Chem 106 Chapter 16: Principles of Reactivity: Chemical Equilibria by unm8F6

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									Chapter 16 Principles of Reactivity: Chemical Equilibria




                                                    1
I. Equilibrium
      A. Equilibrium expressions
      B. Manipulating equilibrium expressions
      C. Magnitude of K
      D. The reaction quotient
II. Applications of Equilibrium Constants
      A. Calculating equilibrium constants
      B. Calculating equilibrium concentrations
III. Le Châtelier's Principle
      A. Change in reactant or product quantity
      B. Change in pressure or volume
      C. Change in temperature
      D. Effect of catalysts



                                                  2
I. Equilibrium

      e.g., H2O + CO2    H2CO3
                                           Dynamic equilibrium

                                              Rateforward = Ratereverse

                                   [H2O]


     conc.
                                   [CO2]

                                 [H2CO3]


                        time
                                                                          3
I. Equilibrium
  General: A            B
                            [B]
                                  Equilibrium lies on side of product.
                                  Rateforward = Ratereverse
    conc.




                                  kf[A] = kr[B]
                                  kf [B]
                            [A]           1      kf > kr
                                  k r [ A]
                 time

                            [A]
                                  Equilibrium lies on side of reactant.
                                  Rateforward = Ratereverse
    conc.




                                  kf[A] = kr[B]
                                  kf [B]
                            [B]           1      kf < kr
                                  k r [ A]
                 time                                                     4
I. Equilibrium
  A. Equilibrium expressions

     For the reaction, aA + bB         cC + dD

     At equilibrium:   Rateforward = Ratereverse

                       kf[A]a[B]b = kr[C]c[D]d

                       kf        [C] c [D]d
                           Kc                     equilibrium expression
                       kr        [A]a [B] b
                                    concentration
                                    equilibrium constant
                                        (at a given temperature)



                                                                             5
I. Equilibrium
  A. Equilibrium expressions

     Also, for the reaction, aA + bB     cC + dD
                                 c d
                                pC p D
                            Kp  a b
                                pA pB
                                  pressure

     e.g., 2N2O5(g)       4 NO2(g) + O2(g)

         What are Kc and Kp?




                                                   6
I. Equilibrium
  A. Equilibrium expressions

  Heterogeneous equilibria:

      1. insoluble solids are not included in Kc or Kp

           e.g., AgCl(s) + 2NH3(aq)               Ag(NH3)2+(aq) + Cl–(aq)

                        [Ag(NH 3 )  ][Cl  ]
               K c'                 2
                        [AgCl ( s )][ NH 3 ]2
                                    constant ( density)

                                         [Ag(NH 3 )  ][Cl  ]
             K c' [AgCl ( s )]    Kc             2
                                             [ NH 3 ]2


                                                                            7
I. Equilibrium
  A. Equilibrium expressions

  Heterogeneous equilibria:

      2. solids, liquids are not included in Kp

           e.g., H2O(g) + SO3(g)             H2SO4(l)
                             pH 2SO 4
                    '
                   Kp                            constant (pvap at T)
                            pH 2O pSO 3

                           '
                          Kp                  1
                                   Kp 
                        pH 2SO 4          pH 2O pSO 3



                                                                         8
I. Equilibrium
  A. Equilibrium expressions
      e.g., Write the equilibrium expression for each of the following reactions.

      Cu(OH)2(s)         Cu2+(aq) + 2 OH–(aq)      (Kc)




      Ni(s) + 2Ag+(aq)        Ni2+(aq) + 2 Ag(s)     (Kc)




      2 NH3(g) + CO2(g) + H2O(g)          (NH4)2CO3(s)       (Kp)



                                                                            9
I. Equilibrium
  B. Manipulating equilibrium expressions
      1. reversing the equation
                                           [HCl]2
         H2 + Cl2        2HCl        Kc =
                                          [H2][Cl2]

                                             [H2][Cl2]
         2HCl          H2 + Cl2      Kc´ =
                                              [HCl]2

                                     1
                            Kc´ =
                                     Kc

         When reversing a chemical equation, take the inverse
         of the equilibrium constant.




                                                                10
I. Equilibrium
  B. Manipulating equilibrium expressions
      2. multiplying equations by some factor
                                             [HCl]4
         2H2 + 2Cl2        4HCl     Kc´´ =
                                           [H2]2[Cl2]2

                            Kc´´ = (Kc)2


                                                [HCl]
         ½H2 + ½Cl2         HCl     Kc´´´ =
                                              [H2]½[Cl2]½

                            Kc´´´ = (Kc)½

         When multiplying a chemical equation by some factor,
         raise the equilibrium constant to that power.

