Chapter 14 Chemical Equilibrium

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Chapter 14 Chemical Equilibrium

• Equilibrium Conditions
• Example Reaction NO2 NO2
  N2O4
• (brown)
  (colorless)
• This reaction proceeds to the right at the start
• But, it never goes all the way to completion. It
  reaches a point where the forward and reverse
  reactions are going at the same rate.
• DNO2 0
• DN2O4 0
• Chemical Equilibrium The state where the
  concentrations of all reactants and products
  remain constant with time.

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• Some reactions almost reach completion before
  equilibrium is achieved. We say the equilibrium
  lies to the right.
• 2H2 O2 2H2O
• Some reactions barely get started before reaching
  equilibrium. We say the equilibrium lies to the
  left.
• 2CaO 2Ca O2
• Equilibria are Dynamic
• Even though DNO2 0, the forward and reverse
  reactions are both occurring.
• At equilibrium, there is no net change in
  concentration, even though individual molecules
  are constantly reacting.
• Reaction 1H2O 1CO
  1H2 CO2

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• At equilibrium, concentrations dont change, even
  though individual molecules do
• The equilibrium lies to the right for this
  reaction. There are more products than reactants
  at equilibrium.
• Why does Equilibrium Occur?
• Molecules react by colliding (Kinetic theory)
• The number of collisions depends on concentration
• H2O and CO are decreased by the forward
  reaction, so the forward reaction slows down.
• H2 and CO2 are increased by the forward
  reaction, so the reverse reaction speeds up.
• Eventually, the forward and reverse reactions
  reach the same rate

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• Factors that determine where the equilibrium
  position is
• Initial concentrations
• Energies of the reactants and products
• Entropy (disorder is favored)
• Reaction rates of the forward and reverse
  reactions

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• Sample Reaction N2 3H2
  2NH3
• When we mix any concentrations of these gases at
  room temperature, there are no changes in
  concentration.
• We might be at equilibrium
• The reaction rates might be very slow at these
  conditions.
• The NN bond is very strong (432 kJ/mol) and hard
  to break
• Entropy favors no forward reaction
• When a catalyst is added and the reaction is
  heated, the concentrations do change until a real
  equilibrium is reached.
• H2 disappears three
• times as fast as N2
• NH3 appears twice as
• fast as N2 disappears.

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• Equilibrium Constants
• The Law of Mass Action
• This is an empirical law discovered in 1864
• Every reaction has a constant associated with it
  telling us where the equilibrium position is.
• aA bB cC dD
• K Equilibrium Constant tells us where the
  equilibrium position is
• K gt 1 tells us the equilibrium lies to the right
• K lt 1 tells us the equilibrium lies to the left
• If we know the concentrations, we can find K from
  its equation
• K is written without units, even in cases where
  there are units left not cancelled. This is
  correct for nonideal behavior of molecules.
• Example Write K for 4NH3 7O2
  4NO2 6H2O

If K 1, halfway
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• Example N2 3H2
  2NH3
• Find K if NH3 3.1 x 10-2, N2 0.85, H2
  3.1 x 10-3
• Find K for the reverse reaction
• Find K for 0.5 N2 1.5 H2
  NH3
• Conclusions
• If K is for a forward reaction, and K is for the
  reverse K 1/K
• If you multiply a reaction by n, K Kn
• Equilibrium Position the set of equilibrium
  concentrations
• There is only one equilibrium constant for a
  given reaction
• The infinite equilibrium positions depend on
  initial concentrations

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• Example 2SO2 O2
  2SO3
• Show that K is the same for 2 different initial
  conditions
• Equilibrium Involving Pressures
• 1. Ideal gas law PV nRT

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• N2 3H2 2NH3
• We can rewrite K using concentration C
• Since C is proportional to P, the partial
  pressure, we can write a similar equilibrium
  constant KP
• Example Find KP for 2NO Cl2
  2NOCl
• PNOCl 1.2 atm, PNO 0.05 atm, PCl2 0.3
  atm
• How are KP and K related?
• KP K(RT)Dn
• Dn (sum of product coefficients) (sum of
  reactant coefficients)
• 5. Example Find K from KP in the example
  directly above

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• Heterogeneous Equilibria
• Types of Equilibria
• Homogeneous all species in the same phase (so
  far, gas phase)
• Heterogeneous species in multiple phases
• 2CO(g) C(s) CO2(g)
• Experiments show that K is not effected by how
  much pure solid or liquid is present.
• Pure solids and liquids are treated as constant
  concentrations, and therefore do not appear in
  the equilibrium expression

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• Examples 2H2O(l) 2H2(g)
  O2(g)
• 2H2O(g) 2H2(g)
  O2(g)
• Applications of Equilibrium Constants
• Equilibrium Constants (K) tell us many things
• Will the reaction go forward? (K gt 1, it will
  move forward)
• Are current conditions at equilibrium?
• What conditions will be present when equilibrium
  is achieved?
• Example A(g) B (g) C (g)
  D (g)
• Start with 9A molecules and 12B molecules
• What are the conditions at equilibrium?
• a) Guess that 5C, 5D, 4A, and 7B

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• Can we find out mathematically instead of
  guessing?
• At Equilibrium
• Try x 8
• Yes, we are at equilibrium when x 8
• This is a general way to solve an equilibrium
  problem ICE Table