Chapter 14 – Chemical Kinetics

1. How fast a chemical reaction occurs Only need to consider the forward reaction Factors that affect rate Concentration of reactants Temperature Catalysts Surface Area Chapter 14 – Chemical Kinetics

2. B. Reaction Rates • Rate is determined in the lab by experiment • Rate determined by measuring (-) disappearance of reactants (+) appearance of products

3. Example 1 – Rates of… A + B  C + D Rate of appearance of C Note the POSITIVE sign!! Rate of disappearance of A Note the NEGATIVE sign!!

4. C4H9Cl + H2O  C4H9OH + HCl Data Table 14.2 – Disappearance of C4H9Cl = 1.64 x 10-4 M/s

5. Example 2 – Rates with Coefficients aA + bB  cC + dD Note that the coefficient becomes a reciprocal value for rate comparison

6. Example 3 – Rate Comparison 2 N2O5  4 NO2 + O2 Given: RN2O5 = 4.2 x 10-7 M/s Calculate the rate of appearance of NO2

7. Example 3 – Rate Comparison 2 N2O5  4 NO2 + O2 Given: RN2O5 = 4.2 x 10-7 M/s The rate of O2 appearance is ½ the rate of N2O5 disappearance

8. Rate Law Expression R = k [reactant]m • R = rate law expression • k = rate constant units are M-1s-1 Note: k depends upon temperature and nature of reaction • m = order of reaction • m=0  rate is independent of [ ]0 • m=1  rate is directly related to [ ]1 • m=2  rate is directly related to [ ]2

9. aA + bB  products • R = k [A]m[B]n m = order with respect to A n = order with respect to B Overall order of reaction is = m + n Note: order of reaction must be determined experimentally in the lab and cannot be simply concluded from the equation coefficients!!!!

10. 2 N2O5 4 NO2 + O2 R = k[N2O5] • CHCl3 + Cl2  CCl4 + HCl R = k [CHCl3][Cl2] • H2 + I2  2 HI R = k [H2] [I2]

11. Method of Initial Rates A + B  C

12. First Order Reactions • Using Calculus… ln [A]t – ln [A]0 = -kt or ln [A]t /[A]0 = -kt [A]0=original conc [A]t=conc @ time, t k = rate constant t = time

13. [A] t ln [A] t Graphing First Order Reactions ln [A]t = -k t + ln [A]0 y = m x + b This is NOT a linear plot…. Scientists like linear plots

14. Example – 1st Order • The decomposition of an insecticide in H2O is first order with a rate constant of 1.45 yr -1. On June 1st, a quantity of 5.0x10-7 g/cm3 washed into a lake. insect  product R = k [insect] • What is the concentration on June 1st next year? ans. [insect]t=1yr = 1.17x10-7 g/cm3 b) How long will it take for the [insect] to drop to 3.0x10-7 g/cm3? ans. t = 0.35 years = 4 months

15. 1st Order Reactions, Half-Life The time that it takes for Original concentration to Drop to ½ of its original concentration.

16. 1 [A] Second Order Reactions y = m x + b • Using Calculus… [A]0=original conc [A]t=conc @ time, t k = rate constant t = time slope=k t

17. 2nd Order Reactions, Half-Life t1/2= 1 k[A]0

18. 1 [A] ln [A] t To Determine Order You Must Graph the Data y = m x + b y = m x + b 1st Order 2nd Order slope=k t

19. Activation Energy, Ea • Molecules must collide to react • Not all collisions result in a reaction • The higher the collision frequency, the faster the reaction rate a. increase temperature b. increase pressure or decrease volume (for gas only) c. catalyst d. increase [conc]

20. Activation Energy, Ea • Activation energy, Ea – the minimum energy needed to start a reaction • Activated complex – intermediate product forming before the reaction is completed

21. Activated Complex A* The bigger Ea, the slower the rate Ea A Energy E B Reaction progress For A  B exothermic E (-) For B  A endothermic E (+) + Ea

22. Arrhenius Equation – Rate and Temperature k=rate const A=frequency Ea=Activation energy R=gas const 8.31 J/mol K T=Temperature (Kelvin)

23. Solving Arrhenius for Two Temperatures

24. Yintercept= ln A ln k Slope = - Ea R 1/ T Graphing Arrhenius Note: to obtain Ea, you must multiply slope by the gas constant