Chemical Kinetics Chapter 15 - PowerPoint PPT Presentation

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Chemical Kinetics Chapter 15

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  1. Chemical KineticsChapter 15 H2O2 decomposition in an insect H2O2 decomposition catalyzed by MnO2

  2. REACTION RATES RR =D [P ] = -D [R ] D t D t P =products R = reactants

  3. Relative Rates Reactant 2A g 4B + C - D [A ] = D [B ] = D [C ] 2D t 4D t D t

  4. Rate Calculations

  5. Collision Theory NO Collisions Collisions Energy NO Collisions YES Energy Orientation

  6. Factors Affecting RXN Rates *Nature of Reactants Temperature Concentration Surface Area/ Physical state Catalysts

  7. Simulation: RATE

  8. MECHANISMSA Microscopic View of ReactionsSections 15.5 and 15.6 Mechanism: how reactants are converted to products at the molecular level. RATE LAW ----> MECHANISM experiment ----> theory

  9. REACTION ORDER In general, for a A + b B --> x X with a catalyst “C” Rate = k [A]m[B]n[C]p The exponents m, n, and p • are the reaction order • can be 0, 1, 2 or fractions • must be determined by experiment!

  10. More on Mechanisms A bimolecular reaction Reaction of cis-butene --> trans-butene is UNIMOLECULAR- only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products Exo- or endothermic?

  11. Collision Theory Reactions require (a) activation energy and (b) correct geometry. O3(g) + NO(g) ---> O2(g) + NO2(g) 1. Activation energy 2. Activation energy and geometry

  12. Mechanisms O3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction.

  13. Mechanisms Most rxns. involve a sequence of elementary steps. 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] NOTE 1. Rate law comes from experiment 2. Order and stoichiometric coefficients not necessarily the same! 3. Rate law reflects all chemistry down to and including the slowest step in multistep reaction.

  14. Mechanisms Most rxns. involve a sequence of elementary steps. 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] Proposed Mechanism Step 1 — slow HOOH + I- --> HOI + OH- Step 2 — fast HOI + I- --> I2 + OH- Step 3 — fast 2 OH- + 2 H+ --> 2 H2O Rate of the reaction controlled by slow step — RATE DETERMINING STEP, rds. Rate can be no faster than rds!

  15. Mechanisms 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] Step 1 — slow HOOH + I- --> HOI + OH- Step 2 — fast HOI + I- --> I2 + OH- Step 3 — fast 2 OH- + 2 H+ --> 2 H2O Elementary Step 1 is bimolecular and involves I- and HOOH. Therefore, this predicts the rate law should be Rate  [I-] [H2O2] — as observed!! The species HOI and OH- are reaction intermediates.

  16. Simulation:” Mechanisms

  17. Rate Laws and Mechanisms NO2 + CO reaction: Rate = k[NO2]2 Two possible mechanisms Two steps: step 1 Single step Two steps: step 2

  18. Ozone Decomposition Mechanism 2 O3 (g) ---> 3 O2 (g) Proposed mechanism Step 1: fast, equilibrium O3 (g) <--> O2 (g) + O (g) Step 2: slow O3 (g) + O (g) ---> 2 O2 (g)

  19. Sovled problems: pg 144