EQUILIBRIUM

1. EQUILIBRIUM CORE (5 HRS) + AHL (4 HRS)

2. IB Core Objective • 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium. Command Term • Outline: Give a brief account or summary. (Obj. 2)

3. N2O4 (g) 2 NO2 (g) 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium. • Many chemical reactions do not go to completion • Once some products are formed the reverse reaction can take place to reform the reactants • Notice the double arrows in the equation!

4. N2O4 (g) 2 NO2 (g) 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium. • In a closed system the concentrations of all the reactants and products will eventually become constant • Such a system is said to be in a state of ‘dynamic equilibrium’ • The forward and reverse reactions continue to occur, but at equilibrium the rate of the forward reaction is equal to the rate of the reverse reaction.

5. N2O4 (g) 2 NO2 (g) Explain each diagram!!!

6. H2O (l) H2O (g) 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium. • Dynamic equilibrium also occurs when physical changes take place • In a closed flask, containing some water, equilibrium will be reached between the liquid water and vapour • The faster moving molecules in the liquid will escape form the surface to become vapour and the slower moving molecules in the vapour will condense back into liquid • Equilibrium will be established when the rate of vaporizationequals the rate of condensation

7. Describe what factors are changing from left to right!!!

8. IB Core • 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. • Deduce: Reach a conclusion from the information given. (Obj. 3)

9. 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. • Forward reaction: N2O4 (g) 2 NO2 (g) • Rate law: Rate = kf [N2O4] • Reverse reaction: 2 NO2 (g) N2O4 (g) • Rate law: Rate = kr [NO2]2

10. [NO2]2 [N2O4] kf kr = 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. • Therefore, at equilibrium Ratef = Rater kf [N2O4] = kr [NO2]2 • Rewriting this, it becomes

11. [NO2]2 [N2O4] Kc = kf kr = 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. The ratio of the rate constants is a constant at that temperature, and the expression becomes Notice that for [NO2] its coefficient in the equation becomes an exponent in the equilibrium equation, also KC is known as the equilibrium constant

12. aA + bB cC + dD [C]c[D]d [A]a[B]b Kc = 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. • To generalize this expression, consider the reaction • The equilibrium expression for this reaction would be

13. What Are the Equilibrium Expressions for These Equilibria?

14. 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogenous reaction. • Homogeneous reactions??? • ‘Homo’ means same! • Therefore, these types of reactions are all in the same phase whether it be gas, liquid, or aqueous. • Equilibrium constants (KC) are calculated from homogenous reactions

15. IB Core 7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant. Deduce: Reach a conclusion from the information given. (Obj. 3)

16. 7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant. • Since the equilibrium expression has the concentration of products on the top and the concentration of reactants on the bottom, it follows that the magnitude of the equilibrium constant is related to the position of equilibrium • When the reaction goes nearly to completion KC >> 1 • If the reaction hardly proceeds then KC << 1

17. 7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant. • If K >> 1, the reaction is product-favored; product predominates at equilibrium. • If K << 1, the reaction is reactant-favored; reactant predominates at equilibrium. • If the value of KC lies between about 10-2 and 102then both reactants and products will be present in the system in noticeable amounts

18. IB Core 7.2.3 Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant. Apply: Use an idea, equation, principle, theory or law in a new situation. (Obj. 2)

19. 7.2.3 Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant. Le Châtelier’s Principle “If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.” No need to ever state this principle!!!

20. Le Châtelier’s Principle (7.2.3) • Provided the temperature remains constant the value of KC must remain constant • If the concentration of the reactants is increased, or one of the products is removed from the equilibrium mixture then more of the reactants must react in order to keep KC constant, i.e. the position of the equilibrium will shift to the right (towards more products)

21. Le Châtelier’s Principle (7.2.3) • Homework: Use your textbook, check out the interactive link on the uaschemistry site, and use any other suitable resources!

