How rapidly reactions proceed - rate of reaction

1. Chemical Kinetics How rapidly reactions proceed - rate of reaction Details of process from reactants to products - mechanism Thermodynamics determines the direction in which reactions proceed spontaneously and the conditions at equilibrium, but not the rate at which equilibrium is reached. For a complete picture of a chemical reaction need both the thermodynamics and kinetics of a reaction.

3. N2(g) + 3H2(g)  2NH3(g) At 298 K DGo = -33.0 kJ DHo = -92 kJ; exothermic reaction K = 6.0 x 105 ; thermodynamically favored at 298 K, Rate is slow at 298K Commercial production of NH3 is carried out at temperatures of 800 to 900 K, because the rate is faster even though K is smaller.

4. Thermodynamic functions are state functions (DG, DH, DE) Thermodynamic functions do not depend on the mechanism of the reaction. The rate of the reaction is very dependent on the path of the process or path between reactants and products. Kinetics reveals information on the rate of the reaction and the mechanism (path) of the reaction.

5. Thermodynamics vs Kinetics A + B  C + D K1 A + B  E + F K2 If K1 > K2 - products C & D are thermodynamically favored over E & F. If products observed are C & D: reaction is thermodynamically controlled If products observed are E & F: reaction is kinetically controlled

6. Zn2+(aq) + S2-(aq)  ZnS(s) K = 1/Ksp = 2.2 x 1023 Fe2+(aq) + S2-(aq)  FeS(s) K = 1/Ksp = 2.7 x 1018 On addition of S2- to an aqueous solution containing both Zn2+ and Fe2+, ZnS precipitates first - reaction is thermodynamically controlled.

7. (1) 2NO(g) + O2(g)  2NO2(g) (2) 2CO(g) + O2(g)  2CO2(g) Both have large values of K; both are thermodynamically favored in the forward direction Reaction (1) is very fast; reaction (2) slow Reactions are kinetically controlled Result is brown color of air due NO2 and buildup of CO in the air

8. Rates of Reactions Zn(s) + 2 H+(aq)  Zn2+(aq) + H2(g) Rate of a reaction: change in concentration per unit time

9. change in concentration change in time [NO]final - [NO]initial average reaction rate = tfinal - tinitial average reaction rate = Units of rate: concentration / time NO2(g) + CO(g)  NO(g) + CO2 (g)

10. NO2(g) + CO(g)  NO(g) + CO2 (g) Time (s) [NO] mol L-1 0 0 50 0.0160 100 0.0240 150 0.0288 200 0.0320 Average rate 1st 50 seconds = 3.2 x 10-4 mol L-1 Average rate 2nd 50 seconds = 1.6 x 10-4 mol L-1 Average rate 3rd 50 seconds = 9.6 x 10-5 mol L-1

11. Instantaneous Rate - rate at a particular moment in time

12. - - - - rate = rate = = = = = = = d[NO2] D[NO2] d t Dt d[A] - d[C] d[D] d[B] - = = rate = = d[CO2] d[CO] d[NO] dt D[CO] D[CO2] D[NO] dt dt dt d t d t d t Dt Dt Dt 1 1 1 1 x y a b NO2(g) + CO(g)  NO(g) + CO2 (g) For an infinitesimally small changes, the instantaneous rate For a general reaction: aA + bB  xC + yD

13. initial rate (t = 0)

14. Rate (at 5 weeks) = -6.3 x 10-3 Rate (at 10 weeks) = -2.6 x 10-3

15. Factors affecting rates of reactions a) Nature of reactants 2NO(g) + O2(g)  2NO2(g) fast 2CO(g) + O2(g)  2CO2(g) slow b) Concentration of reactants: reactions proceed by collisions between reactants c) Temperature: In general, as T increases, rate increases d) Catalyst: increases rate of reaction e) Surface: larger surface area increases rate of reaction f) Nature of solvent

16. Effect of concentration Effect of temperature Effect of surface area

17. Rate Laws and Rate Constant Initial reaction rate: instantaneous rate of change in concentration of a species at the instant the reaction begins. (products present later in the reaction may affect the rate) 2 N2O5 (g)  4 NO2 (g) + O2 (g) Perform a series of experiments with different initial concentrations of N2O5 (g). For each, monitor the concentration of N2O5 (g) versus time Determine the initial rate of reaction at t = 0

18. 2 N2O5 (g)  4 NO2 (g) + O2 (g) For this reaction, experiments indicate: rate  initial concentration of N2O5 (g) rate = k x initial concentration of N2O5 (g)

19. d[NO2] dt Rate = - d[N2O5] / dt = k [N2O5] Rate law k = 5.2 x 10-3 s-1 k is the specific rate constant Units of k depends on the rate law Rate laws are determined experimentally 2 NO2(g)  2 NO(g) + O2 (g) - rate = = k [NO2]2Rate Law k = 0.54 (mol NO2)-1 s-1

20. 2 NO2(g)  2 NO(g) + O2 (g) Rate = k [NO2]2

21. d[A] - d[C] d[D] d[B] - = = rate = = dt dt dt dt 1 1 1 1 d a c b For a general reaction: aA + bB  cC + dD rate = k [A]m [B]n Rate Law For a reaction k has a specific value; k for the reaction changes with temperature Note: m need not equal a; n need not equal b

23. Order of a Reaction rate = k [A]m [B]n Reaction is mth order in A and nth order in B Overall reaction order = m + n The reaction order is determined by the experimentally determined rate law N2O5(g)  N2O4(g) + 1/2 O2(g) Rate = k [N2O5] First order reaction For a 1st order reaction, units of k: time-1

24. C2H6(g)  2 CH3(g) rate = k [C2H6]2second order reaction 2NO2(g)  2NO(g) + O2 (g) Rate = k [NO2]2 second order reaction For 2nd order reactions, units of k: concentration-1 time-1

25. S2O82-(aq) + 3 I-(aq)  2 SO42-(aq) + I3-(aq) (S2O82- persulfate) Rate of disappearance of S2O82- = k [S2O82-(aq)] [I-(aq)] Overall reaction order = 1 + 1 = 2 1st order in S2O82-(aq) and 1st order in I-(aq) If the above reaction was carried out at a high concentration S2O82-(aq), and experimentally: Rate of disappearance of S2O82- = k’ [I-(aq)] pseudo first-order reaction

26. 2 NH3(g)  N2 (g) + 3 H2(g) Rate of disappearance of NH3 = k zero order reaction