1. Reaction Rates
Chapter 18
CP Chemistry

2. Reactions can be…
FAST!
Liquid hydrogen and oxygen reacting to launch a shuttle

3. ..or S L O W
Concrete hardening

4. ..or S L O W
Watching paint dry

5. Expressing rxn rates in quantitative terms

6. Reaction Rates
Red  Blue

7. Collision Theory
Atoms, ions and molecules must collide in order to react.
Reacting substances must collide with the correct orientation.
Reacting substances must collide with sufficient energy to form the activated complex.

8. Orientation and the activated complex
Analogy: if you start with two separate paperclips (reactants) and you wish to link them together (products), not only must the paperclips come into contact, but they also must collide with a specific orientation.

9. Activation energy and reaction
Only collisions with enough energy to react form products

10. Factors affecting reaction rates
Temperature
Concentration
Particle Size
Catalysts

11. TEMPERATURE: Generally, ↑ temp = ↑ rate Why?
Higher temp = faster molecular motion = more collisions and more energy per collision = faster rxn
Analogy: imagine that you are baby-sitting a bunch of 6 year olds. You put them in a yard and you let them run around. Every now and then a couple of kids will run into each other. Now imagine that you decide to feed them some sugar. What happens? They run around faster and of course there are many more collisions. Not only that, the collisions are likely to be a lot harder/more intense.

12. CONCENTRATION
As concentration ↑, frequency of collisions ↑, and therefore rxn rate ↑

13. SURFACE AREA
As surface area ↑, rxn rate ↑
← slow
fast

14. CATALYST pg. 547
Catalyst: a substance that speeds up the rate of a reaction without being consumed in the reaction. (remains unchanged)
Provides an easier way to react
Lowers the activation energy

15. CATALYST
Adding a catalyst speeds up the rxn by lowering the activation energy

16. Chemical Equilibrium

17. Consider a glass of water…
Evaporation

18. Consider a glass of water…
Now, put a lid on it….

19. Consider a glass of water…
Evaporation continues, but condensation also occurs...

20. Consider a glass of water…
The rates equalize, and the system reaches equilibrium.

21. Chemical Equilibrium Consider the following reaction(s):
H2O (liquid)  H20 (gas)
H2O (gas)  H2O (liquid)
H2O (liquid)  H2O (gas)
Equilibrium Symbol

22. Chemical equilibrium occurs when opposing reactions are proceeding at equal rates. The rate at which the products are formed from the reactants equals the rate at which the reactants are formed from the products.
For equilibrium to occur, neither reactant nor products can escape from the system.

23. N H2  2NH KCal
Forward Reaction

24. N H2  2NH KCal
Reverse Reaction

25. Reversible Reactions
REVERSIBLE REACTIONS do not go to completion and can occur in either direction:
aA + bB  cC + dD

26. CHEMICAL EQUILIBRIUM exists when the forward & reverse reactions occur at exactly the same rate
Fig pg. 550
concentration P
Time (reaction progress )
There is no net change in the amounts of reactants and products at equilibrium PAGE 550 (key point)

27. Checkpoint ? Pg. 550
Why is chemical equilibrium called a dynamic state?

28. Answer
Because both the forward and reverse reactions continue

29. At equilibrium:
If there are more products than reactants, the products are said to be favored.
If there are more reactants than products, the reactants are said to be favored.

30. Factors Affecting Equilibrium

31. When a system is at equilibrium, it will stay that way until something changes this condition.

32. Le Chatelier’s Principal
when a change (“stress”) is applied to a system at equilibrium, the system will shift its equilibrium position to counter act the effect of the disturbance.

33. Factors affecting equilibrium include changes in:
Concentrations of reactants or products
Temperature
Pressure (gases)

34. Changes in Concentration:
Consider this reaction at equilibrium:
H2(g) + I2(g)  2HI(g)
What will happen to the equilibrium if we:
add some H2?
Reaction shifts to the right
(forms more product)

35. Changes in Concentration:
Consider this reaction at equilibrium:
H2(g) + I2(g)  2HI(g)
What will happen to the equilibrium if we:
remove some H2?
Reaction shifts to the left
(forms more reactants)

36. Changes in Concentration:
When a substance is added, the stress is relieved by shifting equilibrium in the direction that consumes some of the added substance.
When a substance is removed, the reaction that produces that substance occurs to a greater extent.

37. Changes in Temperature:
Consider this reaction at equilibrium:
2SO2(g) + O2(g)  2SO3(g) kJ
What will happen to the equilibrium if we:
increase the temperature?
Reaction shifts to the left
(forms more reactants)

38. Changes in Temperature:
Consider this reaction at equilibrium:
2SO2(g) + O2(g)  2SO3(g) kJ
What will happen to the equilibrium if we:
decrease the temperature?
Reaction shifts to the right
(forms more products)

39. Changes in Temperature:
Increasing the temperature always favors the reaction that consumes heat, and vice versa.

40. Changes in Pressure: Consider this reaction at equilibrium:
2NO2(g)  N2O4(g)
What will happen to the equilibrium if we:
increase the pressure?
Reaction shifts to the right
(forms more product)

41. Changes in Pressure: Consider this reaction at equilibrium:
2NO2(g)  N2O4(g)
What will happen to the equilibrium if we:
decrease the pressure?
Reaction shifts to the left
(forms more reactant)

42. Changes in Pressure:
Increasing the pressure favors the reaction that produces the fewer moles of gas, and vice-versa.

