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Chapter 14

Chapter 14. Chemical Kinetics. What does ‘kinetics’ mean?. What can make a reaction speed up or slow down? Chemical nature / contact ability (solids / liquids / gases) Concentration Temperature Presence of a Catalyst

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Chapter 14

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  1. Chapter 14 Chemical Kinetics

  2. What does ‘kinetics’ mean?

  3. What can make a reaction speed up or slow down? Chemical nature / contact ability (solids / liquids / gases) Concentration Temperature Presence of a Catalyst Consider a mixture of gaseous molecules A and B in a sealed container. A and B react with each other; what will happen to the reaction rate if the pressure in the container increases? Factors governing reaction rates:

  4. The Reaction Rate • How fast are products formed / reactants used up? • Reaction of A (red)  B (blue) • Reaction rate =  concentration /  time • Units of rate?

  5. Average rate of disappearance ofA over 20 s? Average rate of formation ofB over 40 s? Calculating Average Rates

  6. Instantaneous Rates C4H9Cl (aq) + H2O (l) C4H9OH (aq) + HCl (aq) • Look at the rate of a reaction over time: Is the rate constant? Tangent lines show the rates at specified times (instantaneous rates) What is the instantaneous rate at t=600s? Units of rate?

  7. Looked at one component of a reaction Rate expression that includes all reactants A + B  Products Rate  [A]m [B]n (exponents m and n determined experimentally- look at very soon) Rate Law: Shows how initial reaction rate changes wrt concentration of species involved; introduces rate constant, k Rate = k [A]m[B]n Calculate rate with k, m and n (any concentration of A and B) Instantaneous Rates

  8. Reaction Rates and Stoichiometry • C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) propane • In a particular experiment, rate of loss of C3H8 is 3.55 x 10-3 mol / L.s • How can we calculate the corresponding rate of formation of CO2? • Rate of formation of H2O?

  9. Reaction Rates and Stoichiometry • Dinitrogen pentoxide (N2O5) decomposes to NO2 and O2 (all are gases). Balanced equation for this? • Rate of decomposition of N2O5 at a particular instant is 4.2 x 10-7 M/s – what are the rates of formation of NO2 and O2?

  10. Rate potentially depends on both [NH4+] and [NO2-] Rate Law: Rate = k [NH4+]x [NO2-]y x and y (exponents):indicate how the initial rate is affected by change in concentration of the reactants. Can be determined. k = Rate constant (shows how temperature affects rate) The Rate Law NH4+(aq) + NO2-(aq)  N2(g) + 2H2O(l)

  11. A + B  C; Rate = k[A]n[B]m; values of n and m? Vary concs. of A and B separately, and observe effect on rate. Expt. # Initial Concs (mol/L) Initial Rate (mol L-1 s-1) [A] [B] 1 0.10 0.10 4.0 x 10-5 2 0.10 0.20 4.0 x 10-5 3 0.20 0.10 16.0 x 10-5 Rate Law: Rate = k What is the value of the rate constant? (remember units!) What is the rate of the reaction when [A] = 0.050 M, and [B] = 0.010 M? Exponents and Rate Laws

  12. What is the rate law? • Consider the reaction of nitric oxide with hydrogen: 2 NO(g) + 2H2(g)  N2(g) + 2H2O(g) Expt. # Initial Concs (mol/L) Initial Rate (mol L-1 s-1) [NO] [H2] 1 0.10 0.10 1.23 x 10-3 2 0.10 0.20 2.46 x 10-3 3 0.20 0.10 4.92 x 10-3 (a) What is the rate law for this reaction? (b) Calculate the rate constant. (c) Calculate the rate when [NO]=0.050 M and [H2] = 0.150 M.

  13. The rate constant, k • Let’s compare the rate constants (and their units) for the previous two problems just worked. • What do you notice about the two that looks a little odd?

