Chemical Equilibrium - Scf

“Chemical
Equilibrium”
Chapter 14
Hemoglobin
• Hemoglobin is a protein (Hb) found in
red blood cells that reacts with O 2
– enhances the amount of O 2 that can be
carried through the blood stream
Hb + O 2 ⇔ HbO 2
– the Hb represents the entire protein – it is
not a chemical formula
– the ⇔ represents that the reaction is in
dynamic equilibrium
Hemoglobin
• The concentrations of Hb, O 2 , and HbO 2 are
all interdependent
• The relative amounts of Hb, O 2 , and HbO 2 at
equilibrium are related to a constant called
the equilibrium constant, K
• Changing the concentration of any one of
these necessitates changing the other
concentrations to reestablish equilibrium
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In the lungs, with high
concentration of O 2 ,
the equilibrium shifts
to combine the Hb and
O 2 together to make
more HbO 2
Hemoglobin
In the cells, with low
concentration of O 2 ,
the equilibrium shifts
to break down the
HbO 2 and increase the
amount of free O 2
Fetal Hemoglobin
Fetal hemoglobin’s
equilibrium constant is
larger than adult
hemoglobin
Because fetal hemoglobin
is more efficient at
binding O 2, O 2 is
transferred to the fetal
hemoglobin from the
mother’s hemoglobin in
the placenta
Reactions
• When a reaction starts, the reactants are consumed and
products are made
– In the forward reaction = reactants → products
– reactant concentrations decrease and the product
concentrations increase as reactant concentration
decreases, the forward reaction rate decreases
• Eventually, the products can react to reform some of the
reactants
– In the reverse reaction = products → reactants
– assuming the products are not allowed to escape as product
concentration increases, the reverse reaction rate increases
• Processes that proceed in both the forward and reverse
direction are said to be reversible
– reactants ⇔ products
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Dynamic Equilibrium
• As the forward reaction slows and the
reverse reaction accelerates, eventually
they reach the same rate
• Dynamic equilibrium is the condition
where the rates of the forward and reverse
reactions are equal
• Once the reaction reaches equilibrium, the
concentrations of all the chemicals remain
constant
Equilibrium ≠ Equal
• The rates of the forward and reverse
reactions are equal at equilibrium
• That does not mean the concentrations of
reactants and products are equal
– Some reactions reach equilibrium only after almost
all the reactant molecules are consumed – we say
the position of equilibrium favors the products
– Other reactions reach equilibrium when only a
small percentage of the reactant molecules are
consumed – we say the position of equilibrium
favors the reactants
An Analogy: Population
Changes
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Equilibrium Constant
• Even though the concentrations of reactants and
products are not equal at equilibrium, there is a
relationship between them
• the relationship between the chemical equation
and the concentrations of reactants and products
can be seen in the Equilibrium Expression
• Consider: aA + bB ⇔ cC + dD
– the lowercase letters represent the coefficients of the
balanced chemical equation
• K is called the equilibrium constant is equal to the
product concentration raised to their coefficient
over the reactants raised to their coeffcients K is
unitless
Writing Equilibrium Constant
Expressions
• For the reaction
• aA (aq) + bB (aq) ⇔ cC (aq) + dD (aq) the
equilibrium constant expression is
• For the reaction
2 N 2O 5 ⇔ 4 NO 2 + O 2 the
equilibrium constant
expression is:
What Does the Value of K eq
Imply?
