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Chemical Kinetics CHAPTER 14 Chemistry: The Molecular Nature of Matter, 6 th edition

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  1. Chemical Kinetics CHAPTER 14 Chemistry: The Molecular Nature of Matter, 6th edition By Jesperson, Brady, & Hyslop

  2. CHAPTER 14 Chemical Kinetics • Learning Objectives: • Factors Affecting Reaction Rate: • Concentration • State • Surface Area • Temperature • Catalyst • Collision Theory of Reactions and Effective Collisions • Determining Reaction Order and Rate Law from Data • Integrated Rate Laws • Rate Law  Concentration vs Rate • Integrated Rate Law  Concentration vs Time • Units of Rate Constant and Overall Reaction Order • Half Life vs Rate Constant (1st Order) • Arrhenius Equation • Mechanisms and Rate Laws • Catalysts Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  3. CHAPTER 14 Chemical Kinetics Lecture Road Map: Factors that affect reaction rates Measuring rates of reactions Rate Laws Collision Theory Transition State Theory& Activation Energies Mechanisms Catalysts Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  4. CHAPTER 14 Chemical Kinetics Factors that Affect Reaction Rates Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  5. Kinetics The Speed at Which Reactions Occur Kinetics: Study of factors that govern • How rapidly reactions occur and • How reactants change into products Rate of Reaction: • Speed with which reaction occurs • How quickly reactants disappear and products form Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  6. Kinetics The Speed at Which Reactions Occur • Reaction rate is measured by the amount of product produced or reactants consumed per unit time. • [B] concentration of products will increase over time • [A] concentration of reactants will decrease over time AB Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  7. Kinetics Factors Affecting Reaction Rates 1. Chemical nature of reactants • What elements, compounds, salts are involved? • What bonds must be formed, broken? • What are fundamental differences in chemical reactivity? Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  8. Kinetics Factors Affecting Reaction Rates • Ability of reactants to come in contact (Reactants must meet in order to react) The gas or solution phase facilitates this • Reactants mix and collide with each other easily • Homogeneous reaction • All reactants in same phase • Occurs rapidly • Heterogeneous reaction • Reactants in different phases • Reactants meet only at interface between phases • Surface area determines reaction rate • Increase area, increase rate; decrease area, decrease rate Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  9. Kinetics Factors Affecting Reaction Rates 3. Concentrations of reactants • Rates of both homogeneous and heterogeneous reactions affected by [X ] • Collision rate between A and Bincrease if we increase [A] or increase [B ]. • Often (but not always) reaction rate increases as [X ] increases Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  10. Kinetics Factors Affecting Reaction Rates 4. Temperature • Rates are often very sensitive to temperature • Raising temperature usually makes reaction faster for two reasons: • Faster molecules collide more often and collisions have more energy • Most reactions, even exothermic reactions, require energy to occur • Rule of thumb: Rate doubles if temperatureincreases by 10 °C (10 K) Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  11. Kinetics Factors Affecting Reaction Rates 5. Presence of Catalysts • Substances