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Science Dept.

Chemistry CHE 100 Notes
T. Chemical Equilibria

1. Definitions
   1. Reaction Rates: Rate or speed at which products are produced

   2. Reversible Reactions

   3. Chemical Equilibrium

2. Reaction Rates
   1. Guldberg and Waage - The Law of Mass Action

   2. The Law of Mass Action
      1. Rate of a chemical reaction is proportional to the "Active Masses" of the Reactants (molar concentration for solutions or pressure for gases)

   3. For the general Equation: aA + bB <====> cC + dD
      1. Forward Reaction Rate is proportional to [A]^a[B]^b

      2. Forward Reaction Rate is = K_f[A]^a[B]^b

      3. Reverse Reaction Rate is proportional to [C]^c[D]^d

      4. Reverse Reaction Rate is = K_r[C]^c[D]^d

      5. Forward Reaction Rate = Reverse Reaction Rate

      6. K_f[A]^a[B]^b = K_r[C]^c[D]^d

   4. Using the Example: H_2(g) + I_2(g) <====> 2HI_(g)
      1. K_f[H₂][I₂] = K_r[HI]²

      2. Solving for a general K = K_f/K_r = [HI]²/[H₂][I₂] = K = 45.9

      3. N.B.: K depends on Temperature

      4. Would there be more products than reactants?

      5. How could you tell?

   5. Using the Example: N_2(g) + 3H_2(g) <====> 2NH_3(g) ... the Haber Process
      1. K = [NH₃]²/ [N₂][H₂]³ = K = 2.37 x 10^-3

      2. N.B.: K depends on Temperature

      3. Would there be more reactants than products?

      4. How could you tell?

      5. How could more products be generated?

3. The General Law of Mass Action
   1. For the general Equation: aA + bB <====> cC + dD

   2. [C]^c[D]^d/[A]^a[B]^b = K = Equilibrium Constant

   3. Some Helpful Rules
      1. Solids have a constant composition at a given Temperature and are not considered in the Equilibrium Expression because they are considered in the value for the K constant.

      2. Water has the value = [1]

4. Le Chatelier's Principle
   1. Definition: If there is a change in an equilibrium, (in concentration, temperature or pressure) the system will change to a new equilibrium, if possible, in a direction that will tend to restore the original conditions
      1. Concentration

      2. Temperature

      3. Pressure

5. Terms
   1. Exothermic

   2. Endothermic

6. Use of the Equilibrium Constant as a Dissociation Constant
   1. Weak Electrolytes such as HC₂H₃O₂
      1. HC₂H₃O₂ <====> H¹⁺ + C₂H₃O₂^1-

      2. K = [H¹⁺] [C₂H₃O₂^1-]/ [HC₂H₃O₂] = 1.76 x 10^-5

      3. What is the [H¹⁺] when the HC₂H₃O₂ = 1M?

      4. Helpful Hints:
         1. Let [H¹⁺] = [C₂H₃O₂^1-] = X

         2. Let [HC₂H₃O₂] = 1-X

         3. Since X will be so small (see K constant) 1-X is almost equal to 1, then let [HC₂H₃O₂] = 1

   2. Water and Its Dissociation
      1. H₂O <====> H¹⁺ + OH^1-

      2. K = [H¹⁺][OH^1-]/[H₂O] = 1 x 10^-14

      3. Letting [H₂O] = [1]

      4. [H¹⁺][OH^1-] = 1 x 10^-14

7. Use of the Equilibrium Constant as a Solubility Product
   1. In General
      1. A_aB_b <====> aA^x+ + bB^y-

      2. K_sp = [A^x+]^a [B^y-]^b

   2. Barium Sulfate as an example
      1. BaSO₄ <====> Ba²⁺ + SO₄^2-

      2. K_sp = [Ba²⁺] [SO₄^2-]

      3. If the K_sp is extremely small, would Barium Sulfate be soluble or insoluble?