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Chemical Kinetics

http://college.hmco.com/chemistry/shared/media/animations/oscillatingreaction.html. Chemical Kinetics. Concerns reaction rates: Speed with which reactants are converted to products.

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Chemical Kinetics

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  1. http://college.hmco.com/chemistry/shared/media/animations/oscillatingreaction.htmlhttp://college.hmco.com/chemistry/shared/media/animations/oscillatingreaction.html Chemical Kinetics Concerns reaction rates: Speed with which reactants are converted to products

  2. Fireworks explode/give bright colors due to very fast chemical reactionsThe rusting of a bridge is also a chemical reaction, but it is very slow. The eroding of a mountain takes even longer

  3. Rate: the speed at which something happens Ratechem reaction = ∆ [reactant] or [product] = x mol unit time L⋅s • Always positive • [ ] = molar concentration • Forward Rate: reactants  products • Reverse Rate: products  reactants • Net Rate: forward rate - reverse rate • Average Rate: speed of entire reaction from start to finish • Instantaneous Rate: speed of reaction at one moment in time

  4. The Collision Theory

  5. Atoms, ions, and molecules (reacting substances) must collide in order to react:

  6. Activated complex: temporary, unstable arrangement of atoms that may form products or may break apart to reform reactants • Transition state http://college.hmco.com/chemistry/shared/media/animations/uctransitionstates.html

  7. Enough kinetic energy for reactants to leap over reaction barrier?

  8. Exothermic reactions release energy and form products at lower energy levelEndothermic reactions absorb energy and form products at higher energy level ∆H = (-)

  9. Covers period of time between mixing of reactants and point at which chemical reaction stops or reaches equilibrium How fast chemicals react and factors that influence rate allows chemists to exercise precise control over chemical reactions

  10. Chemical Kinetics • Reaction mechanism • Theory that explains observed products • Understand steps by which reaction takes place • Enables us to predict outcome of similar reactions • Explains how covalent bonds break/form in process • Shows how electron pairs are involved in process • Rate dictates whether reaction can occur • Can adjust reaction conditions to get more suitable rate • If 2 reactions competing for single reagent, lets you favor exclusive formation of single product

  11. Measurement of reaction rate based on rate of appearance of product or disappearance of reactant • Determined by measuring concentration of one or more chemicals at different times during course of reaction • Change can be • Disappearance of reactants (decrease, - rate) • Appearance of product(increase, + rate) • Always define rate as positive quantity • [ ] indicate concentration in mol/L

  12. A  B Rateaverage = -Δ[A] = +Δ[B] Δt Δt RateA disappearing = RateB appearing Rate of appearance (or disappearance) of substance is divided by its stoichiometric coefficient aA + bB  cC + dD Rate = -ΔA = -ΔB = +ΔC = +ΔD aΔt bΔt cΔt dΔt

  13. 2A + B  ½ [rate of disappearance of A] = [rate of disappearance of B] = [rate of appearance of C] Rate of reaction = - 1Δ[A] = - Δ[B] = + Δ[C] 2 ΔtΔtΔt H2 + I2 ↔ 2HI Rate of formation of HI is twice rate of disappearance of H2 or I2 at any given time

  14. 2NO2(g) 2NO(g) + O2(g) Rate of consumption of NO2 = Rate of production of NO = 2(Rate of production of O2) because rate of production of NO is twice that of O2 –Δ[NO2] = Δ[NO] = (Δ[O2] ) 2Δt 2Δt Δt (multiply all by 2) –Δ[NO2] = Δ[NO] = 2(Δ[O2] ) Δt Δt Δt

  15. Factors affecting rate of reaction • Physical state • Liquid/liquid or gas/gas: maximum opportunity to collide • Solid: greater surface area/smaller size of particles, greater area that reaction can take place in • Concentration and pressure • If either increases, more particles within given space • More collisions • Rate of reaction increases

  16. Temperature • Increase in speed of particles means more successful collisions with particles having required activation energy • For many reactions occurring at around room T, rate of reaction doubles for every 10°C (9/11°C rise in T) • # degrees needed to double rate changes gradually as T  • Catalyst • Changes energy pathway for chemical reaction • Provides alternate mechanism that lowers activation energy • More particles have required energy for successful collision

  17. Entropy & temperature • Place 3 drops of blue food coloring in 3 flasks. • Place 3 drops of yellow food coloring in 3 more. • Keep 1 of each at room temperature, and warm 1, and cool the other. • Using note cards, invert the blue flask over the yellow and remove the card. Secure the flask with a clamp. • Allow the flasks to stand and record the time necessary for both flasks to become the same shade of green. • Is the entropy greatest when the colors of the flasks are different or the same? Explain. • Explain the relationship between diffusion and entropy. • What influence does an increase/decrease in temperature have on the diffusion observed?