that increase rates of chemical reactions without being used up • Rate-accelerating agents • Speed up rate dramatically • Rate enhancements of 106 not uncommon • Chemicals that participate in mechanism but are regenerated at the end Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  12. CHAPTER 14 Chemical Kinetics Measuring Reaction Rates Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  13. Rates Measuring Rate of Reaction • Rate = ratio with time unit in denominator • Rate of Chemical Reaction • Change in concentration per unit time. • Always with respect to a given reactant or product • [reactants] decrease with time • [products] increase with time Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  14. Rates Measuring Rate of Reaction • Concentration in M units • Time in s units • Units on rate: • [product] increases by 0.50 mol/L per second  rate = 0.50 M/s • [reactant] decreases by 0.20 mol/L per second  rate = 0.20 M/s Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  15. Rates Rate of Reaction Always positive whether something is increasing or decreasing in [X ] • Reactants • Reactant consumed • So [X ] is negative • Need minus sign to make rate positive • Products • Produced as reaction goes along • So [X ] is positive • Thus rate already positive Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  16. Rates Measuring Rate of Reaction Coefficients indicate the relative rates at which reactants are consumed and products are formed • Related by coefficients in balanced chemical equation • Know rate with respect to one product or reactant • Can use equation to determine rates with respect to all other products and reactants. A + B C + D Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  17. Rates Rate of Reaction: Example C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g) • O2 reacts 5 times as fast as C3H8 • CO2 forms 3 times faster than C3H8 consumed • H2O forms 4/5 as fast as O2 consumed Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  18. Group Problem Clorox bleach is sodium hypochlorite. It should never be mixed with acids, (like vinegar) because it forms chlorine gas: NaClO + 2 HCl → Cl2 + H2O + NaCl If Chlorine gas (Cl2) is formed at a rate of 5.0 x 10-4 mol/Ls what rate is HCl consumed? HCl: Cl2 2:1 Therefore HCl will disappear twice as fast as Cl2 is formed. Rate HCl consumed = 10. x 10-4 mol/Ls Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  19. Rates Change of Reaction Rate with Time Generally reaction rate changes during reaction, it isn’t constant • Often initially fast when lots of reactant present • Slower and slower as reactants are depleted Why? • Rate depends on the concentration of the reactants • Reactants being used up, so the concentration of the reactants are decreasing and therefore the rate decreases Measured in 3 ways: • instantaneous rate, average rate, initial rate Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  20. Rates Instantaneous & Initial Reaction Rate Instantaneous rate • Slope of tangent to curve at some specific time Initial rate • Determined at time = 0 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  21. Rates Average Rate of Reaction Average Rate: Slope of line connecting starting and ending coordinates for specified time frame Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  22. Rates Example Reporting Different Types of Rates Concentration vs. Time Curve for 0.005M phenolphthalein reacting with 0.61 M NaOHat room temperature Rate at any time t = negativeslope (or tangent line) of curve at that point http://chemed.chem.purdue.edu

