1 / 61

Chemical Kinetics

Chemical Kinetics. Chapter 16. Thermodynamics Vs. Kinetics. Thermodynamics - Will the reaction happen under specified Conditions? Thermodynamics and Equilibrium - What will be the extent of the reaction? Kinetics - How quickly will the reaction occur?

cole
Download Presentation

Chemical Kinetics

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemical Kinetics Chapter 16

  2. Thermodynamics Vs. Kinetics • Thermodynamics - Will the reaction happen under specified Conditions? • Thermodynamics and Equilibrium - What will be the extent of the reaction? • Kinetics - How quickly will the reaction occur? • Kinetics - What factors will affect the rate of reaction:?

  3. The Rate of Reaction • Rate of Reaction – describes how fast reactants are used up and how products are formed • Chemical Kinetics – The study of the rates of reactions, what affects them, and the mechanisms (steps) in which they occur.

  4. The Rate of Reactions • Rates are expressed as (molarity / time) • Problem – you need a way to track the change in molarity over time! • Titration – following the acid concentration • Light absorption over time may change • Change in pressure over time

  5. The Rate of Reaction • One Step Mechanism – simplest Case • A(g)  B(g) + C (g) • The rate is proportional to the concentration of the reactant • R  [A] or R = k[A] • K= specific rate constant • 2nd equation only valid at given temperature

  6. The Rate of Reaction • 2A(g)  B(g) + C(g) (coefficients not one) • R  [A]2 orR = k[A]2 • The reaction is second order with respect to A • This relationship is found experimentally • These relationships are only true for simple one step mechanisms – most are NOT

  7. The Rate of Reaction • Rate Law Expressions – found only through experimentation, not through inspection of balanced equations. • Variation on rate law expression: • aA + bB  cC + dD • (-1/a) [A] / t • (-1/b) [B] / t • ( 1/c) [C] / t • ( 1/d) [D] / t

  8. Factors Affecting the Rate of Reaction • The nature of the reactants • The state of matter (temperature related) • Allotropic Form of matter • Diamond vs. graphite, similar G values, but oxidation of graphite is very rapid

  9. Factors Affecting the Rate of Reaction • Chemical Identity • Mg vs. Na in water (sodium has lower ionization energy) • Particle size of the solid • Greater surface area in smaller particles can speed up the reaction • Pulverize the solid • Make an aqueous solution • Evaporate a liquid

  10. Factors Affecting Reaction Rate • Concentration – effect is summarized in the rate law expression. • Increasing the concentration of reactants increases the frequency of the collisions and therefore affects the rate of reactions

  11. The Rate of Reaction • 2A(g) + B (g)  3C(g) • R = k [A]x [B]y • X is the order of reaction with respect to A • Y is the order of reaction with respect to B • The overall rate of reaction is x + y

  12. As slope changes, so does the rate of reaction. Concentration affects the rate. Rate is considered an instantaneous measurement. Generally, we measure the initial rate of reaction. Factors Affecting Reaction Rate

  13. Example: aA + bB  cC

  14. Example • What is the order of reaction with respect to [A]? • What is the order of reaction with respect to [B]? • If the rate is reported as the rate of formation of C, what would be the rate of disappearance of A?

  15. Example 2: aA + bB  cC

  16. Example 2 • What is the order of reaction with respect to [A]? • What is the order of reaction with respect to [B]? • If the rate is reported as the rate of formation of C, what would be the rate of disappearance of A?

  17. Integrated Rate Law Equation • First Order Reaction – aA  products • R = k[A] • First order in [A], 1st order overall • The Integrated Rate Equation is: • ln([A]o/[A]) = akt • Rearranged: ln[A]o – ln[A] = akt • ln[A] = -akt + ln[A]o • Y = mx + b

  18. Plot of ln[A] vs Time

  19. First Order Reaction

  20. Integrated Rate Equations • First Order Reactions: • Useful to approximate the time when half of the reactants are used up because the rate slows down considerably! • Rearrange the equation to solve for T • T=(1/ak) (ln[A]o/[A])

  21. Integrated Rate Equations • When [A] = ½ [A]o • T1/2 = (1/ak) (ln[A]o/1/2[A]o) • T1/2 = (1/ak) ln 2 = 0.693/ak • for 1st order reactions only, t1/2 depends only on the constant and does not change as the reaction progresses. • Practical Example: Half-life!

