • Integrated rate law (section 4) • Shows how concentration changes with time • Graphical determination of the order

16. 14.3 Concentration & Rate aA + bB → cC + dD • General form of rate law: [A], [B] – conc. in M or P rate = k[A]m[B]nk – rate constant; units vary m, n – reaction orders • Reaction orders and, thus, rate laws must be determined EXPERIMENTALLY!!! • Note: m ≠ a and n ≠ b • Overall order = sum of individual orders • Rate constant is independent of concentration.

17. Reaction OrdersFor the reaction: A →B, the rate law is: rate = k[A]m

18. Units of the rate constant, k

19. 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)rate = k[NO]2[H2] • What is the order with respect to NO? • What is the order with respect to H2? • What is the overall order? • If [NO] is doubled, what is the effect on the reaction rate? • If [H2] is halved, what is the effect on the reaction rate? • What are the units of k?

20. PtCl2(NH3)2 + H2O → PtCl(H2O)(NH3)2 + Cl¯rate = k[PtCl2(NH3)2] k = 0.0090 h-1 • Calculate the rate of reaction when the concentration of PtCl2(NH3)2 is 0.020M. • What is the rate of Cl¯ production under these conditions?

21. Practice problem 26 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)rate = k[NO]2[H2] k = 6.0 x 104M-2s-1 @1000K (b) Calculate rate when [NO] = 0.025 M and [H2] = 0.015 M.

22. Determining Rate Laws Initial Rates Method • Find two experiments in which all but one reactant’s concentration is constant. • Observe the relationship between concentration change and rate change to determine the order for that reactant. • Repeat for other reactant(s).

23. NH4+(aq) + NO2¯(aq) → N2(g) + 2 H2O(l) Determine the rate law and calculatek.

24. 2 NO(g) + Cl2(g) → 2 NOCl(g) Determine the rate law and calculatek.

25. BrO3¯(aq) + 5 Br¯(aq) + 6 H+(aq) → Br2(l) + 2 H2O(l) Determine the rate law and calculatek.

26. More Practice

27. 14.4 The Change of Concentration with TimeaA→ products • 1st order y = mx + b Plot: ln[A] vs. t slope = −k integrate

28. 14.4 The Change of Concentration with TimeaA → products • 2nd order y = mx + b Plot: 1/[A] vs. t slope = k integrate

29. 14.4 The Change of Concentration with TimeaA → products • zero order y = mx + b Plot: [A] vs. t slope = −k integrate

30. Integrated rate law analysis (BLB 42)CH3NC(g) → CH3CN(g) - isomerization

31. Integrated rate law analysis (BLB 46)C12H22O11(aq) + H2O(l) → 2 C6H12O6(aq)

32. Half-life, t1/2 • t1/2 – time required for the concentration of a reactant to decrease by half of its initial value

33. 1st order half-life • All half-lives same length of time • Independent of initial concentration Note: Radioactive decay follows 1st order kinetics.

34. 1st order reaction

35. 2nd order half-life • All half-lives different length of time • Dependent on initial concentration • Zero half-life • All half-lives different length of time • Dependent on initial concentration

37. BLB 37

38. BLB 39

40. To Nuclear Chemistry, Ch. 21

41. 14.5 Temperature and Rate • Generally, as temperature increases, so does reaction rate. • This is because k is temperature dependent.

42. 14.5 Temperature and Rate Collision theory • In a chemical reaction, bonds are broken and new bonds are formed. In order for molecules to react, they must collide. • Collisions are either effective or ineffective due to orientation of molecules. • Collisions must have enough energy to overcome the barrier to reaction, the activation energy. • Temperature affects the number of collisions.

43. Molecular Collisions

44. Activation Energy, Ea • Energy barrier (hump) that must be overcome for a chemical reaction to proceed • Activated complex or transition state – arrangement of atoms at the top of the barrier

45. Activation Energy, Ea • Energy difference between the reactant and the highest energy along the reaction pathway • Reaction specific • Rate of reaction is dependent upon the magnitude of Ea; Ea ↓, rate↑ (generally) • Temperature independent

46. Energy Profile Diagram

47. Practice problem

48. Temperature Effects Maxwell-Boltzman Distribution • At higher temperatures, more molecules will have adequate energy to react. • This increases the reaction rate.

49. Arrhenius Equation • Svante Arrhenius developed an equation for the mathematical relationship between k and Ea. • A is the frequency factor, which represents the number of effective collisions.

50. Arrhenius Equation y = m x + b