Reaction Rates & Equilibrium. What determines how fast a reaction takes place? What determines the extent of a reaction?

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1 Reaction Rates & Equilibrium What determines how fast a reaction takes place? What determines the extent of a reaction? Reactants Products 1

2 Reaction Rates Vary TNT exploding. A car rusting. Dead plants turning into coal. 2

3 Rate Units The rate of a chemical reaction is the change in concentration with time. DMolarity Dtime = M t 3

4 Collision Theory B A + B products A & B must collide for a reaction to take place. Aprod. B A 4

5 Collision Theory A + B products A & B must collide with enough kinetic energy for bonds to break. During the collision this KE is converted into PE to break bonds. 5

6 Activation Energy The minimum energy that colliding particles must have in order to react is called activation energy (E a ). 6

7 Potential Energy Potential Energy Diagram activated complex E a reactants DH exothermic products Reaction progress 7

8 Heat of Reaction (Enthalpy) DH = PE products - PE reactants + endothermic vs. - exothermic 8

9 Potential Energy Potential Energy Diagram E a DH prod reactants endothermic Reaction progress 9

10 Activation Energy Some reactions don t proceed at room temperature because they have high activation energy. C + O 2 CO 2 10

11 Try It!!! Sketch the P.E.diagram for: 2CO + O 2 2CO 2 (Get DH from Regents Table I) What is DH of reverse reaction? Is the activation energy of the reverse reaction the same? 11

12 Self-Sustaining Reactions Exothermic Endothermic E a DH E a DH Heat given off Heat absorbed supplies E a does not supply E 12 a

13 PE Diagram: Try It The burning of methane (CH 4 ) has E a = 240 kj. Draw a PE diagram. Is the reaction exo- or endothermic? Is it self-sustaining? What is E a of the reverse reaction? 13

14 Molecular Orientation 2NOBr 2NO + Br 2 reaction reaction no reaction 14

15 Collision Theory 15

16 Reaction Rate What factors affect reaction rate? Hint:Think about collision theory. temperature concentration particle size catalyst 16

17 Temperature Reactions speed up at higher temperature. WHY? Higher T more KE more collisions and more collisions exceed E a 17

18 Temperature Higher T increases reaction rate: food spoiling wood burning corrosion 18

19 Concentration More particles More collisions Faster reaction rate Wood fire in air vs. Wood fire in pure O 2 19

20 Particle Size Why do you start a fire with kindling rather than whole logs? Surface area 20

21 Particle Size Smaller particles larger surface area faster reaction rate Coal dust explosions Flour mill explosions Dissolving sugar 21

22 Catalysts A substance that increases reaction rate without being used up in the reaction. Catalytic converter 22

23 Potential Energy Catalysts: Lower A.E. reactants products Reaction progress 23

24 Catalysts Catalysts work by providing an alternate pathway for the reaction that has lower E a. Catalysts are important in: industrial processes biological reactions (enzymes). 24

25 Collision Theory: Summary 25

26 Reversible Reactions Many reactions go in both the forward and reverse directions simultaneously. 2SO 2 (g) + O 2 (g) forward left right reverse 2SO 3 (g) 26

27 Reversible Reactions N 2 O 4 (g) colorless 2NO 2 (g) brown What happens if you start with only N 2 O 4 in a container? 27

28 N 2 O 4 (g) 2NO 2 (g) t 1 t 2 t 3 t 4 no change 28

29 Chemical Equilibrium N 2 O 4 (g) 2NO 2 (g) At some point: forward rate = reverse rate and there is no net change in reactants or products. 29

30 concentration concentration N 2 O 4 (g) Start with N 2 O 4 2NO 2 (g) Start with NO 2 N 2 O 4 N 2 O 4 NO 2 NO 2 time time 30

31 Chemical Equilibrium Point of no net change in amounts of reactants or products. A B dynamic equilibrium rate forward rxn = rate reverse rxn 31

32 Chemical Equilibrium N 2 (g) + 3H 2 (g) 2NH 3 (g) forward rate = reverse rate Rate of making NH 3 equals rate of making N 2 and H 2 Does NOT mean amount of NH 3, N 2 and H 2 are equal. 32

33 Physical Equilibrium For example: Ice water at 0 o C rate ice melts = rate water freezes H 2 O(s) H 2 O(l) 33

34 Physical Equilibrium Example: saturated solution rate dissolving = rate crystallizing sugar(s) sugar(aq) solution (sugar water) solute (undissolved sugar) 34

35 Equilibrium Constant A quantitative way to determine the amount of reactants and products at equilibrium. 35

36 Equilibrium Constant General reaction: aa + bb K eq = [C]c [D] d [A] a [B] b cc + dd Exponents are the coefficients Square brackets [ ] mean M (concentration in moles per liter) 36

