Successfully reported this slideshow.
Upcoming SlideShare
×

# Chemical kinetics

6,135 views

Published on

This file have little discription about Chemical Kinetics. Please guid me how to make it better.

Published in: Education
• Full Name
Comment goes here.

Are you sure you want to Yes No
• Be the first to comment

### Chemical kinetics

1. 1. Chemical Kinetics Chapter 13
2. 2. Chemical KineticsThermodynamics – does a reaction take place?Kinetics – how fast does a reaction proceed?Reaction rate is the change in the concentration of areactant or a product with time (M/s). A B ∆[A] ∆[A] = change in concentration of A over rate = - ∆t time period ∆t ∆[B] ∆[B] = change in concentration of B over rate = ∆t time period ∆t Because [A] decreases with time, ∆[A] is negative. 13.1
3. 3. A B time ∆[A]rate = - ∆t ∆[B]rate = ∆t 13.1
4. 4. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g) time Br2 (aq) 393 nm393 nm Detector light ∆[Br2] α ∆Absorption 13.1
5. 5. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g) slope of tangent slope of tangent slope of tangent ∆[Br2] [Br2]final – [Br2]initial average rate = - =- ∆t tfinal - tinitialinstantaneous rate = rate for specific instance in time 13.1
6. 6. rate α [Br2] rate = k [Br2] ratek= = rate constant [Br2] = 3.50 x 10-3 s-1 13.1
7. 7. 2H2O2 (aq) 2H2O (l) + O2 (g) PV = nRT n P= RT = [O2]RT V 1 [O2] = P RT ∆[O2] 1 ∆P rate = = ∆t RT ∆tmeasure ∆P over time 13.1
8. 8. 2H2O2 (aq) 2H2O (l) + O2 (g) 13.1
9. 9. Reaction Rates and Stoichiometry 2A BTwo moles of A disappear for each mole of B that is formed. 1 ∆[A] ∆[B] rate = - rate = 2 ∆t ∆t aA + bB cC + dD 1 ∆[A] 1 ∆[B] 1 ∆[C] 1 ∆[D] rate = - =- = = a ∆t b ∆t c ∆t d ∆t 13.1
10. 10. Write the rate expression for the following reaction:CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g) ∆[CH4] 1 ∆[O2] ∆[CO2] 1 ∆[H2O]rate = - =- = = ∆t 2 ∆t ∆t 2 ∆t 13.1
11. 11. The Rate LawThe rate law expresses the relationship of the rate of a reactionto the rate constant and the concentrations of the reactantsraised to some powers. aA + bB cC + dD Rate = k [A]x[B]y reaction is xth order in A reaction is yth order in B reaction is (x +y)th order overall 13.2
12. 12. F2 (g) + 2ClO2 (g) 2FClO2 (g) rate = k [F2]x[ClO2]yDouble [F2] with [ClO2] constantRate doublesx=1 rate = k [F2][ClO2]Quadruple [ClO2] with [F2] constantRate quadruplesy=1 13.2
13. 13. Rate Laws• Rate laws are always determined experimentally.• Reaction order is always defined in terms of reactant (not product) concentrations.• The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation. F2 (g) + 2ClO2 (g) 2FClO2 (g) rate = k [F2][ClO2] 1 13.2
14. 14. Determine the rate law and calculate the rate constant for the following reaction from the following data: S2O82- (aq) + 3I- (aq) 2SO42- (aq) + I3- (aq) [S2O82-] Initial RateExperiment [I-] (M/s) rate = k [S2O82-]x[I-]y 1 0.08 0.034 2.2 x 10-4 y=1 2 0.08 0.017 1.1 x 10-4 x=1 3 0.16 0.017 2.2 x 10-4 rate = k [S2O82-][I-]Double [I-], rate doubles (experiment 1 & 2)Double [S2O82-], rate doubles (experiment 2 & 3) rate 2.2 x 10-4 M/s k= = = 0.08/M•s [S2O8 ][I ] 2- - (0.08 M)(0.034 M) 13.2
15. 15. First-Order Reactions ∆[A] A product rate = - rate = k [A] ∆t ∆[A] rate M/s - = k [A]k= = = 1/s or s-1 ∆t [A] M [A] is the concentration of A at any time t [A]0 is the concentration of A at time t=0 [A] = [A]0exp(-kt) ln[A] = ln[A]0 - kt 13.3
16. 16. The reaction 2A B is first order in A with a rate constant of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease from 0.88 M to 0.14 M ? [A]0 = 0.88 Mln[A] = ln[A]0 - kt [A] = 0.14 M kt = ln[A]0 – ln[A] [A]0 0.88 M ln ln ln[A]0 – ln[A] [A] 0.14 M t= = = = 66 s k k 2.8 x 10 s-2 -1 13.3
17. 17. First-Order ReactionsThe half-life, t½, is the time required for the concentration of areactant to decrease to half of its initial concentration. t½ = t when [A] = [A]0/2 [A]0 ln [A]0/2 ln2 0.693 t½ = = = k k k What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1? t½ = ln2 = 0.693 = 1200 s = 20 minutes k 5.7 x 10 s -4 -1How do you know decomposition is first order? units of k (s-1) 13.3
18. 18. First-order reaction A product # ofhalf-lives [A] = [A]0/n 1 2 2 4 3 8 4 16 13.3
