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Unit 3: Chemical Kinetics

Unit 3: Chemical Kinetics. Reaction Rates Rate Laws Zero Order Reactions First Order Reactions Second Order Reactions Reaction Energy Diagrams Arrhenius Equation. Reaction Rates. Questions to consider: What makes “superglue” bond instantly while Elmer’s glue does not?

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Unit 3: Chemical Kinetics

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  1. Unit 3: Chemical Kinetics Reaction Rates Rate Laws Zero Order Reactions First Order Reactions Second Order Reactions Reaction Energy Diagrams Arrhenius Equation

  2. Reaction Rates • Questions to consider: • What makes “superglue” bond instantly while Elmer’s glue does not? • What factors determine how quickly food spoils? • Why do “glow sticks” last longer when stored in the freezer? • How do catalytic converters remove various pollutants from car exhaust?

  3. Reaction Rates • These types of questions can be answered using chemical kinetics. • The study of the speed or rate at which chemical reactions occur • The rate of a chemical reaction is affected by many factors, including: • Physical state of the reactants: • In order to react, molecules must come in contact with each other. • Molecules react faster under homogeneous conditions.

  4. Reaction Rates • The rate of a chemical reaction is affected by many factors, including (cont): • concentration of reactants • As concentration of reactants increases the rate of reaction generally increases • Greater chance of molecules colliding • reaction temperature • As temperature increases, reaction rate generally increases. • Molecules have more kinetic energy and collide more frequently with enough energy for a reaction to occur.

  5. Reaction Rates • The rate of a chemical reaction is affected by many factors (cont): • presence of acatalyst • a substance that increases the rate of a reaction without being consumed in the reaction • Enzymes • biological catalysts • proteins that increase the rate of biochemical reactions

  6. Reaction Rates • The speed of an object or event is the change that occurs in a given time interval. • Speed of a car = change in distance time interval =Dd Dt Remember, “change” refers to final value minus initial value.

  7. Reaction Rates • The speed or rate of a reaction can be defined in a similar manner. • Average reaction rate: • a positive value that describes the change in either the product or reactant concentration as a function of time • Common units: • M/s, M/min, or M/hr

  8. Reaction Rates • For a general reaction with 1:1 stoichiometry: A  B Avg rate = D [Product] = - D [Reactant] Dt Dt where [ ] is used to indicate the molarity of the material shown within the square brackets Negative sign needed to give a positive value

  9. Reaction Rates Consider the chemical reaction: AB t = 40. min 0.20 M A [B] = ? t = 20. min 0.50 M A [B] = ? Time = 0. 1.0 M A

  10. Reaction Rates For the reaction A  B, the following data can be used to determine the average reaction rate for a particular time interval: Time A B Rate (min)(M) (M)(M/min) 0.0 1.00 0.0 20.0 0.50 40.0 0.20

  11. Reaction Rates • Why are the reaction rates different for each of the time intervals in the previous example???

  12. Reaction Rates Example:Given the following data, what is the average rate of the following reaction over the time interval from 54.0 min to 215.0 min? CH3OH (aq) + HCl (aq) CH3Cl(aq) + H2O (l) Time (min) [HCl] (M) 0.0 1.85 54.0 1.58 107.0 1.36 215.0 1.02

  13. Reaction Rates Time (s) [C4H9OH] 0.0 0.1000 M 50.0 0.0095 100.0 0.0180 150.0 0.0259 200.0 0.0329 300.0 0.0451 Example: Calculate the average rate of reaction during the first 200.0 s of the reaction. C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)

  14. Reaction Rates • On the exam, you will be expected to find the average rate of reaction for a specific time interval when given the concentration (or number of moles) of either a reactant or a product as a function of time.

  15. A Momentary Diversion – A Return to Chem I Time (s) [C4H9Cl] [C4H9OH] 0.0 0.1000 M 0.0000 M 100.0 0.0820 200.0 0.0671 300.0 0.0549 C4H9Cl(aq) + H2O(l)C4H9OH(aq) + HCl(aq) How would you calculate the [C4H9OH] for this reaction at each time interval? (Note: This is NOT a kinetics question!!) You should expect a similar question on your exam!!!

  16. Reaction Rates • So far, all reactions have had a one-to-one stoichiometry. • What happens when the coefficients are not all 1? 2 A  3B

  17. Reaction Rates Consider the following reaction: 2 HI (g)  H2 (g) + I2 (g) Time mol mol mol (min) HI H2 I2 0.0 2.00 0.0 0.0 10.0 1.50 20.0 1.00 30.0 0.75 0.25 0.25 0.50 0.50 0.625 0.625

  18. Reaction Rates Calculate the change in HI and H2 as a function of time (this is NOT the rate) for the first 20.0 minutes of reaction: 2 HI (g)  H2 (g) + I2 (g) DHIDH2 Time mol mol mol Dt Dt (min) HI H2 I2 (mol/min) 0.0 2.00 0.0 0.0 10.0 1.50 0.25 0.25 20.0 1.00 0.50 0.50 30.0 0.75 0.625 0.625

  19. Reaction Rates • The average reaction rate must be numerically the same, regardless of whether you express it as the rate of appearance of a product or the rate of disappearance of a reactant. • HI disappears twice as fast as H2 appears. To make the rates equal: Rate = - 1D [HI] = D [H2] 2 Dt D t

  20. Reaction Rates • In general, for any reaction: a A + b B  c C + d D the rate of the reaction can be found by: Rate = - 1D[A] = - 1D[B] = 1D[C] = 1D[D] a Dt b Dt c Dt d Dt

  21. Reaction Rates • You should be able to use this equation on your exam to find the relationship between the rate of change of one reactant or product and any other reactant or product. • If you know the rate of change of one reactant or product, you should be able to calculate the rate of change of another reactant of product. • Simply using molar ratios is actually easier for this type of problem, however!

