Transcript Chemical Equilibrium - CNYRIC

AP Chapter 15
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Chemical Equilibrium occurs when opposing
reactions are proceeding at equal rates.
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It results in the formation of an equilibrium mixture
of the reactants and products of the reaction.
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The composition mixture does not change with
time.
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N2(g) + 3H2(g) ↔ 2NH3(g)
This reaction involves the presence of a catalyst, a
pressure of several hundred atmospheres and a
temperature of several hundred degrees Celsius.
This equilibrium mixture can be reached
regardless of where one starts.
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The relationship between the concentrations of the
reactants and the products of a system in
equilibrium is given by the law of mass action.
aA + bB ↔ dD + eE
Kc = equilibrium constant
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Kc =
[D]d[E]e ← products
[A]a[B]b ← reactants
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N2O4(g) ↔ 2NO2(g)
Kc =
[NO2]2
[N2O4]
=
[0.0172]2
[0.00140]
= 0.211
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When the equilibrium system consists of gases, it
is convenient to express the concentrations of
reactants and products in terms of gas pressures:
Kp =
(PD)d(PE)e
(PA)a(PB)b
Kc and Kp are related by: Kp = Kc(RT)Δn
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CO(g) + Cl2(g) ↔ COCl2(g)
Kc =
[COCl2]
= 4.56 x 109
[CO][Cl2]
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For the equilibrium constant to be so large, the
numerator must be much larger than the
denominator. Therefore, the equilibrium concentration
of COCl2 must be greater than that of CO or Cl2.
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A large value for the equilibrium constant indicates
that the mixture contains more products than
reactants and therefore lies towards the product
side of the equation.
A small value for the equilibrium constant means
the mixture contains less products than reactants
and therefore lies toward the reactant side.
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Equilibria for which all substances are in the same
phase are called homogeneous equilibria.
Equilibria in which 2 or more phases are present
are called heterogeneous equilibria.
The concentrations of pure solids and liquids are
left out of the equilibrium constant expression for a
heterogeneous equation.
The equilibrium pressure of CO2 (g) is the same in both
bell jars, at the same temperature. The equilibrium
constant expression is Kp = PCO2.
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If the concentrations of all species in an
equilibrium are known, the equilibrium-constant
expression can be used to calculate the value of
the equilibrium constant.
The changes in the concentrations of reactants
and products in the process of achieving
equilibrium are governed by the stoichiometry of
the equation.
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The reaction quotient, Q, is found by substituting
reactant and product concentrations or partial
pressures at any point during a reaction into the
equilibrium-constant expression.
If the system is at equilibrium, Q = K.
Q ≠ K, the system is not at equilibrium.
Q < K, the reaction will move toward equilibrium by
forming more products (it moves from left to right.)
Q > K, the reaction will proceed from right to left.
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Knowing the value of K makes it possible to
calculate the equilibrium amounts of the reactants
and products, often by solving an equation where
the unknown is the change in a partial pressure or
concentration.
Predicting the
direction of a
reaction by
comparing Q
and K.
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LeChatlier’s principle states:
◦ If a system at equilibrium is disturbed by a
change in temperature, pressure or
concentration of one of the components,
the system will shift its equilibrium position
to counteract the effect of the disturbance.
The effect of
adding H2 to an
equilibrium
mixture of N2,
H2 and NH3.
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If a chemical system is at equilibrium and the
concentration of a substance is increased (either
product or reactant), the system responds by
consuming some of the substance.
If some of the concentration is decreased, the
system will respond by producing some of the
substance.
 N2 + 3H2 ↔ 2NH3
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At constant temperature, reducing the volume of a
gaseous equilibrium mixture causes the system to
shift in the direction that reduces the number of
moles of gas.
The system will always favor the side with fewer
moles of a gas in order to re-equilibrate.
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When the temperature of a system in equilibrium
is increased, the system reacts as though a
reactant were added to an endothermic reaction or
a product was added to an exothermic reaction.
It will shift in the direction to consume the excess
reactant or product, which is heat.
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Endothermic:
reactants + heat ↔ products
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Exothermic:
reactants ↔ products + heat
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Adding a catalyst increases the rate at which
equilibrium is achieved, but it does not change the
composition of the equilibrium mixture.