1. CHEMISTRY XL-14A CHEMICAL EQUILIBRIA August 20, 2011Robert Iafe

2. Unit Overview 2  Reactions at Equilibrium  Equilibrium Calculations  Le Châtelier’s Principle  Catalysts

3. Reactions at Equilibrium 3  Reversibility of Reactions  Law of Mass ActionK  Gas Phase EquilibriumK P  Solution Phase EquilibriumK C  Extent of Reaction using K  Direction of ReactionQ vs K

4. Introduction to Chemical Equilibrium 4 We previously assumed all reactions proceeded to completion However: many reactions approach a state of equilibrium Equilibrium – condition of a chemical reaction in which chemical change ceases and no further change occurs spontaneously Equilibrium – a dynamic equilibrium between reactants and products in a chemical reaction. At Equilibrium: Forward and reverse reactions simultaneous with equal rates No further net conversion of reactants to products unless the experimental conditions have been changed The equilibrium state is characterized by the Equilibrium Constant (K eq )

5. Reversibility of Reactions 5 A. N 2 (g) + 3 H 2 (g) 2 NH 3 (g) B. 2 NH 3 (g) N 2 (g) + 3 H 2 (g) AB

6. Reversibility of Reactions 6 Add conc. HCl [Co(H 2 O) 6 ] 2+ Pink [CoCl 4 ] 2- Blue [Co(H 2 O) 6 ] 2+ + [CoCl 4 ] 2- Purple Dissolve CoCl 2. 6H 2 O in pure waterDissolve in 10 M HCl Add H 2 O [Co(H 2 O) 6 ] 2+ + 4 Cl - [CoCl 4 ] 2- + 6 H 2 O

7. The Nature of Chemical Equilibrium 7 [Co(H 2 O) 6 ] 2+ + 4 Cl - [CoCl 4 ] 2- + 6 H 2 O As time progresses, rate = concentration/time slows When rate(s) = constant, reaction is at equilibrium %Co 100 98 2 A B A B C1C1 t1t1 C2C2 t2t2 [Co(H 2 O) 6 ] 2+ [CoCl 4 ] 2-

8. The Nature of Chemical Equilibrium Characteristics of the Equilibrium State 8 1.No Macroscopic Change 2.Reached Spontaneously 3.Dynamic Balance of Forward/Reverse Processes 4.Same regardless of direction of approach [Co(H 2 O) 6 ] 2+ + 4 Cl - [CoCl 4 ] 2- + 6 H 2 O Situations which appear to be equilibrium, but are not: Steady State – macroscopic concentrations are constant, even though system is not at equilibrium 1 process removes species while 2 nd process supplies species Homeostasis – body tries to maintain blood pH, etc…

9. Equilibrium Data 9 2 SO 2 (g) + O 2 (g) 2 SO 3 (g) at 1000. K 1864: Norwegians Cato Guldberg and Peter Waage discovered the mathematical relationship that summarized the composition of a reaction mixture at equilibrium

10. The Empirical Law of Mass Action 10 bB + cC dD + eE At equilibrium (independent of starting conditions) the law of mass action is constant a B = ‘activity’ of species B K = equilibrium constant Magnitude of K tells us about the equilibrium state K >> 1, (activities of products) >> (activities of reactants) at equilibrium K << 1, (activities of products) << (activities of reactants) at equilibrium (unitless)

11. ‘Activity’ 11 a = 1 Pure Solid a = 1 Pure Liquid Ideal Solution Ideal Gas

12. Gas Phase Reactions 12 bB + cC dD + eE a = activity of reacting species For ideal gases: where P º = 1 bar ****Note**** Your textbook refers to K P as K

13. Gas Phase Reactions 13  Write the equilibrium expression for the following reaction:  CO (g) + ½ O 2(g) CO 2 (g)

14. Molarity 14  Molarity is a unit commonly used to describe the concentration of a solution  Solution – homogeneous mixture of two or more substances  Solute – one or more minor components in a solution  Solvent – the major component of a solution, medium for solute  Molarity – the number of moles per unit volume (mol/L = M)  It can also be used to describe the molar volume of a gas

15. Reactions in Solution 15 bB + cC dD + eE a = activity of reacting species For ideal solutions: where c º = 1 M

16. Reactions in Solution 16 Write the equilibrium expression for the following reaction: Cl 2(aq) + 2OH - (aq) ClO - (aq) + Cl - (aq) + H 2 O (l)

17. Pure Substances and Multiples Phases 17 a = activity of reacting species For solids:a (pure solid) = 1 For liquids:a (pure liquid) = 1 H 2 O (l) H 2 O (g) I 2(s) I 2(aq) CaCO 3(s) CaO (s) + CO 2(g)

