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Chapter 22

Chapter 22. Kinetics and Equilibrium. Reaction order. Reaction Rates. Rate laws. Collision theory. Reversible Reactions. Reaction mechanisms. Equilibrium. intermediate. Factors affecting reaction rate. Activated complex. Equilibrium constant. concentration. Factors affecting

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Chapter 22

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  1. Chapter 22 Kinetics and Equilibrium

  2. Reaction order Reaction Rates Rate laws Collision theory Reversible Reactions Reaction mechanisms Equilibrium intermediate Factors affecting reaction rate Activated complex Equilibrium constant concentration Factors affecting equilibrium Evaluating Keq Activation energy temperature Reaction DGo catalysts Surface area LeChatelier’s principle inhibitors

  3. Reaction Rates • Rate = change in property/change in time • Speed: miles/hour • Ddistance/Dtime • What could we measure for a reaction? • Change in mass of a solid • Change in concentration • Temperature changes • pH changes • Gas volume changes • Color changes • We must also measure changes in TIME!

  4. Writing Rate Expressions • For a general reaction aA + bB  cC + dD • General form: • If you are using the same units: regardless of property measured, the overall rate of reaction is the same! • Need to account for coefficients • Convention: all reaction rates are positive

  5. Measuring Rates • Average Rate • Initial Rate • Calculate average rate for early part of data when plot is nearly linear • Instantaneous Rate • Draw a tangent to the curve and determine its slope

  6. What happens to the rate over time? • Compare average rate at beginning vs. average rate at end • Reaction rates typically slow down over time • Why? • There are fewer moles of reactants left, and therefore fewer collisions.

  7. Collision Theory • Molecules must collide in order to react. • They must collide with the correct orientation. • “Effective collision” • Has appropriate orientation; molecules may react. • “Ineffective collision” • Doesn’t have needed orientation; particles will separate.

  8. Collision Theory, cont. • Molecules must collide in order to react. • They must have enough energy to react. • Activation Energy, Ea • The minimum energy that reactants must have for the reaction to occur

  9. Potential Energy Diagrams • Activation Energy: from reactants to top of “hill” • Transition State • Aka Activated Complex • High energy state, where bonds are broken and new bonds are formed • DHrxn = energy of products – energy of reactants

  10. Potential Energy Diagrams • Reactions with a smaller activation energy will occur more quickly than reactions with a larger Ea. Which reaction would you expect to be fastest? Slowest? Why?

  11. Collision Theory • Basic premise: • More collisions = faster reaction rate • More collisions = greater likelihood for effective collisions

  12. How can we speed up the rate of a reaction? • Increase temperature • Particles move more quickly, so more possible collisions • More particles are likely to have enough energy to overcome activation energy barrier • Increase concentration • More particles, so more possible collisions • Increase surface area • More particles are exposed, so more collisions are possible

  13. Catalysts • Speed up reaction rates, without being consumed • Homogeneous vs. heterogeneous catalysts • Enzymes • Catalytic RNA • Catalytic antibodies • Catalytic converter in car engine • Effectively lower the activation energy of the reaction • May even change the mechanism of the reaction

  14. Inhibitors • Cause reaction rate to slow down • May interfere with enzyme active site • May “tie up” reactant in a complex so that it does not react • Food preservatives • Important in biological feedback mechanisms • Thyroid function • Maintaining body temperature

  15. Rate Laws • The rate of the reaction depends in part on the concentrations of the reactants. • Rate law = an expression relating reaction rate to reactant concentration • Rate is directly proportional to reactant concentration (in molarity) raised to some power, n. • Rate  [A]n • n must be determined experimentally • n can be any integer (including zero), or a fraction

  16. Rate Laws, cont. • We can express this proportionality by including a “rate constant,” k, which is experimentally determined. The value of k depends on the reaction conditions. • A general rate law: • Rate = k[A]n • n indicates the extent the rate depends on the concentration • Called the “order of reaction” with respect to A • Rate laws must be experimentally determined by examining the relationship between rate and concentration