                                                                11
I. Equilibrium
  B. Manipulating equilibrium expressions
      3. adding equations

                                                       [NO][SO3]
        SO2 + NO2           NO + SO3             Kc1 = [SO ][NO ]
                                                           2     2
                                                         [H2SO4]
        SO3 + H2O           H2SO4                Kc2 =
                                                       [SO3][H2O]

                                                           [NO][H2SO4]
        SO2 + NO2 + H2O             NO + H2SO4   Kc3 =
                                                         [SO2][NO2][H2O]

                            [NO][SO3]    [H2SO4]    [NO][H2SO4]
        Kc3 = Kc1 · Kc2 =             ·          =
                            [SO2][NO2] [SO3][H2O] [SO2][NO2][H2O]

      When adding equations, multiply the equilibrium constants.


                                                                       12
I. Equilibrium
   B. Manipulating equilibrium expressions
 e.g., Given that for the reaction, N2(g) + 3H2(g)       2NH3(g), Kc = 3.5 x 108,
 what is Kc for the reaction NH3(g)        1/2 N2(g) + 3/2 H2(g)?




 e.g., Given the following:
      H2O(g) + CO(g)         H2(g) + CO2(g) Kp = 1.6
      FeO(s) + CO(g)         Fe(s) + CO2(g) Kp = 0.67
 What is Kp for the reaction Fe(s) + H2O(g)    FeO(s) + H2(g)




                                                                           13
I. Equilibrium
  C. Magnitude of K

      2SO2 + O2         2SO3     K = 3 x 1024

        Equilibrium lies very far to the right; reaction “goes to completion.”


      2HF          H 2 + F2      K = 1 x 10–13

        Equilibrium lies very far to the left; reaction “doesn’t go.”

                                 OH
      H2O + H2C=O              H C OH            K1
                                 H
            “Equilibrium mixture” of reactant and product.

                                                                          14
I. Equilibrium
  D. The reaction quotient, Q

     For aA + bB         cC + dD

          [C]c [D]d
     Qc                 not necessarily at equilibrium
          [A]a [B]b


      If Q > K,   reaction shifts left to achieve equilibrium

      If Q < K,   reaction shifts right to achieve equilibrium

      If Q = K,   reaction is at equilibrium




                                                                 15
I. Equilibrium
  D. The reaction quotient, Q

  e.g., N2 + O2       2NO      Kc = 4.0 x 10–4 at 2000 K

  If [N2] = 0.50 M, [O2] = 0.25 M, and [NO] = 4.2 x 10–3 M,

  Is the system at equilibrium? Which way will it shift?




                                                              16
II. Applications of Equilibrium Constants
  A. Calculating equilibrium constants

   e.g., 2SO2 + O2      2SO3
  1.00 mol of SO2 and 1.00 mol of O2 were placed in a 1.00-L flask at 1000ºC.
  When equilibrium was attained, it was found that the concentration of SO3
  was 0.925 M. What is Kc?

              [SO2]       [O2]           [SO3]
       [ ]i
      D[ ]
      [ ]eq




                                                                        17
II. Applications of Equilibrium Constants
  A. Calculating equilibrium constants

   e.g., CO + Cl2       COCl2
  A mixture of 0.102 M CO and 0.00609 M Cl2 was placed in a flask. When
  the reaction had come to equilibrium at 325 °C, the concentration of Cl2 was
  found to be 0.00301 M. What is Kc for the reaction?




                                                                         18
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  1. SO2 + NO2        NO + SO3         Kc = 85.0
     0.50 mol of SO2 and 0.50 mol of NO2 are placed in a 10.0-L flask and
     allowed to come to equilibrium. What are the equilibrium concentrations
     of all species?

              [SO2]    [NO2]      [NO]    [SO3]
       [ ]i
      D[ ]
      [ ]eq




                                                                       19
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  2. SO2 + NO2        NO + SO3         Kc = 85.0
     Same as problem 1, but this time start with products: 0.050 M NO and
     0.050 M SO3. What are the equilibrium concentrations of all species?