  22. Example 3 • Cyclopentane decomposes to propene in a 1st order reaction. K= 9.2 s-1 at 1000oC. A) calculate the half life at this temperature. B) How much of a 3.0g sample is left after 0.50 seconds? (assume grams are in the same proportionality as molarity)

  23. Second Order Reactions • R = k[A]2 • If a reaction is second order to a particular reactant and second order overall, the Integrated Rate Equation is: 1/[A] – 1/[A]o = akt At t1/2 [A] = 1/2[A]o 1/(1/2)[A]o – 1/[A]o = akt(1/2) 2/[A]o – 1/[A]o = akt(1/2) 1/[A]o = akt(1/2) and t1/2 = 1/ak[A]o • Concentration varies with each passing time period…concentration dependant!

  24. Second Order Reaction • The Half life of a second order reaction depends on the initial concentration at the beginning of THAT time period.

  25. Example 4 • CH3CHO (g)  CH4 (g)+ CO (g) • R = [CH3CHO]2 and k = 2.0 x 10-2 L/mole hr at 527oC a) What is the half life if 0.10 mol is injected into a 1.0L vessel? b) How many moles of CH3CHO remain after 200 hours?

  26. Second Order Reaction

  27. Zero Order Reactions • Zero Order Reaction = aA  products • R = k • Integtrated Rate Law = • [A] = [Ao] – akt • At t1/2…1/2[Ao] = [Ao] – akt • T1/2 = [Ao] / 2ak

  28. Zero Order Reaction

  29. Zero Order Reactions

  30. Graphical Review Can you pick out which is Zero, First, and Second Order?

  31. Investigating Factors Affecting Reaction Rate Collision Theory, Transition State Theory, Temperature, Catalysts and Activation Energy

  32. Collision Theory • Generally, any factor which increases the number of molecular or ionic collisions in solution will increase the rate of reaction • Stirring • Temperature • Concentration • Not Every Collision will guarantee a reaction! Orientation of the collision often affects the outcome!

  33. Transition State Theory • Transition State – short lived high energy complex between reactant and product.

  34. Exothermic Reaction

  35. Endothermic Reaction

  36. Transition State • Ea forward – Ea reverse = E rxn • Activation energy is generally kinetic energy, when equal or greater to the Ea, the reaction will proceed, if not, the reaction will not proceed.

  37. k increases with kelvin Theoretical Explanation Best model to explain is collision theory Kinetic molecular theory says the faster the particles move, the more they will collide Effect of Temperature

  38. Observed – reaction doesn’t increase with temperature as fast as the expected number of collisions Solution – Arrhenius – 1880 Not all collisions are effective, there must be a minimum amount of energy which must be present (Ea) Effect of Temperature

  39. Effect of Temperature # effective collisions = total collisions (e(-Ea/RT)) • Total collisions = Z • -Ea/RT = fraction of collisions with Ea or greater at a given temperature

  40. PROBLEM! • Number of observed collisions were less than calculated •  many of the collisions were ineffective due to orientation. • Fudge Factor P • P = steric factor (<1) = fraction of collisions with correct orientation

  41. Arrhenius’ Equation • k = Z p (e(-Ea/RT)) • Z and p combined into ultimate fudge factor A (frequency factor) • k = A (e(-Ea/RT)) • Alternate form: • lnk = -Ea/RT + lnA • Y = m x + b

  42. Alternate Form 2 • Ln(k1/k2) = -Ea/R (1/T1 – 1/T2) • Use to directly calculate the effect of temperature on the rate constant!

  43. Catalysts • Catalysts – substances added to a reaction which provide an alternative pathway to the reaction, thus lowering the activation energy for the reaction. • Heterogeneous catalysts – exist in the different state as the reactants. • Homogeneous catalysts – exists in the same state as the reactants.

  44. Lowers Ea by facilitating the breaking of bonds • Increases the rate of reaction

  45. Catalysts • Heterogeneous – works by contact (contact catalyst) • Adsorbtion – reactant comes in contact with the catalyst • Desorbtion – newly formed product separates from the catalyst. • Page 692 graphic!

  46. Catalysts • Enzymes – natural protein based catalysts • Work on same principles • Enzyme-substrate complex provides the alternative pathway to high energy biological processes.

  47. Reaction Mechanisms Writing Rate Laws for Multi-Step Reactions

More Related