2. 1 Reaction Rate How Fast Does the Reaction Go

3. 2 Collision Theory l In order to react molecules and atoms must collide with each other. l They must hit each other hard enough to react. (Old bonds must break so new bonds can form.) l Reacting particles must collide with the correct spatial orientation. l Anything that increase these things will make the reaction faster.

4. 3 Energy Reaction coordinate Reactants Products

5. 4 Energy Reaction coordinate Reactants Products Activation Energy - Minimum energy to make the reaction happen

6. 5 Energy Reaction coordinate Reactants Products Activated Complex or Transition State

7. 6 Energy Reaction coordinate Reactants Products Overall energy change

8. 7

9. 8 Things that Effect Rate l Temperature - at higher temperature particles move faster which results in more and harder collisions, the result is faster reaction. l Concentration - more concentrated reacting particles are closer together and will collide more often, resulting in a faster reaction.

10. 9 Things that Effect Rate l Particle size – The only molecules that can collide are the ones at the surface. If the reactants is crushed there are more particles exposed, bigger surface area so the reaction is faster. l To get the particles the smallest possible, the reactants can be dissolved because in solution the reactants exist as molecules or ions. The solution speeds up reactions. (Getting two solids to react with each other is slow.)

11. 10 Things that Effect Rate l Catalysts- substances that speed up a reaction without being used up.(enzyme). l Speeds up reaction by giving the reaction a new path. l The new path has a lower activation energy. l More molecules have this energy. l The reaction goes faster. l Inhibitor- a substance that blocks a catalyst.

12. 11 Energy Reaction coordinate Reactants Products

13. 12 Pt surface HHHH HHHH l Hydrogen bonds to surface of metal. l Break H-H bonds Catalysts

14. 13 Pt surface HHHH Catalysts C HH C HH

15. 14 Pt surface HHHH Catalysts C HH C HH l The double bond breaks and bonds to the catalyst.

16. 15 Pt surface HHHH Catalysts C HH C HH l The hydrogen atoms bond with the carbon

17. 16 Pt surface H Catalysts C HH C HH HHH

18. 17 Reaction Mechanism l Elementary reaction- a reaction that happens in a single step. l Reaction mechanism is a description of how the reaction really happens. l It is a series of elementary reactions. l The product of an elementary reaction is an intermediate. l An intermediate is a product that immediately gets used in the next reaction.

19. 18 + This reaction takes place in three steps

20. 19 + EaEa First step is fast Low activation energy

21. 20 Second step is slow High activation energy + EaEa

22. 21 + EaEa Third step is fast Low activation energy

23. 22 Second step is rate determining

24. 23 Intermediates are present

25. 24 Activated Complexes or Transition States

26. 25 Mechanisms and rates l There is an activation energy for each elementary step. l Slowest step (rate determining) must have the highest activation energy.

27. 26

28. 27 Equilibrium l When I first put reactants together the forward reaction starts. l Since there are no products there is no reverse reaction. l As the forward reaction proceeds the reactants are used up so the forward reaction slows. l The products build up, and the reverse reaction speeds up.

29. 28 Equilibrium l Eventually you reach a point where the reverse reaction is going as fast as the forward reaction. l This is dynamic equilibrium. l The rate of the forward reaction is equal to the rate of the reverse reaction. l The concentration of products and reactants stays the same, but the reactions are still running.

30. 29 Equilibrium l Equilibrium position- how much product and reactant there are at equilibrium. l Shown with the double arrow. l Reactants are favored l Products are favored l Catalysts speed up both the forward and reverse reactions so don’t affect equilibrium position.