22. IB Core 7.2.4 State and explain the effect of a catalyst on an equilibrium reaction. State: Give a specific name, value or other brief answer without explanation or calculation. (Obj. 1) Explain: Give a detailed account of causes, reasons or mechanisms. (Obj. 3)

23. 7.2.4 State and explain the effect of a catalyst on an equilibrium reaction. • A catalyst will increase the rate at which equilibrium is reached, as it will speed up both the forward and reverse reactions equally, but it will have no effect on the position of equilibrium and hence on the value of KC

24. IB Core 7.2.5 Apply the concepts of kinetics and equilibrium to industrial processes. Apply: Use an idea, equation, principle, theory or law in a new situation. (Obj. 2)

25. 7.2.5 Apply the concepts of kinetics and equilibrium to industrial processes. • Suitable examples include the Haber and Contact processes • Homework: Use your textbook and any other suitable resources!

26. AHL: Liquid-vapour equilibrium (17.1) • Describe the equilibrium established between a liquid and its own vapour and how it is affected by temperature changes (17.1.1) • Sketch graphs showing the relationship between vapour pressure and temperature and explain them in terms of the kinetic theory (17.1.2) • State and explain the relationship between enthalpy of vaporization, boiling point and intermolecular forces (17.1.3) • Homework: Use your textbook and any other suitable resources! Present this data to the class…random selection!

27. IB AHL 17.1.1 Describe the equilibrium established between a liquid and its own vapour and how it is affected by temperature changes. Describe: Give a detailed account. (Obj 2)

28. 17.1.1 Describe the equilibrium established between a liquid and its own vapour and how it is affected by temperature changes. Things to Consider • How would you state this in as few words as possible? • What is happening during this equilibrium?

29. IB AHL 17.1.2 Sketch graphs showing the relationship between vapour pressure and temperature and explain them in terms of kinetic theory. Sketch: Represent by means of a graph showing a line and labelled but unscaled axes but with important features (for example, intercept) clearly indicated. (Obj. 3)

30. 17.1.2 Sketch graphs showing the relationship between vapour pressure and temperature and explain them in terms of kinetic theory. Things to consider • Describe how kinetics are involved in the rate of vapourisation. • How would a graph look showing this relationship if you had to sketch this on an IB exam?

31. IB AHL 17.1.3 State and explain the relationship between enthalpy of vaporization, boiling point and intermolecular forces. State: Give a specific name, value or other brief answer without explanation or calculation. (Obj. 1) Explain: Give a detailed account of causes, reasons or mechanisms. (Obj. 3)

32. 17.1.3 State and explain the relationship between enthalpy of vaporization, boiling point and intermolecular forces. Things to consider • What are the definitions of enthalpy of vaporization and boiling point? • Describe how intermolecular forces are related to the enthalpy of vapourisation and boiling point. • Describe how pressure is related to temperature and equilibrium. (Real world: What problems would I encounter cooking a pot of pasta high in the mountains around Bend, Oregon? How could I solve this?)

33. IB AHL 17.2.1 Solve homogenous equilibrium problems using the expression for Kc. Solve: Obtain an answer using algebraic and/or numerical methods. (Obj. 3)

34. H2 (g) + I2 (g) 2 HI (g) The equilibrium law (17.2) A closed system initially containing 1.000 x 10−3 M H2 and 2.000 x 10−3 M I2 At 448C is allowed to reach equilibrium. Analysis of the equilibrium mixture shows that the concentration of HI is 1.87 x 10−3 M. Calculate Kc at 448C for the reaction taking place, which is

35. The equilibrium law (17.2) What do we know?

36. The equilibrium law (17.2) [HI] Increases by 1.87 x 10-3M

37. The equilibrium law (17.2) Stoichiometry tells us [H2] and [I2] decrease by half as much

38. The equilibrium law (17.2) We can now calculate the equilibrium concentrations of all three compounds…

39. Kc= [HI]2 [H2] [I2] (1.87 x 10-3)2 (6.50 x 10-5)(1.07 x 10-3) = = 51 The equilibrium law (17.2) …and, therefore, the equilibrium constant What does the equilibrium constant value mean regarding the reaction?