  17. Reaction Order

  18. Example • A  B • As [A] doubles (2-fold change) from 0.050 M to 0.10 M, the rate quadruples (4-fold change) • 4/2 = 2, so the reaction is second order in A • Rate = k[A]2 • Can find k by substituting rate and concentration from any trial

  19. If more than one reactant is involved… • Example: A + B  C • Rate = k [A]n[B]m • The order of each reactant must be determined independently • Hold one constant, change the other; how is the rate affected? • The overall order of the reaction is the sum of the exponents n + m

  20. Example:NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l) • Trials 1-4: hold [NH4+] constant in large excess, vary [NO2-] • What is the reaction order for [NO2-]? • Trials 5-8: hold [NO2-] constant in large excess, vary [NH4+] • What is the reaction order for [NH4+]?

  21. Example:NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l) • Compare trials 1 and 2 • When [NO2-] doubles, the reaction rate doubles • 2/2 = 1 • Reaction is first order in [NO2-]

  22. Example:NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l) • Compare trials 5 and 6 • When [NH4+] doubles, rate doubles • 2/2 = 1 • Reaction is first order in [NH4+]

  23. Example: • NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l) Based on the experimental data, • Rate = k [NO2-]1 [NH4+]1 • The reaction is first order in ammonium ion, first order in nitrite ion, and second order overall

  24. How does a reaction take place? • Most reactions occur as a series of simple steps • “elementary steps” • unimolecular • Bimolecular • Termolecular (rare--unlikely that three particles will collide simultaneously)

  25. Reaction Mechanism • A step by step description of the steps that occur in a chemical reaction • Includes all of the elementary steps that add up to the overall reaction • Proposed mechanism must be consistent with experimental rate law

  26. Chemical Intermediates • Made in one step, consumed in another step • Consider the reaction A  D • In this potential energy diagram, B and C are intermediates • Three activated complexes • Possible mechanism: • A  B • B  C • C  D

  27. Rate limiting step • The slow step in a multi-step mechanism • “rate determining step” • RLS • Has the greatest effect on the rate of a multi-step mechanism • Molecularity of RLS should match the experimentally determined reaction order

  28. Why study reaction kinetics? • Even “simple” reactions can have multi-step mechanisms. • A balanced equation gives the reactants and products, and stoichiometric relationships, but NO direct information on how reactants become products • Goal of kinetics: eliminate unlikely mechanisms

  29. Example • Consider the reaction: NO2(g) + CO(g)  NO(g) + CO2 • From experiments, we know that Rate = k[NO2]2

  30. Proposed mechanism • Step 1: (slow) NO2(g) + NO2(g)  NO3(g) + NO(g) • Step 2: (fast) NO3(g) + CO(g)  NO2(g) + CO2(g)

  31. Proposed mechanism • Step 1: (slow) NO2(g) + NO2(g)  NO3(g) + NO(g) • Step 2: (fast) NO3(g) + CO(g)  NO2(g) + CO2(g) • Sum of steps 1 and 2: NO2(g) + CO(g)  NO(g) + CO2

  32. Example, cont. • What is the intermediate? • How do you recognize a reaction intermediate? • How many activated complexes must form? • Which step is the RLS? • If additional CO is added, what effect will this have on the observed rate? • If additional NO2 is added, what effect will this have on the observed rate?

  33. Example, cont. Note that the sum of the steps in the mechanism adds up to the overall, balanced equation • Note that the slow step of the mechanism is bimolecular (2 colliding molecules), and the overall reaction order is second order • This mechanism is consistent with the experimental rate law • We can say a mechanism is possibly correct, and we can disprove mechanisms. Deducing mechanisms is difficult!