  14. Rate Law: Shows how initial rate varies with concentration. Integrated Rate Law: Shows what the concentration of products or reactants is at a particular time during the reaction. 1st order reactions: Rate = - [A] / t = k [A] ln [A]t = -kt + ln[A]o 2N2O5 2N2O4 + O2; Rate = k[N2O5] At 45 °C, k = 6.22 x 10-4 s-1. If initial conc. Of N2O5 is 0.10M, how many minutes for [N2O5] to drop to 0.01 M? The Integrated Rate Law

  15. 2nd order reactions; Rate = - [A] / t = k [A]2 1/ [A]t = kt + 1/[A]o Decomposition of NO2(g)  NO(g) + ½ O2(g) 2nd order in NO2, with k = 0.543 L / mol.s. [NO2]o = 0.050 M What is the remaining concentration after 0.500 h? Integrated Rate Law – 2nd order

  16. Graphical representations • How to tell if a reaction is 1st order: • Insert Figure 14.8

  17. Graphical Representations • How to tell if a reaction is 2nd order:

  18. Half-life of a particular chemical t for ½ reactant to disappear 3g A  t1/2 = 0.693 / k Half-life (t1/2) – 1st order reactions

  19. t1/2 = 0.693 / k 131I; t1/2 = 8 days. 10g sample of 131I, after 8 days….decayed down to…. Radioactive emission (-, -, -, X-rays) to stable isotope Fraction of 131I present after 24 days? Amount of 131I? Half-life (t1/2) – 1st order reactions

  20. t1/2 = 1/ k[A]o 2HI(g)  I2(g) + H2(g); Rate = k[HI]2; k = 0.079 Lmol-1 s-1 Initial conc. of HI = 0.050M; t1/2 = ? Half-life (t1/2) – 2nd order reactions

  21. Temperature and Rate • Value of k varies with change in T • Light sticks in hot vs. ice water

  22. Model for Chemical Kinetics Collision Model (Collision Theory) Molecules must successfully collide to react. 2 reasons for ineffective collisions: Insufficient K.E to overcome Ea (Activation Energy) barrier Orientation issues How do reactions occur?

  23. Orientation Issues Cl + NOCl  NO + Cl2

  24. Activation Energy (Ea)

  25. Changing T changes rate; also changes k Extent of k variation with T depends on Ea. To calculate Ea for a reaction: ln (k1/k2) = -Ea/RT (1/T2 – 1/T1) Decomposition of HI has rate constants: k = 0.079 L mol-1 s-1 at 508 °C k = 0.24 L mol-1 s-1 at 540 °C What is the Ea for the reaction, in kJ / mol? Activation Energy (Ea)

  26. What happens during a chemical reaction? NO(g) + O3(g)  NO2(g) + O2(g) Occurs in a single step. Rate Law = k This is an elementary reaction What is the molecularity of this elementary reaction? NO(g) + Cl2(g)  NOCl(g) + Cl(g) The Reaction Mechanism

  27. Multi-Step Reactions NO2(g) + CO(g)  NO(g) + CO2(g) Rate = k[NO2]2; Single step? Possible Mechanism: Step 1: NO2(g) + NO2(g)  NO3(g) + NO(g) Step 2: NO3(g) + CO(g)  NO2(g) + CO2(g) If mechanism is correct, Step 1 + Step 2 = overall reaction Intermediates? Rate Determining Step? The Reaction Mechanism

  28. Decomposition of N2O (nitrous oxide) occurs by a 2-step mechanism: Step 1: N2O(g)  N2(g) + O(g) (slow) Step 2: N2O(g) + O(g)  N2(g) + O2(g) (fast) What is the overall equation for the reaction? Write the rate law for the reaction. Mechanism problem

  29. Change rates of reactions without being used up Lowers activation energy of reaction (changing mechanism) Homogeneous catalysts Heterogeneous catalysts Catalysts

  30. Catalytic Convertors

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