31. 30 Equilibrium Constant l aA + bB ↔ cC + dD l R f = k f [A] a [B] b l R r = k r [C] c [D] d l At equilibrium R f = R r l k f [A] a [B] b = k r [C] c [D] d k f = [C] c [D] d k r [A] a [B] b K eq = [C] c [D] d [A] a [B] b

32. 31 Measuring equilibrium l At equilibrium the concentrations of products and reactants are constant. l We can write a constant that will tell us where the equilibrium position is. l K eq equilibrium constant l K eq = [Products] coefficients [Reactants] coefficients l Square brackets [ ] means concentration in molarity (moles/liter)

33. 32 Writing Equilibrium Expressions l General equation aA + bB cC + dD l K eq = [C] c [D] d [A] a [B] b l Write the equilibrium expressions for the following reactions. l 3H 2 (g) + N 2 (g) 2NH 3 (g) l 2H 2 O(g) 2H 2 (g) + O 2 (g)

34. 33 Calculating Equilibrium l K eq is the equilibrium constant, it is only effected by temperature. l Calculate the equilibrium constant for the following reaction. 3H 2 (g) + N 2 (g) 2NH 3 (g) if at 25ºC there 0.15 mol of N 2, 0.25 mol of NH 3, and 0.10 mol of H 2 in a 2.0 L container.

35. 34 What it tells us l If K eq > 1 Products are favored l If K eq < 1 Reactants are favored

36. 35 LeChâtelier’s Principle Regaining Equilibrium

37. 36 LeChâtelier’s Principle l If something is changed in a system at equilibrium, the system will respond to relieve the stress. l Three types of stress are applied.

38. 37 Changing Concentration l If you add reactants (or increase their concentration). l The forward reaction will speed up. l More product will form. l Equilibrium “Shifts to the right” Reactants  products

39. 38 Changing Concentration l If you add products (or increase their concentration). l The reverse reaction will speed up. l More reactant will form. l Equilibrium “Shifts to the left” Reactants  products

40. 39 Changing Concentration l If you remove products (or decrease their concentration). l The forward reaction will speed up. l More product will form. l Equilibrium “Shifts to the right” Reactants  products

41. 40 Changing Concentration l If you remove reactants (or decrease their concentration). l The reverse reaction will speed up. l More reactant will form. l Equilibrium “Shifts to the left”. Reactants  products l Used to control how much yield you get from a chemical reaction.

42. 41 Changing Temperature l Reactions either require or release heat. l Endothermic reactions go faster at higher temperature. l Exothermic go faster at lower temperatures. l All reversible reactions will be exothermic one way and endothermic the other.

43. 42 Changing Temperature l As you raise the temperature the reaction proceeds in the endothermic direction. l As you lower the temperature the reaction proceeds in the exothermic direction. Reactants + heat  Products at high T Reactants + heat  Products at low T

44. 43 Changes in Pressure l As the pressure increases the reaction will shift in the direction of the least gases. At high pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g) At low pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g)

45. 44 E N D

46. 45 Thermodynamics Will a reaction happen?

47. 46 Energy l Substances tend react to achieve the lowest energy state. l Most chemical reactions are exothermic. l Doesn’t work for things like ice melting. l An ice cube must absorb heat to melt, but it melts anyway. Why?

48. 47 Entropy l The degree of randomness or disorder. lSlS l The first law of thermodynamics. The energy of the universe is constant. l The second law of thermodynamics. The entropy of the universe increases in any change. l Drop a box of marbles. l Watch your room for a week.

49. 48 Entropy Entropy of a solid Entropy of a liquid Entropy of a gas l A solid has an orderly arrangement. l A liquid has the molecules next to each other. l A gas has molecules moving all over the place.

50. 49 Entropy increases when... l Reactions of solids produce gases or liquids, or liquids produce gases. l A substance is divided into parts -so reactions with more reactants than products have an increase in entropy. l the temperature is raised -because the random motion of the molecules is increased. l a substance is dissolved.

51. 50 Entropy calculations l There are tables of standard entropy (pg 407 and the index). l Standard entropy is the entropy at 25ºC and 1 atm pressure. l Abbreviated Sº, measure in J/K. The change in entropy for a reaction is  Sº= Sº(Products)-Sº(Reactants). Calculate  Sº for this reaction CH 4 (g) + O 2 (g)  CO 2 (g) + H 2 O(g)

52. 51 Spontaneity Will the reaction happen, and how can we make it?