  18. Homework: Read 17.1-17.2, pp. 529-541 Q pg. 535, #6 Q pp. 554-555, #34, 37, 48-49, 62

  19. Rate Laws The rate of a chemical reaction dictates whether a reaction can occur

  20. Rate Laws: An introduction • Chemical reactions are reversible • With time, enough products accumulate • Reverse reaction becomes important • Concentration of reactants depends on difference in rates of forward/reverse reactions • Reaction rate depends only on [ ] of reactants

  21. Rate = k[A]x[B]y[C]z Must be determined by experiment x/y/z: reaction orders May be zero, positive, negative, integer, or fraction Must be determined by laboratory experiments No relationship to stoichiometric coefficients of balanced chemical equation Proportionality constant "k“ (rate constant) Relates [ ] and orders to rate of reaction Units of rate constant depend on order of reaction

  22. Orders of reactions • If reaction is zero order with respect to reactant • Rate does not depend on concentration of that reactant • Doubling concentration will not increase or decrease rate • If first order with respect to reactant • Rate directly proportional to concentration of that reactant • Doubling concentration doubles rate • If second order with respect to a reactant • Rate is directly proportional to square of concentration of that reactant • Doubling concentration quadruples rate • Overall order of reaction: sum of individual reaction orders

  23. http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=chapter_14&folder=reaction_orderhttp://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=chapter_14&folder=reaction_order

  24. NH4+ + NO2-® N2 + 2 H2O rate = k[NH4+]1[NO2-]1 or simply k[NH4+][NO2-] 1.35 x 10-7 M/s = k(0.100 M)(0.005 M) solving, k = 2.7 x 10-4/Ms or 2.7 x 10-4 L/mol·s

  25. NO2 + CO ® NO + CO2 rate = k[NO2]2[CO]0 or simply k[NO2]2 0.18 M/s = k(0.10 M)2(0.0010 M)0 k = 18/Ms or 18 L/mol·s

  26. 2 NO2 + 2 H2 ® N2 + 2 H2O rate = k[NO2]2[H2]1 or simply k[NO2]2[H2] 1.3 x 10-5 M/s = k(0.0050 M)2(0.0020 M)1 k = 2.6 x 102/M2s

  27. Example • The initial rate of the reaction BrO3-(aq) + 5 Br-(aq) + 8 H+(aq)  3 Br2(l) + H2O(l)  has been measured at the reactant concentrations shown (in mol/L):  • Experiment [BrO3-] [Br-]     [H+] Initial rate (mol/Ls)         1           0.10    0.10    0.10     8.0 x 10-4       2           0.20    0.10    0.10     1.6 x 10-3       3           0.10    0.20    0.10     1.6 x 10-3       4           0.10    0.10    0.20     3.2 x 10-3 • Rate = k[A]x[B]y so Rate = k[BrO3-]1[Br-]1[H+]2 • Rate = 1 + 1 + 2 = 4

  28. Example • The reaction of iodide ion with hypochlorite ion, OCl- (which is found in liquid bleach), follows the equation  OCl- + I-  OI- + Cl- • It is a rapid reaction that gives the following rate data: Initial Concentrations      Rate of Formation (mol/Ls)  [OCl-]            [I-]           [Cl-] 1.7 X 10-3   1.7 X 10-3        1.75 X 1043.4 X 10-3    1.7 X 10-3        3.50 X 1041.7 X 10-3    3.4 X 10-3        3.50 X 104 • Determine value of rate constant.  Rate = k[A]x[B]y • 1.75 x 104 = k(1.7 x 10-3)1 (1.7 x 10-3)1 • rate = 6.06 x 109 L/mol sec [OCl-] [I-] • Overall Rate = 1 + 1 = 2

  29. Homework: • Read 17.3, pp. 542-545 • Q pg. 545, #20 • Q pp. 554-556, #56, 69-70, 74

  30. Complex reaction: two or more elementary steps • Intermediates usually so reactive, can’t be isolated • Produced in one elementary step and consumed in another • Do not appear in overall reaction • Determines overall rate of reaction

  31. Use rate laws to determine mechanism for • A + B  Y + Z • Overall reaction rate expressed by Rate = k[B] • Predict if following mechanism is valid or invalid • Step 1: 2A  Y + D Slow process • Step 2: B + D  A + Z Fast process A + B  Y + Z • Rate law for rate-determining step should match rate law for overall reaction • 2A  Y + D • Rate = k[A]2 (doesn’t match rate = k[B])-invalid

  32. Homework: • Read 17.4, pp. 546-549 • Q pp. 554-556-#57, 71, 76 • Do test practice, pg. 557 • Use link for quiz and submit as before. http://www.glencoe.com/qe/science.php?qi=1001

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