  22. Reaction Rates Example: Write a mathematical expression (equation) that shows how the rate of disappearance of N2O5 is related to the rate of appearance of NO2 in the following reaction? 2 N2O5 (g)  4 NO2 (g) + O2 (g)

  23. Reaction Rates Example: If the rate of decomposition of N2O5 in the previous example at a particular instant is 4.2 x 10-7M /s, what is the rate of appearance of NO2? 2 N2O5 (g)  4 NO2 (g) + O2 (g)

  24. Reaction Rates • Recall that the average reaction rate changes during the course of the reaction. • Until now, we have calculated average reaction rates. • The reaction rate at a particular time (not time interval) is called the instantaneous reaction rate.

  25. Reaction Rate • The instantaneous reaction rate is found by determining the slope of a line tangent to the curve at the particular time of interest. • Fortunately (for you), you won’t have to do this on the exam or HW!

  26. Rate Laws • The average reaction rate decreases with time. • The reaction slows down as the concentration of reactants decreases. CH3OH (aq) + HCl (aq) CH3Cl(aq) + H2O (l) Time (min) [HCl] (M) Avg. Rate (M /min) 0.0 1.85 54.0 1.58 0.0050 107.0 1.36 0.0042 215.0 1.02 0.0031

  27. Rate Laws • In general, the rate of any reaction depends on the concentration of reactants. • The way in which the reaction rate varies with the concentration of the reactants can be expressed mathematically using a rate law. • An equation that shows how the reaction rate depends on the concentration of the reactants

  28. Rate Laws • For a generalized chemical reaction: w A + x B  y C + z D the general form of the rate law is: Rate = k[A]m [B]n where k = rate constant m, n = reaction order

  29. Rate Laws • Rate Constant (k) • a proportionality constant that relates the concentration of reactants to the reaction rate • Reaction Order • the power to which the concentration of a reactant is raised in a rate law • Overall reaction order • The sum of all individual reaction orders

  30. Rate Laws • Rate laws must be determined experimentally. • Measure the instantaneous reaction rate at the start of the reaction (i.e. at t = 0) for various concentrations of reactants. • You CANNOT determine the value of “m” or “n” by looking at the coefficients in the balanced chemical equation for the overall reaction!

  31. Rate Laws • First Order Reaction • Overall reaction order = 1 • Rate = k[A] Expt [A] (M) Rate (M/s) 1 0.50 1.00 2 1.00 2.00 3 2.00 4.00 When [A] doubles, the rate doubles. When [A] increases by a factor of 4 the rate increases by a factor of 4.

  32. Rate Laws • Second Order Reaction • Overall reaction order = 2 • Rate = k[A]2 Expt [A] (M) Rate (M/s) 1 0.50 0.50 2 1.00 2.00 3 1.50 4.50 When [A] doubles, the rate increases by a factor of 22=4. When [A] increases by a factor of 3 the rate increases by a factor of 32 = 9.

  33. Rate Laws • Third Order Reaction • Overall reaction order = 3 • Rate = k[A]3 Expt [A] (M) Rate (M/s) 1 0.50 0.25 2 1.00 2.00 3 1.50 6.75 When [A] doubles, the rate increases by a factor of 23=8. When [A] increases by a factor of 3 the rate increases by a factor of 33 = 27.

  34. Rate Laws • Zero Order Reaction • Overall reaction order = 0 • Rate = k[A]0 so Rate = k Expt [A] (M) Rate (M/s) 1 0.50 2.00 2 1.00 2.00 3 1.50 2.00 When [A] doubles, the rate stays constant. When [A] increases by a factor of 3, the rate stays constant.

  35. Rate Laws REMEMBER • Rate laws must be determined experimentally. • Determine the instantaneous reaction rate at the start of the reaction (i.e. at t = 0) for various concentrations of reactants. • You CANNOT determine the value of “m” or “n” by looking at the coefficients in the balanced chemical equation!

  36. Rate Laws • To determine the rate law from experimental data, • identify two experiments in which the concentration of one reactant has been changed while the concentration of the other reactant(s) has been held constant • determine how the reaction rate changed in response to the change in the concentration of that reactant.

  37. Rate Laws • To determine the rate law from experimental data (cont) • Repeat this process using another set of data in which the concentration of the first reactant is held constant while the concentration of the other one is changed.

  38. Rate Laws Example: The initial reaction rate of the reaction A + B  C was measured for several different starting concentration of A and B. The following results were obtained. Determine the rate law for the reaction. Expt # [A] (M) [B] (M) Initial rate (M /s) 1 0.150 0.100 4.0 x 10-5 2 0.150 0.200 8.0 x 10-5 3 0.450 0.100 3.6 x 10-4

  39. Rate Laws

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