53. 52 Spontaneous reaction l Reactions that will happen. l Nonspontaneous reactions don’t. l Even if they do happen, we can’t say how fast. l Two factors influence. l Enthalpy (heat) and entropy(disorder).

54. 53 Two Factors l Exothermic reactions tend to be spontaneous. negative  H. l Reactions where the entropy of the products is greater than reactants tend to be spontaneous. Positive  S. A change with positive  S and negative  H is always spontaneous. A change with negative  S and positive  H is never spontaneous.

55. 54 Other Possibilities l Temperature affects entropy. l Higher temperature, higher entropy. l For an exothermic reaction with a decrease in entropy (like rusting). l Spontaneous at low temperature. l Nonspontaneous at high temperature. l Entropy driven.

56. 55 Other Possibilities l An endothermic reaction with an increase in entropy like melting ice. l Spontaneous at high temperature. l Nonspontaneous at low temperature. l Enthalpy driven.

57. 56 Gibbs Free Energy l The energy free to do work is the change in Gibbs free energy.  Gº =  Hº - T  Sº (T must be in Kelvin) l All spontaneous reactions release free energy. So  G <0 for a spontaneous reaction.

58. 57  G=  H-T  S HH SS Spontaneous? -+- At all Temperatures GG ++? At high temperatures, “entropy driven” --? At low temperatures, “enthalpy driven” +-+ Not at any temperature, Reverse is spontaneous

59. 58 Problems l Using the information on page 407 and pg 190 determine if the following changes are spontaneous at 25ºC. 2H 2 S(g) + O 2 (g)  2H 2 O(l) + S(rhombic) l At what temperature does it become spontaneous?

60. 59 2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2S l From Pg. 190 we find  H f ° for each component – H 2 S = -20.1 kJO 2 = 0 kJ – H 2 O = -285.8 kJS = 0 kJ l Then Products - Reactants l  H =[2 (-285.8) - 0] - [2 (-20.1) + 1(0)] = -531.4 kJ

61. 60 2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2 S l From Pg. 407 we find  S for each component – H 2 S = 205.6 J/K O 2 = 205.0 J/K – H 2 O = 69.94 J/K S = 31.9 J/K l Then Products - Reactants l  S=[2 (69.94) - 2(31.9) ] - [2 (205.6) + 205] = -412.5 J/K

62. 61 2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2 S l  G =  H - T  S l  G = -531.4 kJ - 298K (-412.5 J/K) l  G = -531.4 kJ - -123000 J l  G = -531.4 kJ - -123 kJ l  G = -408.4 kJ l Spontaneous l Exergonic- it releases free energy. l At what temperature does it become spontaneous?

63. 62 Spontaneous l It becomes spontaneous when  G = 0 l That’s where it changes from positive to negative. l Using 0 =  H - T  S and solving for T l 0 -  H = - T  S l -  H = -T  S l T =  H =  S = 1290 K -531.4 kJ -412.5 J/K = -531400 kJ -412.5 J/K

64. 63 There’s Another Way l There are tables of standard free energies of formation compounds.(pg 414)  Gº f is the free energy change in making a compound from its elements at 25º C and 1 atm. for an element  Gº f = 0 l Look them up.  Gº=  Gº f (products) -  Gº f (reactants) l Check the last problems.

65. 64 2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2S l From Pg. 414 we find  H f ° for each component – H 2 S = -33.02 kJO 2 = 0 kJ – H 2 O = -237.2 kJS = 0 kJ l Then Products - Reactants l  H =[2 (-237.2) - 2(0)] - [2 (-33.02) + 1(0)] = -408.4 kJ

66. 65 Reversible Reactions Reactions are spontaneous if  G is negative. If  G is positive the reaction happens in the opposite direction. 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy 2H 2 O(g) + energy   H 2 (g) + O 2 (g) 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy