2. Ch 17 Reaction Rates & Ch 18 Equilibrium Part 1. How Fast Is the Reaction?

3. Collision Theory l In order to react, molecules and atoms must touch each other. l And they must hit each other hard enough to react. l Anything that increases these things will make the reaction faster. l There is a certain amount of energy needed to start the reaction. This energy is called the activation energy.

4. Collision Theory l Collision theory - atoms, ions, and molecules must collide in order to react. l For any reaction to occur, the particles must come in contact with each other. l Only a small amount of collisions produce reactions.

5. Orientation l Even though particles collide, the particles must have the correct orientation to react. l When correct orientation does occur a short lived intermediate substance is formed is called an activated complex l activated complex - a temporary, unstable arrangement of atoms that may form products or may break apart to re-form reactants.

6. Activation Energy l Even when particles collide and with correct orientation, the reaction still might not occur l An activated complex will not form if there is insufficient energy. l Activation energy (E a ) - the minimum amount of energy that reacting particles must have to form the activated complex and lead to a reaction.

7. Correct Orientation and Energy

8. Energy Reaction coordinate Reactants Products

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

10. Energy Reaction coordinate Reactants Products Activated Complex or Transition State

11. Energy Reaction coordinate Reactants Products Overall energy change

12. Reaction Pathway l Shows the change in energy during a chemical reaction

13. Exothermic Reaction l reaction that releases energy l products have lower PE than reactants 2H 2 (l) + O 2 (l)  2H 2 O(g) + energy energy released

14. Endothermic Reaction l reaction that absorbs energy l reactants have lower PE than products 2Al 2 O 3 + energy  4Al + 3O 2 energy absorbed

15. Factors Affecting Reaction Rates Movie

16. Things that Affect Reaction Rate l Temperature is important. l Higher temperature means faster particles. l Faster means more and harder collisions. l This leads to faster reactions. l Concentration is important. l More concentrated the reactants, the closer together the molecules are. l Closer molecules means colliding more often. l Again leading to faster reactions.

17. Things that Affect Reaction Rate l Particle size can be a factor. l Molecules can only collide at the surface. l Smaller particles, bigger the surface area. l Smaller particles means faster reactions. l Smallest possible molecules or ions result from dissolving compounds. Dissolving speeds up reactions. l Getting two solids to react with each other is slow.

18. Things that Affect 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.

19. Things That Affect Rate 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 slows or even blocks a reaction.

20. The Nature of Reactants l The reactive nature of reactants can determine the rate of reactions l Some substances react more readily than others such as the alkali metals. l For example, both calcium and sodium are reactive metals, however when placed in water, sodium reacts much more readily and produces much more energy than calcium.

21. Concentration l Reactions speed up when the concentrations of reacting particles are increased l The reaction speeds up because the amount of collisions between to reacting particles increases.

22. Surface Area l When the surface area increase the reaction rate increases l Increasing the surface area allows more particles to collide with one another, thus increasing the reaction rate.

23. Temperature l Usually increasing the temperature of a reaction generally increases the reaction rate. l This occurs because when temperature is increased, the particles move more quickly thus more collisions occur. l Also, when temperature is increased the energy of the collisions is greater.

24. Energy Reaction coordinate Reactants Products

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

26. Pt surface HHHH Catalysts C HH C HH

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

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

29. Pt surface H Catalysts C HH C HH HHH

30. Catalytic Converter

31. Inhibitors l Inhibitors slow down reaction rates and some stop reactions from occurring all together. l The food industry uses inhibitors to keep foods fresher, longer l Example is the chemical that is put on apples to keep apples from browning.

32. Use an Inhibitor to Slow Down the Reaction

33. Expressing Reaction Rates l Chemical reaction occur at certain rates. l The reaction between vinegar and baking soda occurs relatively fast where as the reactions that occur in the formation of fossil fuels occurs much slower l reaction rate - the change in concentration of a reactant or a product per unit time, expressed as M/s.

34. Calculating Reaction Rates l Average reaction rate = Δ quantity / Δ time l Δ quantity = final molar - initial molar l Δ time = final time - initial time l The value is never negative, so take the absolute value

35. Sample Question l In a reaction between butyl chloride (C 4 H 9 Cl) and water the initial concentration of butyl chloride was 0.220M at time 0.00 s and the concentration at time 4.00 s was 0.100M. Calculate the average reaction rate over this time period. l Known: t 1 = 0.00 s C 4 H 9 Cl at t 1 = 0.220M t 2 = 4.00 s C 4 H 9 Cl at t 2 = 0.100M l Unknown: average reaction rate = ? M/s

36. Answer l Average reaction rate = Δ quantity / Δ time l = 0.100 M - 0.220 M / 4.00 s - 0.00 s l = |-0.120 M/4.00 s| note: never negative l = 0.0300 M/s

37. Reaction Rate Laws l The equation that expresses the mathematical relationship between the rate of a chemical reaction and the concentration of reactants is called the rate law. l Rate = k[A] m [B] n l Rates are determined experimentally l k is the specific rate constant and varies with temperature.

