Chapter 8: Kinetics and Equilibrium

1. Chapter 8:Kinetics and Equilibrium

2. Do Now: Answer these questions • 1. What type of energy is related to the temperature of a substance? (think about definition of temperature) • 2. If the temperature of a substance is increased, what do you think happens to the speed of its particles? • 3. In the diagram below, which container has a higher concentration of Na+ and Cl- ions dissolved in water?

3. https://www.youtube.com/watch?v=WTwE7xDZkPk

4. Small Difference Between Success and Failure

5. What is Kinetics? • Kinetics is the branch of chemistry that deals with the rates of chemical reactions • Collision theory: • 1. In order for a reaction to occur, reactants must collide with each other • 2. Effective collisions are when reactants come together with the correct energy and position to form a product

6. Reading Reaction Equations A + B  AB • A & B are reactants, AB is product • “Double arrow” means reaction can move in either direction • Reverse reaction: AB  A + B

7. Aim: What factors affect rate of reaction? • As the amount of effective collisions increases, product formation increases (reaction rate increases) • More successful interactions make more product • How can we increase the reaction rate (by increasing number of effective collisions)?

8. What Affects Rate of Reaction? • Factors affecting reaction rates: • Nature of reactants • Concentration • Surface area • Pressure • Temperature • Catalysts

9. Reaction Rate Factors:Nature of Reactants Reactions involve the breaking of old bonds and making of new bonds. In general: • 1. Covalent bonds are slower to react then ionic bonds • 2. Breaking bonds requires more energy than making bonds through collisions

10. Reaction Rate Factors:Concentration • Generally, increasing concentration increases the rate of reaction • Especially if volume is decreased

11. Reaction Rate Factors:Surface Area • Generally, the more surface area that is exposed, the more chances there are for effective collisions that increase rate of reaction

12. Surface Area Demos • Sugar cube vs. granulated sugar • Alka Seltzer™, chunk vs. crushed

13. Reaction Rate Factors:Pressure • No effect on solids and liquids, only gases • Increase pressure  smaller volume  higher rate of effective collisions  higher rate of reaction

14. Reaction Rate Factors:Temperature • Generally, increasing temperature increases kinetic energy of molecules • More kinetic energy  more effective collisions  higher rate of reaction

15. Demo with temperature-dependent glow stick reaction rates

16. Catalysts are substances that alter the speed of a chemical reaction without being permanently changed themselves Adding a catalyst increases the rate of reaction by providing a different and easier pathway for the reaction Reaction Rate Factors: Catalysts

17. Review Questions Complete review book pages 137 & 138, #1-10

18. Conclusion Questions • Describe what happened to the glow sticks • Explain what occurred in the hot water and the cold water (include: speed of the particles, effective collisions, and rate of reaction in your answer)

19. Aim: How can we classify energy in chemical reactions? • Do Now: Answer the following questions: 1. What do scientists study in the branch of chemistry called kinetics? 2. What does the Collision Theory describe? 3. What are effective collisions?

20. 3 Ways to Classify Energy in Reactions • 1. Look at the reaction (where is the Energy term?) • CH4 + 2 O2 CO2 + 2 H2O + 890.4 kJ • N2 + O2 + 66.4 kJ  2 NO2 • If energy is with the reactants, it’s being added in and the process is endothermic • If energy is with the products, it’s being released and the process is exothermic

21. 3 Ways to Classify Energy in Reactions • 2. Look at the ∆H, the heat of reaction Listed on Table I in reference tables

22. Using PE Values to Classify Reactions Example: A + B  AB Compare the PE of the reactants (left) to the PE of the products (right) in the forward reaction PE reactants: 75 Joules PE products: 25 Joules Energy of products is lower than reactants, so energy must have been released, making the reaction exothermic

23. 3 Ways to Classify Energy in Reactions • 3. Potential Energy (PE) Diagrams

24. How to Interpret Potential Energy Diagrams • Show how the potential energy of reactants changes to chemical energy in bonds • PE diagrams track PE changes during chemical reactions

25. How to Read PE Diagrams

26. Let’s Label the PE Diagram • PE of reactants • PE of products • ∆H (heat of reaction): ∆H = PE products - PE reactants • If ∆H is negative, (PE prod < PE reac) and reaction is EXOTHERMIC