18. Mixed Phase Reaction 18 Write the equilibrium expression for the following reaction: Zn (s) + 2H 3 O + (aq) Zn 2+ (aq) + H 2(g) + 2H 2 O (l)

19. Extent of a Reaction 19 Phosgene (COCl 2 ) is an important intermediate in the manufacture of certain plastics. It is produced by the reaction: CO (g) + Cl 2(g) COCl 2(g) a) Use the law of mass action to write the equilibrium expression for this reaction b) At 600 C, K = 0.20. Calculate the partial pressure of phosgene in equilibrium with a mixture of P CO = 0.0020 atm and P Cl2 = 0.00030 atm

20. Direction of Change in Reactions: Reaction Quotient (Q) 20 To determine the direction of a reaction, use the Reaction Quotient: Q is the same expression as K Q valid all the time K valid only at equilibrium The relationship between Q and K will determine the direction of a reaction

21. Direction of Change in Reactions: Q vs K 21 When Q = K: Reaction is at equilibrium When K > Q: Reaction proceeds forward (right) (too much reactant, not enough product) When K < Q: Reaction proceeds in reverse (left) (too little reactant, too much product K < Q K > Q K = Q time Q

22. Equilibrium Calculations 22  K vs K C  Relationships between K’s of reactions  Using K and K c  ICE box calculations

23. K vs K C 23

24. K vs K C 24

25. K vs K C 25 At 1132 ºC, for the following reaction: 2 H 2 S (g) 2 H 2(g) + S 2(g) K C = 2.26 x 10 -4. What is the value of K at this temperature?

26. Relationships among Equilibrium Expressions 26 Case 1: Compare forward and reverse reactions Case 2: Multiply a chemical reaction by a constant x Case 3: Add or subtract chemical reactions (Hess’s Law) 2 H 2(g) + O 2(g) 2 H 2 O (g) 2 H 2 O (g) 2 H 2(g) + O 2(g)

27. Relationships among Equilibrium Expressions 27 Case 1: Compare forward and reverse reactions Case 2: Multiply a chemical reaction by a constant x Case 3: Add or subtract chemical reactions (Hess’s Law) 2 H 2(g) + O 2(g) 2 H 2 O (g) H 2(g) + ½ O 2(g) H 2 O (g)

28. Relationships among Equilibrium Expressions 28 Br 2(g) + I 2(g) 2 IBr (g) Case 1: Compare forward and reverse reactions Case 2: Multiply a chemical reaction by a constant x Case 3: Add or subtract chemical reactions (Hess’s Law) 2 BrCl (g) Cl 2(g) + Br 2(g) 2 BrCl (g) + I 2(g) 2 IBr (g) + Cl 2(g) If adding reactions: 1 + 2  K 3 = K 1 x K 2 If subtracting reactions: 1 – 2  K 3 = K 1 / K 2

29. Relationships among Equilibrium Expressions 29 At 1330 K, germanium(II) oxide (GeO) and tungsten(VI) oxide (W 2 O 6 ) are both gases. The following two equilibria are established simultaneously: 2GeO (g) + W 2 O 6(g) 2 GeWO 4(g) K = 7000 GeO (g) + W 2 O 6(g) GeW 2 O 7(g) K = 38,000 Compute K for the reaction: GeO (g) + GeW 2 O 7(g) 2 GeWO 4(g) K = ?

30. Equilibrium Calculations: (R)ICE Box 30 For a Chemical Reaction: bB + cC dD + eE Initial P a Change in P a Equilibrium P a -bx -cx + dx +ex 1 atm 1 atm 0 atm 0 atm 1 - bx 1 - cx 0 + dx 0 + ex Reaction

31. Evaluating K from Reaction Data 31 ? ? 0.10 atm At 600 ºC, a gas mixture of CO and Cl 2 has initial partial pressures of P (CO) = 0.60 atm and P (Cl2) = 1.10 atm. At equilibrium, the P (COCl2) = 0.10 atm. Calculate K for this reaction: CO (g) + Cl 2(g) COCl 2(g) Initial P a Change in P a Equilibrium P a 0.60 atm 1.10 atm 0 atm -x -x + x 0.60 atm - x 1.10 atm - x 0.10 atm - 0.10 - 0.10 + 0.10 (0.60-0.10) atm (1.10-0.10) atm 0.10 atm

32. Evaluating K from Reaction Data 32 Must take into account stoichiometric coefficients: 2 As(OH) 6 3- (aq) + 6 CO 2(g) As 2 O 3(s) + 6 HCO 3 - (aq) + 3 H 2 O (l) I [A] C in [A] E [A] 2 As(OH) 6 3- (aq) + 6 CO 2(g) As 2 O 3(s) + 6 HCO 3 - (aq) + 3 H 2 O (l) 1 M 1 M - 1 M - -2x -6x +6x 1 M – 2x 1 M – 6x 1 M + 6x