38. Reaction Rate Laws l m and n are reaction orders and define how the rate is affected by the concentration (molarity) of reactants A and B. l The overall reaction order is equal to the sum of the individual reactant reaction orders. l If doubling the concentration doubles the reaction, then the order is first order or one, ie [A] 1 l If doubling the concentration quadruples the reaction, then the order is second order or two, [A] 2

39. Overall Reaction Order l We can use the method of initial rates to determine the overall reaction order. l Examine the experimental data below to determine the reaction order.

40. Overall Reaction Order l Recall the general rate law is l Rate = k[A] m [B] n l Looking at trial 1 and 2 we see doubling A doubles the initial rate. Therefore the reaction is 1 st order with A concentration (m=1). l Looking at trials 2 and 3 we see doubling B quadruples the initial rate. Therefore the reaction is 2 nd order with B (n=2). l The overall rate is 3 rd order which is the sum of m and n.

41. Your Turn l Determine the general rate law formula and overall reaction order based on the experimental data below.

42. 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.

43. Reaction Mechanism l An intermediate is a product that immediately gets used in the next reaction.

44. Reaction Mechanisms l Many chemical reactions consist of a sequence of two or more reactions l Such is evident in the earth’s stratosphere where 2O 3 3O 2 l This is the overall reaction after three steps occur which is started when intense UV radiation from the sun liberates chlorine atoms from certain compounds.

45. Reaction l Elementary step: Cl + O 3 O 2 + ClO l Elementary Step: O 3 O 2 + O l Elementary Step: ClO + O Cl + O 2 l Complex reaction: 2O 3 3O 2

46. Reaction Mechanisms l The destruction of ozone on the previous slide is a complex reaction l The complete sequence of elementary steps that make up a complex reaction is known as a reaction mechanism. l In a reaction mechanism some substances are produce and remain while others known as intermediates are produced by one reaction but consumed in a subsequent reaction.

47. + This reaction takes place in three steps

48. + EaEa First step is fast Low activation energy

49. Second step is slow High activation energy + EaEa

50. + EaEa Third step is fast Low activation energy

51. Second step is rate determining

52. Intermediates are present

53. Activated Complexes or Transition States

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

55. Thermodynamics Ch.16.5 Will a reaction happen?

56. Energy l Substances tend to 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. Why? l Thermite Reaction moviemovie

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

58. Entropy l Defined in terms of probability. l Substances take the arrangement that is most likely. l The most likely is the most random. l Calculate the number of arrangements for a system.

59. l 2 possible arrangements l 50 % chance of finding the left empty

60. l 4 possible arrangements l 25% chance of finding the left empty l 50 % chance of them being evenly dispersed

61. l 16 possible arrangements l 6.25% chance of finding the left empty l 37.5 % chance of them being evenly dispersed

62. Gases l Gases completely fill their chamber because there are many more ways to do that than to leave half empty. l S solid <S liquid <<S gas l there are many more ways for the molecules to be arranged as a liquid than a solid. l Gases have a huge number of positions possible.

63. 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.

64. Entropy increases when... l reactions of solids produce gases or liquids, or liquids produce gases. l a substance is divided into parts -reactions with less reactants than products increase in entropy. l the temperature is raised -the random motion of the molecules is increased. l a substance is dissolved. –increase in # of pieces l There are tables of standard entropy but our book doesn’t have one. I will give them to you.

65. Entropy Calculations l Standard entropy is the entropy at 25ºC and 1 atm pressure. l Abbreviated Sº, measure in J/K. l The change in entropy for a reaction is  Sº= Sº(Products) - Sº(Reactants) l Determine if the reaction is positive or negative entropy. If entropy increases then ΔS is positive. l A solid changing to a gas is an increase in entropy. Making more products or dissolving a solid is a positive change as well.

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

67. Spontaneous Reaction l Physical or chemical change that will occur without outside intervention. l Nonspontaneous reactions don’t. l Even if they do happen, we can’t say how fast. l Two factors influence spontaneity, l Enthalpy (heat) and Entropy (disorder).

68. Enthalpy and Entropy Factors Exothermic reactions tend to be spontaneous. *Negative  H. 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.

69. 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. l A negative ΔG o like a negative ΔH o represents a release in energy. So  G <0 for all spontaneous reactions.

70.  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

72. 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

73. 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.

74. 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.

75. 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.

76. 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)

77. 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)

78. 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 are 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.

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

80. LeChâtelier’s Principle Regaining Equilibrium

81. 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.

82. 1a. 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

83. 1b. 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

84. 1c. 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

85. 1d. 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.

86. 2a. 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.

87. 2b. 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

88. 3. 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)