43. Catalysts
A catalyst increases the rate at which equilibrium is reached, but it does not change the composition of the equilibrium mixture.

44. Action of a Catalyst
Activation Energy
Without a catalyst

45. Lower Activation Energy
Action of a Catalyst
Lower Activation Energy
With a catalyst: A catalyst lowers the activation energy.

46. How would the equilibrium be influenced by:
Example: consider the rxn at equilibrium: N2(g) + 3H2(g)  2NH3(g) + 94 kJ
How would the equilibrium be influenced by:
Increasing the temp:
Decreasing the temp:
Increasing the pressure:
Adding more H2:
Removing some NH3:
Decreasing the pressure:
Adding a catalyst:
rxn shifts to the left
rxn shifts to the right
rxn shifts to the right
rxn shifts to the right
rxn shifts to the right
rxn shifts to the left
no change in equilibrium position

47. Example: How will an increase in pressure affect the equilibrium in the following reactions:
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
RXN SHIFTS LEFT
2H2(g) + O2(g)  2H2O(g)
RXN SHIFTS RIGHT

48. Example: How will an increase in temperature affect the equilibrium in the following reactions:
2NO2(g)  N2O4(g) + heat
RXN SHIFTS LEFT
H2(g) + Cl2(g)  2HCl(g) + 92 KJ
H2(g) + I2(g) + 25 kJ  2HI(g)
RXN SHIFTS RIGHT

49. To close, let’s talk about Fritz Haber and the “process” that made him famous…
Let’s say you want to manufacture ammonia gas (NH3). What are the optimum conditions?
N H2  2 NH3 + Heat
We wish to favor the forward reaction, thus producing more NH3 gas.

50. For example N2 + 3H2  2 NH3 + Heat
By cooling the reaction, the reactions counters by producing heat. It does this by shifting to the right, producing heat, and more NH3 gas.

51. For example N2 + 3H2  2 NH3 + Heat 4 moles of gas 2 moles of gas
By increasing the pressure, the reaction tries to reduce the pressure. It does this by shifting to the side with the fewest moles of gas. This is the product side with 2 moles of gas. Thus the reaction shifts to the right reducing pressure, and producing more NH3 gas.

52. For example N2 + 3H2  2 NH3 + Heat
By adding additional N2 and H2, the reactions tries to use them up. In doing so, the reaction shifts to the right.
By removing NH3 as soon as its formed, the reactions tries to produce more. Shifting the reaction to the right.

53. For example N2 + 3H2  2 NH3 + Heat
So if you want to produce maximum ammonia gas you should
Cool the reaction
Conduct the reaction under high pressure
Add N2 and H2
Remove NH3

54. The Equilibrium Constant, Keq page 556
For the reaction: aA + bB  cC + dD at equilibrium, the constant, kc or keq:
Keq is a measure of the extent to which a reaction occurs; it varies with temperature.

55. Example (a): Write the equilibrium expression for…
PCl5  PCl3 + Cl2

56. Example (b): Write the equilibrium expression for…
4NH O2  4NO + 6H2O

57. Ex: One liter of the equilibrium mixture from example (a) was found to contain mol PCl3, mol Cl2 and mol PCl5. Calculate Keq.
PCl5  PCl3 + Cl2

58. What does keq=0.53 mean to me???
When Keq > 1, most reactants will be converted to products.
When Keq < 1, most reactants will remain unreacted.

59. The equilibrium constant allows us to ….
Predict the direction in which a reaction mixture will proceed to achieve equilibrium.
Calculate the concentrations of reactants and products once equilibrium has been reached.

60. Ex:. PCl3(g) + Cl2(g)  PCl5(g) Keq=1
Ex: PCl3(g) + Cl2(g)  PCl5(g) Keq=1.9 In a system at equilibrium in a 1.00 L container, we find 0.25 mol PCl5, and 0.16 mol PCl3. What equilibrium concentration of Cl2 must be present?

61. Reaction Quotient (Q)
Reaction Quotient (Q) is calculated the same as Keq, but the concentrations are not necessarily equilibrium concentrations.
Comparing Q with Keq enables us to predict the direction in which a rxn will proceed if it is NOT at equilibrium.

62. Comparing Q to Keq
Forward rxn predominates – “reaction proceeds to the right”(until equil. is reached)
System is at equilibrium
Reverse reaction predominates – “reaction proceeds to the left” (until equilibrium is reached)
When Q < Keq:
When Q = Keq:
When Q > Keq:

63. Ex:. H2(g) + I2(g)  2HI(g) Keq for this reaction at 450 C is 49
Ex: H2(g) + I2(g)  2HI(g) Keq for this reaction at 450 C is 49. If 0.22 mol I2, 0.22 mol H2, and 0.66 mol HI are put into a 1.00-L container, would the system be at equilibrium? If not, what must occur to establish equilibrium.
Q < Keq; therefore forward reaction predominates until equilibrium is reached.