  34. Reversible Reactions • Some reactions essentially go “to completion” or “to an end” • Form a precipitate • Form a gas • NaCl(aq) + AgNO3(aq)  NaNO3(aq) + AgCl(s) • Many processes are reversible • Can be returned to their original state • Melting ice/refreezing water • Rechargeable batteries

  35. Writing reversible processes • Use a double-headed arrow to indicate reversible processes • sugar(s) ↔ sugar(aq) • Two reactions are occuring: • Sugar(s)  sugar(aq) • Sugar(aq)  sugar(s) • Both the forward and reverse processes occur simultaneously

  36. Equilibrium • Eventually, the forward and reverse reactions may reach a balance point. • At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. • Both reactions continue to occur. • No observable, macroscopic changes will be visible.

  37. Equilibrium • At equilibrium, the ratio of product to reactant is a constant • This ratio is called the “equilibrium constant” • Square brackets indicate molar concentrations

  38. Equilibrium Constant • If Keq > 1 there are more products than reactants present at equilibrium • “products are favored” • If Keq < 1 there are more reactants than products present at equilibrium • “reactants are favored”

  39. Types of equilibria • Phase equilibrium • Ex. Solid ↔ liquid • Solution equilibrium • Ex. Dissolving a solid in a solute • Chemical equilibrium • Ex. H2(g) + I2(g) ↔ 2 HI(g)

  40. Law of Mass Action • To write Keq expressions • Start with a balanced equation • Ignore solids and pure liquids • In the numerator, the products are multiplied together • Coefficients become exponents • In the denominator, the reactants are multiplied together • Coefficients become exponents

  41. Law of Mass Action • For the equation aA + bB ↔ cC + dD

  42. Law of Mass Action • Reaction: 2 SO2(g) + O2(g) ↔ 2 SO3(g)

  43. The equilibrium constant • Each reaction has its own characteristic Keq value • The only factor that affects the value of Keq is temperature • Many combinations of concentrations can result in the same value for Keq • Treat Keq as if it were dimensionless

  44. Reaction quotient, Q • Use to predict how a system will adjust to reach equilibrium • Use non-equilibrium/current concentrations in Keq expression • Comparing Q to Keq predicts how non-equilibrated system will respond to achieve equilibrium

  45. Using the reaction quotient • If Q = Keq, the system is at equilibrium • If Q > Keq, products must react to form reactants in order to reach equilibrium • Too many products! • System must “shift left” • If Q < Keq, reactants must form products to reach equilibrium • Too many reactants! • System must “shift right”

  46. Example • Consider the synthesis of ammonia at 500oC: • N2(g) + 3 H2(g) ↔ 2 NH3(g) Keq = 6.0 x 10-2 • If [NH3] = [N2] = 1.0 x 10-3 M and [H2] = 2.0 x 10-3 M, is the system at equilibrium? If not, how will the system shift to react equilibrium?

  47. Example • Step 1: Write a Keq expression • The value of Keq is given • Step 2: Calculate Q • Substitute the given concentrations into Keq expression and evaluate • Step 3: Compare Q to Keq • Therefore, the system will shift left to make more reactants

  48. Factors affecting equilibrium • A system at equilibrium will stay at equilibrium. • Systems can be stressed in a number of ways: • Adding reactants or products • Removing reactants or products • Changing the temperature • Changing the pressure • Reactions can’t do yoga, but they do need a way to deal with the stress

  49. Henry-Louis LeChâtelier • Trained as a mining engineer • Became a professor of chemistry at the École des Mines (1877), of mineral chemistry at the Collège de France (1898) and finally of chemistry at the Sorbonne (1907) • An authority on metallurgy, cements, glasses, fuels and explosives. • Developed a platinum-rhodium thermocouple for measuring high temperatures • Suggested the use of the oxyacetylene flame for welding and metal cutting.

  50. LeChâtelier’s Principle • When a system at equilibrium is disturbed by the application of a stress, the system will attain a new equilibrium position that minimizes the stress. • Two ways to approach it: • Reaction rates • “see saw” analogy

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