              [SO2] [NO2] [NO]      [SO3]
       [ ]i
      D[ ]
      [ ]eq




                                                                       20
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  3. COCl2       CO + Cl2         Kc = 2.2 x 10-10
     When 0.20 mol of COCl2 is placed in a 10.0-L flask, what are the
     equilibrium concentrations of all species?
             [COCl2]   [CO]    [Cl2]
      [ ]i
      D[ ]




                                                                        21
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  4. 2CO2        2CO + O2             Kc = 6.4 x 10-7 at 2000ºC
     Starting with 0.010 M CO2, what are the [ ]eq’s of all species?
             [CO2]    [CO]     [O2]
      [ ]i
      D[ ]




                                                                       22
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  5. CO + H2    H2CO      Kc = 4.6 x 109
    0.15 0.20 M
            [CO]   [H2]    [H2CO]
     [ ]i
     D[ ]




                                              23
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  5. (cont.) CO + H2       H2CO   Kc = 4.6 x 109
            0.15 0.20 M

             [CO]   [H2]     [H2CO]
      [ ]i
      [ ]f
      D[ ]




                                                   24
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  6. 2SO2 + O2   2SO3         Kp = 1.1 x 1012
     1.0 1.0 atm
     Carry reaction to completion:

            pSO2     pO2      pSO3
      pi
      pf
      Dp




                                                25
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  e.g., For the decomposition of gaseous water at 500ºC,
            2H2O(g)        2H2(g) + O2(g), Kc = 6.0 x 10–28.
       If 0.20 mol of H2O is placed in a 5.0-L container and allowed to come
       to equilibrium at 500ºC, what will be the concentrations of all species?




                                                                           26
II. Applications of Equilibrium Constants
  B. Calculating equilibrium concentrations
  e.g., For the reaction, CO(g) + Cl2(g)      COCl2(g), Kc = 4.6 x 109. If
  0.15 mol of CO and 0.30 mol of Cl2 are placed in a 1.0-L flask and allowed
  to react, what would be the concentration of each species at equilibrium?




                                                                        27
III. Le Châtelier’s Principle

    “Any change imposed on a system at equilibrium will cause the system
    to shift in such a way to re-establish equilibrium.”




                     Reactants             Products




                                                                      28
III. Le Châtelier’s Principle
   A. Change in quantity of reactant or product

     A        B at equilibrium

                              increase [B]
               [B]
         K=             reaction shifts to left
               [A]

      A reaction will shift: away from a substance added,
                              toward a substance removed.

     e.g., 3H2(g) + N2(g)              2NH3(g)
                   2
                  pNH3                   Increase p H 2 : reaction shifts ______
          Kp     3
                 pH 2 p N 2              Increase p NH3 : reaction shifts ______
                                         Decrease p N 2 : reaction shifts ______

                                                                               29
III. Le Châtelier’s Principle
   A. Change in quantity of reactant or product
                                                 CH3
   e.g., CH3   CH2    CH2     CH3         CH3    CH CH3        Kc = 2.5
                butane (B)                   isobutane (I)
   At equilibrium, [B] = 0.20 M and [I] = 0.50 M. If 0.10 mol of butane is added
   and the system returns to equilibrium,
       a. Which way will the reaction shift to re-establish equilibrium?
       b. What will be the new [ ]eq’s?
       c. Is there be more or less butane after the reaction reaches equilibrium?




                                                                           30
III. Le Châtelier’s Principle
   B. Change in pressure or volume
                                       pC
      A(g) + B(g)       C(g)     Kp = p p
                                       A B


      Decrease V:     reaction shifts to side with
      (increase PT)    fewer moles of gas
        Increase V:    reaction shifts to side with
       (decrease PT)   more moles of gas


       e.g., 3H2(g) + N2(g)        2NH3(g)

         Increase V: reaction shifts ________
         Increase PT: reaction shifts ________


                                                      31
III. Le Châtelier’s Principle
   C. Change in temperature

      2NO2(g)       N2O4(g) + heat       DH = -57.2 kJ/mol
                           “product”

        DH < 0: increasing T shifts equilibrium left (K gets smaller)
                 decreasing T shifts equilibrium right (K gets larger)

          DH > 0: opposite effect


       e.g., 2NOBr(g)       2NO(g) + Br2(g)     DH = +16.1 kJ/mol

         Increase T: reaction shifts ________
         Decrease T: reaction shifts ________



                                                                         32
III. Le Châtelier’s Principle
   D. Effect of catalysts
       No effect on position of equilibrium
           (neither a reactant nor a product of the reaction)
        - increases rate at which equilibrium is established

   e.g., 2N2O(g) + O2(g)       4 NO (g)           DH = 199 kJ/mol

   Will the amount of NO at equilibrium be greater, less, or the same if we:
       •add N2O?       __________
       •remove O2?      __________
       •add NO?         __________
       •increase V?     __________
       •raise T?        __________
       •add a catalyst? __________
                                                                           33
III. Le Châtelier’s Principle

    e.g., CaCO3(s)     CaO(s) + CO2(g)         DH = 278 kJ/mol

    How will the amount of CO2 change if we:
       •add CaCO3(s)? __________
        •remove CO2(g)? __________
                                                     K p  pCO 2
        •decrease V?     __________
        •decrease T?     __________




                                                                   34

								
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