33. Evaluating K from Reaction Data 33 Graphite (solid carbon) is added to a vessel that contains CO 2(g) at a pressure of 0.824 atm. The pressure rises during the reaction. The total pressure at equilibrium is 1.366 atm. Calculate the equilibrium constant: C (s) + CO 2(g) 2 CO (g)

34. Calculating Equilibrium Compositions 34 H 2(g) and I 2(g) are sealed in a flask with partial pressures: P(H 2 ) = 1.980 atm and P(I 2 ) = 1.710 atm. The sealed flask is heated to 600 K, and the gases quickly reach equilibrium. K 600 K = 92.6 H 2(g) + I 2(g) 2 HI (g) Calculate the equilibrium partial pressures of H 2, I 2 and HI at 600 K Can solve this 3 ways: 1) Solve exactly using quadratic equation 2) Solve by approximation 3) Solve by successive approximation if 2 doesn’t work ax 2 + bx + c = 0

35. Le Châtelier’s Principle 35  Adding and Removing Reagents  Changing Pressure  Effect of Temperature on Equilibrium  Effect of Temperature on K  Catalysts

36. Direction of Change in Reactions: Le Châtelier’s Principle 36 Le Châtelier’s Principle: A system in equilibrium that is subjected to a stress will react in a way that tends to counteract the stress At equilibrium, the macroscopic properties of a system remain constant When you perturb the equilibrium, the system will counteract the perturbation Changing the concentration of a reactant or product Changing the Volume or Pressure Changing the Temperature Can maximize the yield by perturbing the system

37. Le Châtelier’s Principle Adding or Removing Reagent 37 Change the concentration (or Partial Pressure) of a reagent Consider how the change in concentration affects Q Increase a Product (decrease a Reactant) Q > KReverse (to the left) Increase a Reactant (decrease a Product) Q < KForward (to the right) P X (atm) 3.0 2.0 1.0 2HI (g) H 2(g) + I 2(g)

38. Le Ch â telier’s Principle Changing Volume or Pressure 38 Specific to gas phase reactions 2 P 2(g) P 4(g) Increase V (decrease P) Increase P (decrease V) Reaction  direction with greater n gas Reaction  direction with fewer n gas

39. Le Châtelier’s Principle Changing the Temperature 39 Endothermic or Exothermic reactions Endothermic Reaction – absorbs heat Exothermic Reaction – gives off heat Increase T (“adding heat”) Reaction will proceed endothermically Trying to absorb excess heat Increase T Endothermic reaction  forward Exothermic reaction  reverse Decrease T Endothermic reaction  reverse Exothermic reaction  forward 2 NO 2(g) N 2 O 4(g) Δ H < 0  exothermic

40. Direction of Change in Reactions: Le Ch â telier’s Principle 40 3 Al 2 Cl 6(g) 2 Al 3 Cl 9(g) Increase P(Al 2 Cl 6 ) Increase Volume Decrease Temperature (assume Δ H < 0) H 2 S (g) + l 2(g) 2 HI (g) + S (s) The equilibrium constant at 110 ºC is K = 0.0023 For the following conditions, calculate Q and determine the direction of the reaction: a)P(I 2 ) = 0.461 atm, P(H 2 S) = 0.050 atm, P(HI) = 0.1 atm b) P(I 2 ) = 0.461 atm, P(H 2 S) = 0.050 atm, P(HI) = 1.0 x 10 -3 atm

41. Temperature Dependence of K 41 Δ Hº is the Enthalpy of a Reaction Δ Hº < 0 – Exothermic Δ Hº > 0 – Endothermic R = 8.31451 J/mol K Van’t Hoff Equation

42. Temperature Dependence of K 42 If Δ H < 0 (exothermic reaction), then Increase in T reduces K If Δ H > 0 (endothermic reaction), then Increase in T increases K

43. Direction of Change in Reactions: van’t Hoff Equation 43 3 H 2(g) + N 2(g) 2 NH 3(g) K(298 K) = 5.9 x 10 5 If Δ H º = -92.2 kJ/mol, calculate K at T = 600 K

44. Catalysts 44 A Catalyst is a substance that increases the rate of a chemical reaction and is not consumed in the net reaction Haber Process: N 2(g) + 3 H 2(g) 2 NH 3(g) Δ Hº = -92.4 kJ/mol Catalysts do not affect the equilibrium composition They only change the path Catalysts speed up both the forward and reverse reactions You will learn more about catalysts in 14B when you study Reaction Kinetics Fe 3 O 4