27. Interpreting PE Diagrams Example: AB  A + B Compare the PE of the reactants (left) to the PE of the products (right) in the forward reaction PE reactants: 25 Joules PE products: 75 Joules Energy of products is higher than reactants, so energy must have been absorbed, making the reaction endothermic

28. How to Read PE Diagrams

29. Let’s Label the PE Diagram • PE of reactants • PE of products • ∆H (heat of reaction): ∆H = PE products - PE reactants • If ∆H is positive, (PE prod > PE reac) and reaction is ENDOTHERMIC

30. Activated Complex • Activated Complex: intermediate molecule that forms after effective collision • Unstable and temporary • Activation Energy: energy needed to start forward reaction

33. Catalyzed Reaction • Activation Energy (with catalyst): energy needed to start reaction if a catalyst is added

34. PE Diagram Review • Complete review book pages 140 & 141, # 11-17

35. Do Now: Answer questions 4 NH3 + 5 O2 4 NO + 6 H2O + 425 kJ • 1. Is the reaction endothermic of exothermic? • 2. If you placed your hand on the container where this reaction was taking place, would it feel hot or cold? • 3. What is the ∆H for the reaction? • 4. Draw a sketch of what the PE diagram for this reaction could look like

36. Aim: How can we look at systems at equilibrium? • Equilibrium: a state of balance • Dynamic Equilibrium: when the rate of a forward reaction is equal to the rate of its reverse reaction; doesn’t mean amounts are equal, only rates • Reversible Equilibrium: many reactions in equilibrium can be reversed, as indicated by a double arrow

37. Types of Equilibrium • Physical Equilibrium is balance between changes taking place during physical processes like changes of state or dissolving • Phase equilibrium: equilibrium between phases • Solid/liquid phases: in a closed system, rate of ice melting equals rate of water freezing: H2O (s)  H2O (l) • Liquid/gas phases: in a closed system, rate of evaporation equals rate of condensation: H2O (l)  H2O (g)

38. Solution Equilibrium • Solid/liquid solutions: in closed systems, saturated solutions are in equilibrium when rate of dissolving equals rate of recrystallization • Extra Kool-Aid powder at the bottom of a glass • Liquid/gas solutions: in closed systems, equilibrium exists between gaseous and dissolved states • CO2 dissolved in soda

39. Aim: How will a system react when a stress if applied?

40. Le Chatelier’s Principle • If a stress is applied to a system at equilibrium, the equilibrium will shift to release the effects of the stress • Stressors: • Temperature • Concentration • Pressure

41. Le Chatelier’s Principle:Temperature • The side you shift towards will increase and the side you shift away from will decrease • Increase in temperature favors the endothermic reaction (where heat/energy is absorbed) • Decrease in temperature will favor the exothermic reaction (where heat/energy is released)

42. Cobalt chloride demo? Do demo before showing lesson as motivation

43. Le Chatelier’s Principle:Concentration • Make chart like in notes packet to show effects

44. Another cobalt chloride demo?

45. Temperature/Concentration • Add Away (increase) • Take Towards (decrease) • Pen/pencil will go up on whichever side you shift towards

46. Do Now: Answer the following questions • 1. What is Le Chatelier’s Principle? • 2. Using the following equation, label the forward and reverse reactions as endothermic or exothermic CO (g) + 3 H2 CH4 (g) + H2O (g) + heat • 3. In which direction will equilibrium shift if the temperature is increased? • 4. In which direction will equilibrium shift if the temperature is decreased?

47. Pressure • Need to know how many gas molecules are on the reactant side and on the product side • If equal # of gas molecules on each side, changing pressure will have no effect • Increasing pressure shifts from side with more gas molecules towards side with fewer gas molecules • Decreasing pressure shifts from side with fewer gas molecules towards side with more gas molecules

48. Pressure Example • 4 NH (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g) 4 + 5 = 9 gas molecules 4 + 6 = 10 gas molecules • Increase pressure: shift left (more to less) 9  10 • Decrease pressure: shift right (less to more) 9 10

49. Conclusion Questions • 1. How does a system at equilibrium respond to a stress? List the factors that can be stresses on an equilibrium system.

50. Conclusion Questions • 2. Given the reaction at equilibrium: N2 (g) + 3 H2 (g)  2 NH2 (g) + heat • Which stress would cause the equilibrium to shift to the left? • (a) Increasing the temperature • (b) Adding N2 (g) to the system • (c) Adding H2 (g) to the system