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

Transcription

1 CHEMISTRY XL-14A CHEMICAL EQUILIBRIA August 20, 2011 Robert 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 Action K Gas Phase Equilibrium K P Solution Phase Equilibrium K C Extent of Reaction using K Direction of Reaction Q 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) A B

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

7 The Nature of Chemical Equilibrium 7 %Co [Co(H 2 O) 6 ] Cl - [CoCl 4 ] H 2 O A [Co(H 2 O) 6 ] 2+ [Co(H 2 O) 6 ] 2+ A 2 B C 1 C 2 [CoCl 4 ] 2- [CoCl 4 ] 2- t 1 t 2 As time progresses, rate = concentration/time slows When rate(s) = constant, reaction is at equilibrium B

8 The Nature of Chemical Equilibrium 8 Characteristics of the Equilibrium State 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 ] Cl - [CoCl 4 ] 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 : Norwegians Cato Guldberg and Peter Waage discovered the mathematical relationship that summarized the composition of a reaction mixture at equilibrium 2 SO 2 (g) + O 2 (g) 2 SO 3 (g) at K ( P D P ) d P E P ( ) e ( P B P ) b P C P ( ) c = K

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 (unitless) 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

11 Activity 11 Ideal Gas Ideal Solution c =1M Pure Liquid Pure Solid a = 1 a = 1

12 Gas Phase Reactions 12 bb + cc dd + ee a = activity of reacting species For ideal gases: where P º = 1 bar ( P D P ) d P E P P D d P E e ( ) e ( P B P ) b P C P P B b P C c = K P ( ) c = K P D d P E e P B b P C c = K P ( ) ( d +e b c ) K P = K P ( ) d +e b c ( ) ****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 ([D] c ) d ([E] c ) e ([B] c ) b [C] c ( ) c = K [D] [D] d [E] e [B] b [C] c = K C d [E] e [B] b [C] c = K ( c ) ( d +e b c ) K C = K( c ) ( d +e b c )

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) K = ([ 2+ Zn ] /1M)( P H 2 /1atm)( 1) 2 1 ([ ] /1M) 2 ( ) H 3 O + K = [ Zn 2+ ]P H 2 [ H O + ] 2 3

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 = Calculate the partial pressure of phosgene in equilibrium with a mixture of P CO = atm and P Cl2 = atm

20 20 Direction of Change in Reactions: Reaction Quotient (Q) To determine the direction of a reaction, use the Reaction Quotient: Q = P d P e D E P b P c B C [ Q = D ] d [ E] e [ B] b [ C] c 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 21 Direction of Change in Reactions: Q vs K K < Q When K < Q: Reaction proceeds in reverse (left) (too little reactant, too much product Q K = Q When Q = K: Reaction is at equilibrium K > Q When K > Q: Reaction proceeds forward (right) (too much reactant, not enough product) time

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 [D] d e eq [E] eq [B] b eq [C] eq c = K C P D,eq d b P B,eq e P E,eq c P C,eq = K P X = n X RT V [ X] = n X V R = L bar mol K P = [ X]RT X

24 K vs K C 24 d P D,eq e P E,eq b P B,eq P C,eq c = K P X = X K = P d e D P E P b B P = c C [ ]RT ([ X D]RT) d [ X E ]RT ( ) e ([ X B ]RT) b [ X C ]RT ( ) c [ K = X ] d [ D X ] e E ( RT) d +e b c = K ( [ X ] b [ B X ] c c RT) d +e b c C K = K c RT ( ) d +e b c R = L bar mol K

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 What is the value of K at this temperature? K = K c ( RT) d +e b c L bar R = mol K

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) K 1 = P 2 H 2 O P 2 H 2 P O2 2 H 2 O (g) 2 H 2(g) + O 2(g) K 1 ' = P 2 H 2 P O2 2 P H 2 O K 1 K 1 ' = P 2 H 2 O P 2 H 2 P O2 =1 P 2 2 H 2 P O2 P H 2 O

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) K 1 = P 2 H 2 O P 2 H 2 P O2 H 2(g) + ½ O 2(g) H 2 O (g) K 1 ' = P H 2 O 1/ 2 P H 2 P O2 K 1 ' = 2 P H 2 O P H 2 2 P O2 = K 1 K 1 ' = K 1 x

28 Relationships among Equilibrium Expressions 28 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) K 1 = P Cl 2 P Br2 K 2 2 = P IBr P BrCl P Br2 P I 2 Br 2(g) + I 2(g) 2 IBr (g) 2 2 BrCl (g) + I 2(g) 2 IBr (g) + Cl 2(g) K 3 = P 2 IBrP Cl2 2 P BrCl P I 2 = P Cl 2 P Br 2 P BrCl P 2 IBr 2 P I 2 P Br2 = K 1 K 2 K 3 = P 2 IBrP Cl2 2 P BrCl P I 2 If adding reactions: 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 Reaction Initial P a Change in P a Equilibrium P a bb + cc dd + ee 1 atm 1 atm 0 atm 0 atm -bx -cx + dx +ex 1 - bx 1 - cx 0 + dx 0 + ex K = P d P e D E P b P c B C K = ( dx) d ( ex) e ( 1 bx) b ( 1 cx) c

31 Evaluating K from Reaction Data 31 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 ( )? atm - atm x ( )? atm - x atm 0.10 atm K = P COCl 2 P CO P Cl2 K = ( 0.10) ( )( ) = 0.10 ( 0.50) ( 1.00) = 0.20

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) 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] 1 M 1 M - 1 M - -2x -6x +6x 1 M 2x 1 M 6x 1 M + 6x K C = [ HCO ] 6 3 [ 3 As(OH) ] 2 [ 6 CO ] 6 2 K C = [ 1+ 6x] 6 [ 1 2x] 2 [ 1 6x] 6

33 Evaluating K from Reaction Data 33 Graphite (solid carbon) is added to a vessel that contains CO 2(g) at a pressure of atm. The pressure rises during the reaction. The total pressure at equilibrium is 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 ) = atm and P(I 2 ) = 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 x = b ± b2 4ac 2a

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

36 36 Direction of Change in Reactions: Le Châtelier s Principle 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 37 Adding or Removing Reagent Change the concentration (or Partial Pressure) of a reagent Consider how the change in concentration affects Q 2HI (g) H 2(g) + I 2(g) 3.0 Increase a Product (decrease a Reactant) Q > K Q = P H 2 P I 2 2 P HI Reverse (to the left) Increase a Reactant (decrease a Product) P X (atm) Q < K Forward (to the right)

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

39 39 Le Châtelier s Principle Changing the Temperature 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) ΔH < 0 exothermic

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

41 Temperature Dependence of K 41 Van t Hoff Equation ln K 2 = ΔH K 1 R 1 1 T 2 T 1 ΔHº is the Enthalpy of a Reaction ΔHº < 0 Exothermic ΔHº > 0 Endothermic R = J/mol K

42 Temperature Dependence of K 42 ln K 2 = ΔH K 1 R 1 1 T 2 T 1 If ΔH < 0 (exothermic reaction), then Increase in T reduces K If ΔH > 0 (endothermic reaction), then Increase in T increases K

43 43 Direction of Change in Reactions: van t Hoff Equation 3 H 2(g) + N 2(g) 2 NH 3(g) K(298 K) = 5.9 x 10 5 If ΔH º = 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 2 Haber Process: Fe 3 O 4 K = P NH 3 3 P N2 P H 2 N 2(g) + 3 H 2(g) 2 NH 3(g) Hº = 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