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Chemical Kinetics

Chemical Kinetics. BLB 11 th Chapter 14. 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. Expectations. Review: concentration units, graphing (line, slope) Work with reaction rates.

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Chemical Kinetics

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  1. Chemical Kinetics BLB 11thChapter 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. What happens to concentration as the reaction proceeds?N2(g) + 3 H2(g) → 2 NH3(g)

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

  7. 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.

  8. Reaction rates vary significantly.

  9. 14.2 Reaction Rates • Rate – change in some variable per unit time • Reaction rate – change in concentration (M), moles, or pressure per unit time; M/s or M·s-1 • Rates are determined by monitoring concentration as a function of time. • Rates are positive quantities; for reactant A:

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

  11. 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.

  12. C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) (Fig. 14.4, p. 561)

  13. 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)

  14. C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) t=0 0.1000 t=0 0 t=50 0.0905 t=50 ? Δ= Δ=

  15. 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

  16. 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 ? ? ?

  17. Types of Rate Laws • Differential rate law or rate law (section 3) • Shows how the initial reaction rate changes with starting concentrations • Must be specific in how defined (temp., etc.) • Initial rates used to determine reactant order • Integrated rate law (section 4) • Shows how concentration changes with time • Graphical determination of the reaction order

  18. 14.3 Concentration & Rate Laws NH4+(aq) + NO2-(aq) → N2(g) + 2 H2O(l) This relationship is summarized as a rate law. rate = k[NH4+][NO2-]

  19. Rate Laws 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 reaction order = sum of individual orders • Rate constant is independent of concentration.

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

  21. Units of the rate constant, k

  22. 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?

  23. 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?

  24. Practice problem 30 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] = 0.035 M and [H2] = 0.015 M.

  25. 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). See samples on pp. 568-9.

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

  27. 2 NO(g) + Cl2(g) → 2 NOCl(g) Determine the rate law and calculatek.

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

  29. More Practice

  30. 14.4 The Change of Concentration with TimeaA→ products • 1st order y = mx + b Plot: ln[A] vs. t slope = −k integrate

  31. 14.4 The Change of Concentration with TimeaA→ products • 2nd order y = mx + b Plot: 1/[A] vs. t slope = k integrate

  32. 14.4 The Change of Concentration with TimeaA → products • Zero-order y = mx + b Plot: [A] vs. t slope = −k integrate

  33. Integrated rate law analysis (BLB 48)isomerization: CH3NC(g) → CH3CN(g) To Excel for analysis

  34. Integrated rate law analysis (BLB 52)C12H22O11(aq) + HCl(l) → 2 C6H12O6(aq) To Excel for analysis

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

  36. 1st order half-life • All half-lives same length of time • Independent of initial concentration Note: Radioactive decay follows 1st order kinetics.

  37. 1st order reaction

  38. 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

  39. BLB 43a

  40. BLB 45

  41. For the reaction A → B, k = 4.15x10-3 M-1·s-1. If the initial concentration of A is 0.100 M, how many seconds will it take for the concentration to decrease to 0.0250?

  42. To Nuclear Chemistry, Ch. 21

  43. 14.5 Temperature and Rate • Generally, as temperature increases, so does reaction rate. • This is because k is temperature dependent.

  44. Rate constant and TemperatureCH3NC(g) → CH3CN(g)

  45. 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.

  46. Molecular Collisions Cl + NOCl→NO + Cl2

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