37 Equilibrium Constant K eq = [C]c [D] d [A] a [B] b products on top For a given reaction, K eq depends only on temperature. 37

38 Equilibrium Constant aa + bb cc + dd K eq = [C]c [D] d [A] a [B] b K eq > 10 mostly products K eq < 0.1 mostly reactants 38

39 Try It: Write K eq N 2 O 4 (g) 2NO 2 (g) N 2 (g) + 3H 2 (g) 2NH 3 (g) 39

40 Equilibrium Constant N 2 O 4 (g) 2NO 2 (g) At equilibrium in a 2.0 L flask, there are moles of N 2 O 4 and mol of NO 2. What is K eq? 40

41 K eq : Try It!!! N 2 + O 2 2NO mol The number of moles of each species in a 4.0 L container at equilibrium are shown above. Calculate K eq. 41

42 Factors Affecting Equilibrium Equilibrium is a delicate balance. H 2 CO 3 CO 2 + H 2 O 42

43 Le Chatelier s Principle If a stress (change) is applied to a system at equilibrium, the equilibrium will shift to offset the stress! concentration temperature pressure (gas) volume (gas) 43

44 Le Chatelier s Principle 1. Concentration stress If a substance is added to a reaction at equilibrium (the stress ), equilibrium will shift to use up the added substance. 44

45 Concentration Stress H 2 CO 3 CO 2 + H 2 O If add CO 2 (stress), equilibrium shifts to use up some CO 2. Reaction proceeds to the left. 45

46 Concentration Stress H 2 CO 3 CO 2 + H 2 O If remove CO 2 (stress), equilibrium shifts to make more CO 2. Reaction proceeds to the right. 46

47 Conc. Stress: Try It N 2 (g) + O 2 (g) 2NO(g) Which way does the equilibrium shift if: O 2 is removed? NO is added? What happens to O 2 if N 2 is added? 47

48 Concentration Stress 3H 2 (g) + N 2 (g) 2NH 3 (g) What happens if H 2 is added to this system at equilibrium? 48

49 Concentration 3H 2 (g) + N 2 (g) 2NH 3 (g) Initial H 2 NH 3 H 2 added Final N 2 Time 49

50 Temperature Stress Raising temperature (adding heat) causes a reaction to shift to absorb heat. N 2 O 4 (g) colorless 2NO 2 (g) brown DH = 58 kj 50

51 Temperature Stress 58 kj + N 2 O 4 2NO 2 Endothermic: thus reaction proceeds to the right to use up heat. 20 o C 80 o C 51

52 Temperature Stress CoCl H 2 O Co(H 2 O) Cl - + heat Cooling the mixture causes the solution to turn pink. Is DH + or -? Exothermic 52

53 Pressure Stress Increase in pressure (or decrease the volume) will shift equilibrium to the side of the reaction with fewer moles of gas. 53

54 Pressure Stress N 2 (g) + 3H 2 (g) 2NH 3 (g) 4 moles gas 2 moles gas Increase pressure (or decrease volume) shifts equilibrium to the right. 54

55 Le Chatelier: Try It!!! N 2 F 4 (g) 2NF 2 (g) DH = 38.5 kj What happens if: NF 2 is added Pressure is decreased N 2 F 4 is removed Temp. is increased 55

56 Le Chatelier: Try It!!! N 2 (g) + 3H 2 (g) 2NH 3 (g) What happens if: N 2 is added P is increased from 1 to 2 atm NH 3 is removed Temp. lowered from 0 o C to -5 o C Helium gas is added A catalyst is added 56

57 Thermodynamics How do you know which reactions will occur? light Bunsen burner spontaneous 57

58 Nonspontaneous Reactions Many reactions won t go even if you can write an equation. 2H 2 O 2H 2 + O 2 Forget about it. 58

59 Spontaneous vs. Nonspontaneous Reactions How do you know? Two factors: Enthalpy and Entropy 59

60 energy 1.Enthalpy (Heat of Reaction) A spontaneous reaction is favored by giving off heat. (exothermic) DH - reactants products 60

61 1. Enthalpy (Heat) For example: 4Al + 3O 2 2Al 2 O 3 DH = kj Gives off heat (exothermic) and is spontaneous. 61

62 Enthalpy: Not the Only Factor! An ice cube melts at 25 o C even though melting is endothermic. Why? H O(s) H O(l) 2 2 DH = +6.0 kj 62

