Reaction Rate & Equilibrium

1. Reaction Rate & Equilibrium

2. Collision Theory • When one substance is mixed with another, the two substances do not react on a macroscopic basis, but react as their individual particles (atoms, molecules, or ions) come together.

3. The factors that affect how these particles come together are the factors that influence the rate of a reaction.

4. In order for the reaction to occur, the particles involved must collide with each other, the more often the particles collide, the faster the reaction occurs.

5. Not every collision results in a reaction.

6. NO3(g) + CO(g)  NO2(g) + CO2(g)

7. The colliding molecules must have enough energy to react and form an activated complex.

8. This minimum amount of energy is called the activation energy.

9. The activated complex is the temporary arrangement of atoms as they change from reactants into products.

11. Reaction Rate • Is affected by the following: • Temperature • Concentration of reactants • Particle Size/Surface Area • Catalysts • Nature of the Reactant

12. Effect of Temperature on Reaction Rate • Increasing the temperature increases the rate of a reaction.

13. At higher temperatures, the particles have more energy, move faster and collide more frequently.

14. More importantly, the higher energy of the particles means that • more of the particles have sufficient energy (activation energy) • to have a successful collision (form an activated complex).

15. In an exothermic reaction, the addition of a small amount of heat can often have a dramatic effect on the rate of reaction.

16. This small addition of energy gives a few particles the activation energy, they react and release more energy, which is absorbed by more particles, which react and release more energy, …

17. Effect of Concentration on Reaction Rate • Increasing the concentration of the reactants increases the rate of a reaction. • When there are more particles per unit volume, they will collide more frequently, causing the rate to increase.

18. Effect of Particle Size/Surface Area • Decreasing particle size/increasing surface area increases the rate of a reaction. • When the surface area is increases, there is more contact between the reactants, the number of collisions increase, and therefore the rate of reaction also increases.

19. Effect of Catalysts • A catalyst is a substance that increases the rate of a chemical reaction without being permanently consumed by the reaction.

20. A catalyst provides a different pathway or mechanism for the reaction that has a lower activation energy. • When the activation energy is lower, more of the particles will have sufficient energy to react, more of the collisions are successful therefore the reaction rate increases.

21. A catalyst increases the rate of a reaction by lowering the activation energy.

22. Effect of the Nature of the Reactant • The nature of the reactants involved will determine the kind of reaction that occurs.

23. Reactions with bond rearrangements or electron transfer take longer than reactions without these changes. • Ionic reactions (such as double displacement or neutralization) occur almost instantly.

24. In reactions with acids, stronger acids (H2SO4, HNO3, HClO4, HCl, HF, and HI) will occur faster than reactions with weak acids (all other acids).

25. In reactions with metals, more active metals (see activity series) will react faster than less active metals.

26. Reaction Mechanism • Most reactions occur in a series of steps. Each step normally involves the collision of only two particles.

27. There is little chance of three or more particles colliding with the proper position and sufficient energy to cause a reaction.

28. If a reaction consists of several steps such as the following: • A … B • B … C • C … final product

29. One of the steps will be slower than all the others. This step is called the rate determining step. • The other faster steps will not affect the rate of the reaction.

30. The series of steps that must occur for a reaction to go to completion is called the reaction mechanism.

31. CATALYSIS • · A Catalyst is a substance that speeds up a reaction without being consumed.

32. In theory the catalyst may be used over and over again.

33. In practice, however, there is some loss of catalyst through other reactions that occur at the same time (side reactions).

34. Importance of Catalysts: • · Catalysts allow reactions to occur much faster. • · Catalysts allow reactions to occur at lower temperatures (energy savings).

35. Characteristics of Catalysts: • Catalysts are often quite specific. They increase the rate of some reactions but not others. • Enzymes (catalysts used in biological organisms) are extremely selective.

36. How Catalysts Work: • A catalyst is an active participant in the reaction. • A catalyst must participate in one step of a reaction and be regenerated in a later step.

38. The catalyzed reaction mechanism makes available a reaction path having an increased • overall rate of reaction by: - increasing the frequency factor, or - decreasing the Activation Energy, Ea (most dramatic effect)

39. Example:The commercial preparation of sulfuric acid, H2SO4 from SO2 NO catalyst First step: 2 SO2g) + O2(g) 2 SO3(g) Proposed Mechanism: • 2NO + O2 2 NO2 • NO2 + SO2NO + SO3(must occur twice each time the first step occurs once) • catalyst • 2NO + O2 2NO2 reaction intermediate • 2NO2 + 2SO22NO + 2SO3 ___________________________________ Overall:2 SO2 + O2 2SO3

40. Equilibrium Systems • In some chemical reactions, all of the reactants are converted to products: 2Cu(s) + O2(g)  2CuO(s) This reaction is said to go to completion.

41. In other processes, however, the final mixture consists of both products and reactants. Reactants  Products Reactants  Products Reactants  Products This reaction is at equilibrium.

42. The reactants are colliding to form products while the products are also colliding to reform reactants.

43. This is known as a dynamic equilibrium – the rate of the forward reaction is equal to the rate of the reverse reaction.

44. Equilibrium is reached in a closed system when all properties of the system are constant. • The forward and reverse reactions are happening at the same rate.

45. The concentrations of products and reactants may be different, but they are no longer changing.

46. Le Châtelier’s Principle • If a stress is applied to a system at equilibrium, the equilibrium will shift to relieve the stress.

47. Concentration • Increase [reactant] = shifts to use up the added reactants and produce more products • 􀂾 Shifts right • Increase [product] = shifts to use up added product and product more reactants • 􀂾 Shifts left • Decrease [reactant] = shifts to produce more reactants • 􀂾 Shifts left • Decrease [product] = shifts to produce more products • 􀂾 Shifts right • • • • •

48. Temperature • Increase = favours the endothermic direction. • 􀂾 The reaction shifts in the direction that uses up the added energy. • Decreasing = favours the exothermic direction. • 􀂾 The reaction shifts in the direction that produces energy. • • •

49. Pressure/Volume • Increasing pressure/Decreasing volume favours fewer gas molecules. • 􀂾 The reaction shifts to relieve the pressure. • Decreasing pressure/Increasing volume favours more gas molecules. • 􀂾 The reaction shifts to restore the pressure. • • •

50. Ta Da!