the double arrow implies the process is dynamic. consider
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The Concept of Equilibrium. The double arrow implies the process is dynamic. Consider Forward reaction: A B Rate = k f [A] Reverse reaction: B A Rate = k r [B] At equilibrium k f [A] = k r [B]. The Concept of Equilibrium. For an equilibrium we write - PowerPoint PPT PresentationTRANSCRIPT
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• The double arrow implies the process is dynamic.• Consider
Forward reaction: A B Rate = kf[A]
Reverse reaction: B A Rate = kr[B]
• At equilibrium kf[A] = kr[B].
The Concept of EquilibriumThe Concept of Equilibrium
N2O4(g) 2NO2(g)
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• For an equilibrium we write • As the reaction progresses
– [A] decreases to a constant,
– [B] increases from zero to a constant.
– When [A] and [B] are constant, equilibrium is achieved.
• Alternatively:– kf[A] decreases to a constant,
– kr[B] increases from zero to a constant.
– When kf[A] = kr[B] equilibrium is achieved.
The Concept of EquilibriumThe Concept of Equilibrium
A B
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The Equilibrium ConstantThe Equilibrium Constant
• For a general reaction in the gas phase
the equilibrium constant expression is
where Keq is the equilibrium constant.
aA + bB cC + dD
ba
dc
eqPP
PPK
BA
DC
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The Equilibrium ConstantThe Equilibrium Constant
• For a general reaction
the equilibrium constant expression for everything in solution is
where Keq is the equilibrium constant.
aA + bB cC + dD
ba
dc
eqKBA
DC
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The Equilibrium ConstantThe Equilibrium Constant
• Keq is based on the molarities of reactants and products at equilibrium.
• We generally omit the units of the equilibrium constant.• Note that the equilibrium constant expression has
products over reactants.• The same equilibrium is established no matter how the
reaction is begun.
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The Equilibrium ConstantThe Equilibrium ConstantThe Magnitude of Equilibrium
Constants• The equilibrium constant, K, is the ratio of products to
reactants.• Therefore, the larger K the more products are present at
equilibrium.• Conversely, the smaller K the more reactants are present
at equilibrium.• If K >> 1, then products dominate at equilibrium and
equilibrium lies to the right.
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The Equilibrium ConstantThe Equilibrium ConstantThe Magnitude of Equilibrium
Constants• If K << 1, then reactants dominate at equilibrium and the
equilibrium lies to the left.
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The Equilibrium ConstantThe Equilibrium ConstantThe Direction of the Chemical
Equation and Keq
• An equilibrium can be approached from any direction.• Example:
• has
N2O4(g) 2NO2(g)
46.6
42
2
ON
2NO
P
PKeq
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The Direction of the Chemical Equation and Keq
• In the reverse direction:
The Equilibrium ConstantThe Equilibrium Constant
46.61
155.02NO
ON
2
42 P
PKeq
2NO2(g) N2O4(g)
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Other Ways to Manipulate Chemical Equations and Keq Values
• The reaction
has
which is the square of the equilibrium constant for
The Equilibrium ConstantThe Equilibrium Constant
2N2O4(g) 4NO2(g)
2ON
4NO
42
2
P
PKeq
N2O4(g) 2NO2(g)
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Other Ways to Manipulate Chemical Equations and Keq Values
• Equilibrium constant for the reverse direction is the inverse of that for the forward direction.
• When a reaction is multiplied by a number, the equilibrium constant is raised to that power.
• The equilibrium constant for a reaction which is the sum of other reactions is the product of the equilibrium constants for the individual reactions.
The Equilibrium ConstantThe Equilibrium Constant
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• We ignore the concentrations of pure liquids and pure solids in equilibrium constant expressions.
Heterogeneous EquilibriaHeterogeneous Equilibria
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Predicting the Direction of Reaction• We define Q, the reaction quotient, for a general reaction
as
• Q = K only at equilibrium.
Applications of Applications of Equilibrium ConstantsEquilibrium Constants
aA + bB cC + dD
ba
dc
PP
PPQ
BA
DC
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• Proceed as follows:– Tabulate initial and equilibrium concentrations (or partial
pressures) given.
– If an initial and equilibrium concentration is given for a species, calculate the change in concentration.