63 Entropy: Gedanken Gas Y at 1 atm Gas B at 1 atm What happens when the valve is opened? Why? 63

64 Entropy (S) The gases mix because of entropy. 64

65 Entropy (S) If you drop a glass, what happens? If you drop the broken pieces of glass, do they reform the glass? 65

66 Changing Entropy (S) Increasing entropy solid liquid gas 66

67 Increasing Entropy (S) 1. Solid Liquid Gas 2. Dissolving a solute. NaCl(s) H 2 O Na + (aq) + Cl - (aq) 67

68 Increasing Entropy (S) 3.Reactions where the moles of gas products is greater than the moles of gas reactants. CaCO 3 (s) CaO(s) + CO 2 (g) 0 moles gas 1 mole gas 68

69 Increasing Entropy (S) 4.Raising temperature increases kinetic energy which increases randomness. 69

70 Entropy; Try It!! 1.Is S of Na + (aq) and Cl - (aq) ions increased or decreased when salt water evaporates? 2.What about the S of the water in this process? 3.Is S higher or lower for: 2H 2 (g) + O 2 (g) 2H 2 O(g) 70

71 Spontaneous Reactions The combination of enthalpy (DH) and entropy (DS) determines whether a reaction is spontaneous. Exothermic (-DH) favors spontaneous More random (+DS) favors 71

72 Spontaneous Reaction? DH + (endo) DH (exo) DS + (more random products) DS (less random products) 72

73 Predict: Spontaneous? (See Table I) N 2 (g) + 2O 2 (g) 2NO 2 (g) C(s) + O 2 (g) CO 2 (g) 73

74 Spontaneous Reactions Need a quantitative way to determine if a reaction is spontaneous. Gibbs Free Energy DG 74

75 Gibbs Free Energy DG = DH - TDS (T in Kelvin) DG is (-) spontaneous DG is (+) nonspontaneous 75

76 Gibbs Free Energy DG = DH - TDS Note the units: DG usually in kj or kcal DH usually in kj or kcal DS usually in J/K or cal/k T must be in K 76

77 Free Energy: Break It Down! DG = DH - TDS CaCO 3 (s) DH = 178 kj CaO(s) + CO 2 (g) DS = 161 J/K Is this reaction spontaneous at 25 o C? at 1000 o C? 77

78 Summary 1.Thermodynamics (energy): will a reaction take place? (DH & DS) 2.Kinetics: speed of a reaction (collision theory; PE diagram) 3.Equilibrium: extent of a reaction (LeChatelier) 78

79 79

80 Warm-up What is collision theory? What two things must be true about the collisions for a chemical reaction to occur? 80

81 Warm-up Draw a PE diagram for the formation of NH 3 from its elements (assume E a = 310kJ). What is the value of E a for the reverse reaction? 81

82 Warm-up Why does increasing the temperature make reactions speed up? 82

83 Warm-up What factors affect the rate of a reaction? Explain each using collision theory. 83

84 Warm-up 2CO + O 2 2CO 2 Write an expression for the equilibrium constant for this reaction. 84

85 What is meant by Warm-up chemical equilibrium? Moles of each chemical at equilibrium in a 3.0L flask at 150 o C are: N 2 (g) + 3H 2 (g) 2NH 3 (g).37mol 1.2mol 0.52mol What is K eq? 85

86 Warm-up Draw a PE diagram for the reaction of carbon and hydrogen to form C 2 H 2.(E a is 352 kj) What is E a for the reverse reaction? 86

87 Warm-up 2CO + O 2 2CO 2 For this system at equilibrium: Which way does the equilib. shift if O 2 is removed? What happens to O 2 if CO is added? 87

88 Warm-up 2CO + O 2 2CO 2 For this system at equilibrium: What happens if the system is cooled? 88

89 Warm-up For the reaction of nitrogen and oxygen forming nitrogen dioxide, what is the effect on equilibrium of lowering the temperature? 89

90 Warm-up What is entropy? Give a synonym. Is S higher (+entropy) or lower (-entropy) when: water is frozen? you clean your room? salt is dissolved? aluminum is burned? 90

91 Warm-up (Honors) Al reacts with O 2 to form Al 2 O 3 : Draw PE diagram (assume AE is 890 kj) What is the AE of reverse reaction? What 4 things would speed up the rxn? Honors: Write the expression for K eq Which way does equilib. shift if: more O 2 is added? If T is raised? As the reaction proceeds, does entropy increase or decrease? 91

92 Warm-up (Regents) Al reacts with O 2 to form Al 2 O 3 : Draw PE diagram (assume AE is 890 kj) What is the AE of reverse reaction? Name 4 things would speed up the rxn? Which way does equilib. shift if: more O 2 is added? If T is raised? As the reaction proceeds, does entropy increase or decrease? 92