19. 19. Second-Order Reactions ∆[A] A product rate = - rate = k [A]2 ∆t rate M/s ∆[A]k= = 2 = 1/M•s - = k [A]2 [A] 2 M ∆t 1 1 [A] is the concentration of A at any time t = + kt [A] [A]0 [A]0 is the concentration of A at time t=0 t½ = t when [A] = [A]0/2 1 t½ = k[A]0 13.3
20. 20. Zero-Order Reactions ∆[A] A product rate = - rate = k [A]0 = k ∆t rate ∆[A]k= = M/s - =k [A] 0 ∆t [A] is the concentration of A at any time t [A] = [A]0 - kt [A]0 is the concentration of A at time t=0 t½ = t when [A] = [A]0/2 [A]0 t½ = 2k 13.3
21. 21. Summary of the Kinetics of Zero-Order, First-Order and Second-Order Reactions Concentration-Time Order Rate Law Equation Half-Life [A]0 0 rate = k [A] = [A]0 - kt t½ = 2k 1 rate = k [A] ln[A] = ln[A]0 - kt t½ = ln2 k 1 1 1 2 rate = k [A] 2 = + kt t½ = [A] [A]0 k[A]0 13.3
22. 22. A+B C+D Exothermic Reaction Endothermic ReactionThe activation energy (Ea) is the minimum amount ofenergy required to initiate a chemical reaction. 13.4
23. 23. Temperature Dependence of the Rate Constant k = A • exp( -Ea/RT ) (Arrhenius equation) Ea is the activation energy (J/mol) R is the gas constant (8.314 J/K•mol) T is the absolute temperature A is the frequency factor Ea 1 lnk = - + lnA R T 13.4
24. 24. Ea 1lnk = - + lnA R T 13.4
25. 25. Reaction MechanismsThe overall progress of a chemical reaction can be representedat the molecular level by a series of simple elementary stepsor elementary reactions.The sequence of elementary steps that leads to productformation is the reaction mechanism. 2NO (g) + O2 (g) 2NO2 (g) N2O2 is detected during the reaction! Elementary step: NO + NO N 2 O2 + Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 13.5
26. 26. Intermediates are species that appear in a reactionmechanism but not in the overall balanced equation.An intermediate is always formed in an early elementary stepand consumed in a later elementary step. Elementary step: NO + NO N 2 O2+ Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2The molecularity of a reaction is the number of molecules reacting in an elementary step.• Unimolecular reaction – elementary step with 1 molecule• Bimolecular reaction – elementary step with 2 molecules• Termolecular reaction – elementary step with 3 molecules 13.5
27. 27. Rate Laws and Elementary StepsUnimolecular reaction A products rate = k [A]Bimolecular reaction A+B products rate = k [A][B]Bimolecular reaction A+A products rate = k [A]2Writing plausible reaction mechanisms:• The sum of the elementary steps must give the overall balanced equation for the reaction.• The rate-determining step should predict the same rate law that is determined experimentally.The rate-determining step is the slowest step in thesequence of steps leading to product formation. 13.5
28. 28. The experimental rate law for the reaction between NO 2 and CO to produce NO and CO2 is rate = k[NO2]2. The reaction is believed to occur via two steps: Step 1: NO2 + NO2 NO + NO3 Step 2: NO3 + CO NO2 + CO2What is the equation for the overall reaction? NO2+ CO NO + CO2What is the intermediate? NO3What can you say about the relative rates of steps 1 and 2? rate = k[NO2]2 is the rate law for step 1 so step 1 must be slower than step 2 13.5
29. 29. A catalyst is a substance that increases the rate of achemical reaction without itself being consumed. k = A • exp( -Ea/RT ) Ea k uncatalyzed catalyzed ratecatalyzed > rateuncatalyzed Ea‘ < Ea 13.6
30. 30. In heterogeneous catalysis, the reactants and the catalystsare in different phases. • Haber synthesis of ammonia • Ostwald process for the production of nitric acid • Catalytic convertersIn homogeneous catalysis, the reactants and the catalystsare dispersed in a single phase, usually liquid. • Acid catalysis • Base catalysis 13.6
31. 31. Haber Process Fe/Al2O3/K2ON2 (g) + 3H2 (g) 2NH 3 (g) catalyst 13.6
32. 32. Ostwald Process Pt catalyst4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g) 2NO (g) + O2 (g) 2NO2 (g)2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq) Hot Pt wire Pt-Rh catalysts used over NH3 solution in Ostwald process 13.6
33. 33. Catalytic Converters catalyticCO + Unburned Hydrocarbons + O2 converter CO2 + H2O catalytic 2NO + 2NO2 converter 2N2 + 3O2 13.6
34. 34. Enzyme Catalysis 13.6
35. 35. enzymeuncatalyzed catalyzed 13.6