– Use stoichiometry on the change in concentration line only to calculate the changes in concentration of all species.
– Deduce the equilibrium concentrations of all species.
• Usually, the initial concentration of products is zero. (This is not always the case.)
Calculating Equilibrium Calculating Equilibrium ConstantsConstants
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Predicting the Direction of Reaction• If Q > K then the reverse reaction must occur to reach
equilibrium (i.e., products are consumed, reactants are formed, the numerator in the equilibrium constant expression decreases and Q decreases until it equals K).
• If Q < K then the forward reaction must occur to reach equilibrium.
Applications of Applications of Equilibrium ConstantsEquilibrium Constants
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• Proceed as follows:– Tabulate initial and equilibrium concentrations (or partial
pressures) given.
– If an initial and equilibrium concentration is given for a species, calculate the change in concentration.
– Use stoichiometry on the change in concentration line only to calculate the changes in concentration of all species.
– Deduce the equilibrium concentrations of all species.
• Usually, the initial concentration of products is zero. (This is not always the case.)
Calculating Equilibrium Calculating Equilibrium ConstantsConstants
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Calculating Equilibrium Constants• The same steps used to calculate equilibrium constants
are used.• Generally, we do not have a number for the change in
concentration line.• Therefore, we need to assume that x mol/L of a species is
produced (or used).• The equilibrium concentrations are given as algebraic
expressions.
Applications of Applications of Equilibrium ConstantsEquilibrium Constants
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• H2(g) + I2(g) 2HI(g) Keq = 51
• Predict in which direction the reaction will proceed to reach equilibrium at 448 oC if we start with 0.02 mol HI, 0.01 mol H2 and 0.03 mol I2 in a 2.00 L container.
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Use of Approximation
• Approximations can be used if the value of K is so small that the terms (-x) or (+x) will barely have an effect.
• Validity of approximation: < 5%
• ______[x]______ = < 5%[initial concentration]
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Sample Problem
• 2NOCl(g) 2NO(g) + Cl2(g) K = 1.6 x 10-5
• 1.0 mole NOCl is placed in a 2.0-L flask. What are the equilibrium concentrations?
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• Le Châtelier’s Principle: if a system at equilibrium is disturbed, the system will move in such a way as to counteract the disturbance.
Le ChLe Châtelier’s Principleâtelier’s Principle
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Change in Reactant or Product Concentrations
• Adding a reactant or product shifts the equilibrium away from the increase.
• Removing a reactant or product shifts the equilibrium towards the decrease.
• To optimize the amount of product at equilibrium, we need to flood the reaction vessel with reactant and continuously remove product (Le Châtelier).
• We illustrate the concept with the industrial preparation of ammonia.
Le ChLe Châtelier’s Principleâtelier’s Principle
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Effects of Volume and Pressure Changes
• Le Châtelier’s Principle: if pressure is increased the system will shift to counteract the increase.
• That is, the system shifts to remove gases and decrease pressure.
• An increase in pressure favors the direction that has fewer moles of gas.
• In a reaction with the same number of product and reactant moles of gas, pressure has no effect.
Le ChLe Châtelier’s Principleâtelier’s Principle
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Effects of Volume and Pressure Changes
• An increase in pressure (by decreasing the volume) favors the formation of colorless N2O4.
• The instant the pressure increases, the system is not at equilibrium and the concentration of both gases has increased.
• The system moves to reduce the number moles of gas (i.e. the forward reaction is favored).
Le ChLe Châtelier’s Principleâtelier’s Principle
N2O4(g) 2NO2(g)
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Effect of Temperature Changes• Adding heat (i.e. heating the vessel) favors away from
the increase:– if H > 0, adding heat favors the forward reaction,
– if H < 0, adding heat favors the reverse reaction.
• Removing heat (i.e. cooling the vessel), favors towards the decrease:– if H > 0, cooling favors the reverse reaction,
– if H < 0, cooling favors the forward reaction.
Le ChLe Châtelier’s Principleâtelier’s Principle
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The Effect of Catalysis• A catalyst lowers the activation energy barrier for the
reaction.• Therefore, a catalyst will decrease the time taken to reach
equilibrium.• A catalyst does not effect the composition of the
equilibrium mixture.
Le ChLe Châtelier’s Principleâtelier’s Principle