Reaction Rates

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

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. N2 + 3H2 2NH3 + 22 KCal Forward Reaction

24. N2 + 3H2 2NH3 + 22 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 EQUILIBRIUM R Fig. 18.10 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) + 198 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) + 198 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. Action of a Catalyst Lower Activation Energy With a catalyst: A catalyst lowers the activation energy.

46. 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 • RXN SHIFTS LEFT • 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? N2 + 3H2 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.