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MECHANISMS A Microscopic View of Reactions Sections 15.5 and 15.6

MECHANISMS A Microscopic View of Reactions Sections 15.5 and 15.6. How are reactants converted to products at the molecular level? Want to connect the RATE LAW ----> MECHANISM experiment ----> theory. MECHANISMS. For example Rate = k [trans-2-butene]

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MECHANISMS A Microscopic View of Reactions Sections 15.5 and 15.6

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  1. MECHANISMSA Microscopic View of ReactionsSections 15.5 and 15.6 How are reactants converted to products at the molecular level? Want to connect the RATE LAW ----> MECHANISM experiment ----> theory

  2. MECHANISMS For example Rate = k [trans-2-butene] Conversion requires twisting around the C=C bond.

  3. MECHANISMS Conversion of trans to cis butene

  4. MECHANISMS Energy involved in conversion of trans to cis butene See Figure 15.15

  5. MECHANISMS Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ACTIVATION ENERGY, Ea = energy req’d to form activated complex. Here Ea = 233 kJ/mol

  6. MECHANISMS Also note that trans-butene is MORE STABLE than cis-butene by about 4 kJ/mol. Therefore, trans ---> cis is ENDOTHERMIC. This is the connection between thermo-dynamics and kinetics.

  7. Activation Energy A flask full of trans-butene is stable because only a tiny fraction of trans molecules have enough energy to convert to cis. In general, differences in activation energy are the reason reactions vary from fast to slow.

  8. MECHANISMS 1. Why is reaction observed to be 1st order? As [trans] doubles, number of molecules with enough E also doubles. 2. Why is the reaction faster at higher temperature? Fraction of molecules with sufficient activation energy increases with T.

  9. MECHANISMS Reaction of trans --> cis is UNIMOLECULAR- only one reactant is involved.

  10. MECHANISMS Reaction of trans --> cisis UNIMOLECULAR- only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products

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

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

  13. 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

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

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

  16. MECHANISMS Most rxns. involve a sequence of elementary steps. 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 Rate of the reaction controlled by slow step — RATE DETERMINING STEP, rds. Rate can be no faster than rds!

  17. 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 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.

  18. Arrhenius Equation • Reaction rates depend on energy, frequency of collisions, temperature, and geometry of molecules given by: • A = frequency of collisions with correct geometry at concentration of 1M (L/mol*s) • R = gas constant (8.314 x 10-3 kJ/K*mol) • e-Ea/RT is fraction of molecules having the minimum energy required for reaction

  19. Arrhenius Equation • Calculate the value of the activation energy from the temp. dependence of the rate constant • Calculate the rate constant for a given temp. (if activation energy and A are known)

  20. Arrhenius Equation • Taking the natural log and rearranging: • Straight line plot of ln k vs 1/T • Slope of –Ea/R

  21. CATALYSIS Catalysts speed up reactions by altering the mechanism to lower the activation energy barrier.

  22. CATALYSIS Catalysts speed up reactions by altering the mechanism to lower the activation energy barrier. Dr. James Cusumano, Catalytica Inc. What is a catalyst? Catalysts and society

  23. CATALYSIS In auto exhaust systems — Pt, NiO 2 CO + O2 ---> 2 CO2 2 NO ---> N2 + O2

  24. CATALYSIS 2. Polymers: H2C=CH2 ---> polyethylene 3. Acetic acid: CH3OH + CO --> CH3CO2H 4. Enzymes — biological catalysts

  25. CATALYSIS MnO2 catalyzes decomposition of H2O2 2 H2O2 ---> 2 H2O + O2 Catalysis and activation energy Uncatalyzed reaction Catalyzed reaction

  26. Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.18

  27. Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.19

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