1. Chapter 14 Kinetics

2. Defined:

3. Factors that affect Rate 1. 2. 3. 4.

4. Review… In any equation –Reactants  products –N 2(g) + 3H 2(g)  2NH 3(g) Equilibrium –

5. RATE Speed of the Reaction – Measure either – –

6. Rate Equation Rate =

7. Rate Equation Rate =

9. Rate Equation Average Rate =

10. Instantaneous Rate

11. Rate and Concentration Generally, rate decreases as time passes. –Why?

12. Study rate by looking at NH 4(aq) + + NO 2 - (aq)  N 2(g) + 2H 2 O (l) See table 14.3 pg 516

13. Rate Law Rate = Must determine the value of k, the rate constant Choose one set of data… Look at the units

14. Practice Using previous rate law, what is the rate when the concentration of each reactant is 0.334 M? K= 2.7 x 10 -4 M -1 s -1

15. Reaction Order If rate law = k[react 1] m [react 2] n The exponents are

16. Units of rate constants Depend on the Rate units are Rate constant units must allow that

17. For example Second order reaction Units of rate =

18. Determining Rate Law Must be determined experimentally Compare rates as concentration of reactants is changed

19. Condition 1 Change concentration =

20. Condition 2 Change in concentration = Double concentration = Triple concentration =

21. Condition 3 Double concentration = – Triple concentration =

22. Rate aA + bB  cC + d D

23. Dependence of Rate on Concentration

24. Conditions for Rxn Orientation of molecules Energy, specifically, Kinetic Energy

25. Reactive Collisions Both previous conditions met – – KE is easy to change – –

26. Activation Energy, E a Minimum energy needed for reactive collision – Increase T means E a is also energy required to make

27. Activated complex When reacting molecules ‘stick’ together before completing reaction They must have – – Intermediate species

28. Reaction completes When activated complex breaks apart New molecules result from this process No activated complex left in system

29. Exothermic Reactions Stored chemical energy Disorder Products at

30. Endothermic Reactions Energy is –Increase of Disorder Products at

31. Direction of Reaction Forward Reverse One is favorable Tendency to increase Tendency to decrease

32. Heat of reaction  H rxn Enthalpy of reaction Measure of the energy change Formula –  H rxn = E f – E r

33. Exo or Endothermic? If  H > 0 kJ –Reaction is –+ sign tells you If  H < 0 kJ –Reaction is –- sign tells you

34. Practice What is the  H rxn if the Energy of the forward reaction is 65 kJ and that of the reverse is 32? Is this reaction exo or endothermic?

35. Catalysts Something that increases the rate of reaction Not consumed Provides a ‘Helps’ in the formation of the activated complex

36. Formation of Acid Rain Coal and sulfur Produce SO 2 in cars and energy plants Reacts with water in air Needs NO as

37. Catalysts Decrease E a Do not change the Only change the Provides a

38. Enzymes CO 2 + H 2 O  H 2 CO 3 Carbonic anhydrase –Enzyme (-ase is the key) –Facilitates conversion –Changes the rate by a factor of 3.5 x 10 6

39. Enzyme/Substrate Complex Decreases Therefore increases High E a = Low E a =