  23. Rates Example Reporting Different Types of Rates Initial rate = Average rate between first two data points Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  24. Rates Example Reporting Different Types of Rates Instantaneous Rate at 120.4 s (90,0.0028) (160,0.0018) Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  25. Rates Example Reporting Different Types of Rates Average Rate between 0 and 120.4 s Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  26. Group Problem A reaction was of NO2 decomposition was studied. The concentration of NO2 was found to be 0.0258 M at 5 minutes and at 10 minutes the concentration was 0.0097 M. What is the average rate of the reaction between 5 min and 10 min? A. 310 M/min B. 3.2 × 10–3M/min C. 2.7 × 10–3M/min D. 7.1 × 10–3M/min Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  27. CHAPTER 14 Chemical Kinetics Rate Laws Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  28. Rate Laws Rates Based on All Reactants A + BC + D • Rate Law or Rate expression • k is the rate constant • Dependent on Temperature & Solvent • m and n = exponents found experimentally • No necessary connection between stoichiometric coefficients (,) and rate exponents (m, n) • Usually small integers • Sometimes simple fractions (½, ¾) or zero = k[A]m[B]n Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  29. Rate Laws Rates Based on All Reactants Below is the rate law for the reaction 2A +B → 3C rate= 0.045 M–1s–1 [A][B] If the concentration of A is 0.2 M and that of B is 0.3 M, and the reaction is 1st order (m & n = 1) what will be the reaction rate? rate=0.045 M–1 s–1 [0.2][0.3] rate=0.0027 M/s  0.003 M/s Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  30. Rate Laws Order of Reactions Rate = k[A]m[B]n Exponents specify the order of reaction with respect to each reactant Order of Reaction • m = 1 [A]1 1st order in [A] • m = 2 [A]2 2nd order in [A] • m = 3 [A]3 3rd order in [A] • m = 0 [A]0 0th order in [A] [A]0 = 1  means A doesn't affect rate Overall order of reaction = sum of orders (m and n) of each reactant in rate law Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  31. Rate Laws Order of Reactions: Example 5Br– + BrO3– + 6H+ 3Br2 + 3H2O x = 1 y = 1 z = 2 • 1st order in [BrO3–] • 1st order in [Br –] • 2nd order in [H+] • Overall order = 1 + 1 + 2 = 4 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  32. Rate Laws Order of Reaction & Units for k Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  33. Group Problem The following rate law has been observed: Rate = k[H2SeO][I–]3[H+]2. The rate with respect to I– and the overall reaction rate is: A. 6, 2 B. 2, 3 C. 1, 6 D. 3, 6 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  34. Rate Laws Calculating k If we know rate and concentrations, can use rate law to calculate k From Text Example of decomposition of HI at 508 °C • Rate= 2.5 × 10–4M/s • [HI] = 0.0558 M Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  35. Rate Laws Determining Exponents in Rate Law Experimental Determination of Exponents • Method of initial rates • If reaction is sufficiently slow • or have very fast technique • Can measure [A] vs. time at very beginning of reaction • Before it slows very much, then • Set up series of experiments, where initial concentrations vary Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  36. Rate Laws Determining Rate Law Exponents: Example 3A + 2Bproducts Rate = k[A]m[B]n • Convenient to set up experiments so • The concentration of one species is doubled or tripled • And the concentration of all other species are held constant • Tells us effect of [varied species] on initial rate Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  37. Rate Laws Determining Rate Law Exponents • If reaction is 1storder in [X], • Doubling [X]1 21 • Doubles the rate • If reaction is 2ndorder in [X], • Doubling [X]2 22 • Quadruples the rate • If reaction is 0thorder in [X], • Doubling [X]0 20 • Rate doesn't change • If reaction is nthorder in [X] • Doubling [X]n 2n times the initial rate Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  38. Rate Laws Determining Rate Law Exponents: Example • Comparing Expt. 1 and 2 • Doubling [A] • Quadruples rate • Reaction 2ndorder in A = [A]2 2m = 4 or m = 2 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  39. Rate Laws Determining Rate Law Exponents: Example • Comparing Expt. 2 and 3 • Doubling [B] • Rate does not change • Reaction 0thorder in B = [B]0 = 1 2n = 1 or n = 0 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  40. Rate Laws Determining Rate Law Exponents: Example • Conclusion: rate = k[A]2 • Can use data from any experiment to determinek • Let’s choose first experiment Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  41. Rate Laws Determining Rate Law Exponents: Ex 2 2 SO2 + O2 2 SO3 Rate = k[SO2]m[O2]n Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  42. Rate Laws Determining Rate Law Exponents: Ex 2 4 = 2mor m = 2 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  43. Rate Laws Determining Rate Law Exponents: Ex 2 3 = 3n or n = 1 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  44. Rate Laws Determining Rate Law Exponents: Ex 2 Rate = k[SO2]2[O2]1 • 1st order in [O2] • 2nd order in [SO2] • 3rd order overall • Can use any experiment to find k Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  45. Group Problem Using the following experimental data, determine the order with respect to NO and O2 . A. 2, 0 B. 3,1 C. 2, 1 D. 1, 1 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  46. Group Problem Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

  47. Example :Method of Initial Rates BrO3– + 5Br– + 6H+ 3Br2 + 3H2O

  48. Ex.:Method of Initial Rates Compare 1 and 2 Compare 2 and 3

  49. Ex.:Method of Initial Rates Compare 1 and 4 • First order in [BrO3–] and [Br–] • Second order in [H+] • Overall order = m + n + p = 1 + 1 + 2 = 4 • Rate Law is: Rate = k[BrO3–][Br–][H+]2

  50. CHAPTER 14 Chemical Kinetics Integrated Rate Laws Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E