• When the value of K eq >> 1, we know that
when the reaction reaches equilibrium there
will be many more product molecules present
than reactant molecules
• When the value of K eq

Large K
Small K
Equilibrium Constants for
Reactions Involving Gases
• The concentration of a gas in a mixture is
proportional to its partial pressure
• Therefore, the equilibrium constant can be expressed
as the ratio of the partial pressures of the gases
• For aA(g) + bB(g) ⇔ cC(g) + dD(g) the equilibrium
constant expressions are
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K c and K p
• In calculating K p, the partial pressures are always in
atm
• The values of K p and K c are not necessarily the same
– because of the difference in units
– K p = K c when Δn = 0
• The relationship between them is: K p = K c (RT) Δn
Example
Heterogeneous Equilibria
• Pure solids and pure liquids are materials
whose concentration doesn’t change during
the course of a reaction
• Because their concentration doesn’t change,
solids and liquids are not included in the
equilibrium constant expression
– the reaction aA(s) + bB(aq) ⇔ cC(l) + dD(aq)
– the equilibrium constant expression is:
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Heterogeneous Equilibria
Calculating Equilibrium Constants from
Measured Equilibrium Concentrations
• The most direct way of finding the equilibrium
constant is to measure the amounts of
reactants and products in a mixture at
equilibrium
• The equilibrium mixture may have different
amounts of reactants and products, but the
value of the equilibrium constant will always
be the same
– as long as the temperature is kept constant
Calculating Equilibrium
Concentrations
• Stoichiometry can be used to determine the
equilibrium concentrations of all reactants and
products if you know initial concentrations
and one equilibrium concentration
– suppose you have a reaction 2 A (aq) + B (aq) ⇔ 4 C (aq) with
initial concentrations [A] = 1.00 M, [B] = 1.00 M, and [C] = 0.
You then measure the equilibrium concentration of C as [C]
= 0.50 M.
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Example
The Reaction Quotient
• If a reaction mixture, containing both reactants and
products, is not at equilibrium; how can we determine
which direction it will proceed?
• The answer is to compare the current concentration
ratios to the equilibrium constant
• The concentration ratio of the products (raised to the
power of their coefficients) to the reactants (raised to
the power of their coefficients) is called the reaction
quotient, Q (a snapshot of the reaction)
The Reaction Quotient:
Predicting the Direction of Change
• Q > K, the reaction will proceed fastest in the reverse
direction
• Q < K, the reaction will proceed fastest in the forward
direction
• Q = K, the reaction is at equilibrium
• Note:
– if a reaction mixture contains just reactants, Q = 0, and the
reaction will proceed in the forward direction
– if a reaction mixture contains just products, Q = ∞, and the
reaction will proceed in the reverse direction
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Example
• A sample of PCl 5(g) is placed in a 0.500 L container
and heated to 160°C. The PCl 5 is decomposed into
PCl 3(g) and Cl 2(g). At equilibrium, 0.203 moles of
PCl 3 and Cl 2 are formed. Determine the equilibrium
concentration of PCl 5 if K c = 0.0635
Finding Equilibrium Concentrations When
Given the Equilibrium Constant and Initial
Concentrations or Pressures
• First decide which direction the reaction will
proceed
– compare Q to K
• Define the changes of all materials in terms of x
• Solve for x
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Example
• For the reaction I 2 (g) ⇔ 2 I(g) the value of
K c = 3.76 x 10 -5 at 1000 K. If 1.00 moles of I 2 is placed into a
2.00 L flask and heated, what will be the equilibrium
concentrations of [I 2 ] and [I]?
(Hint: you will need to use the quadratic formula to solve for x)
Approximations to Simplify the Math
• When the equilibrium constant is very small,
the position of equilibrium favors the
reactants
• For relatively large initial concentrations of
reactants, the reactant concentration will not
change significantly when it reaches
equilibrium
– the [X] equilibrium = ([X] initial − ax) ≈ [X] initial
– assuming the reaction is proceeding forward
Checking the Approximation and
Refining as Necessary
• We can check our approximation
afterwards by comparing the
approximate value of x to the initial
concentration
• If the approximate value of x is less
than 5% of the initial concentration, the
approximation is valid
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Example
• For the reaction 2 H 2 S(g) ⇔ 2 H 2 (g) + S 2 (g) @ 800°C, K c = 1.67
x 10 -7 . If a 0.500 L flask initially containing 1.25 x 10 -4 mol H 2S
is heated to 800°C, find the equilibrium concentrations.
Disturbing and Re-establishing
Equilibrium
• Once a reaction is at equilibrium, the
concentrations of all the reactants and
products remain the same
• However if the conditions are changed, the
concentrations of all the chemicals will
change until equilibrium is re-established
• The new concentrations will be different, but
the equilibrium constant will be the same
– unless you change the temperature
Le Châtelier’s Principle
• Le Châtelier's Principle guides us in
predicting the effect various changes in
conditions have on the position of equilibrium
• It says that if a system at equilibrium is
disturbed, the position of equilibrium will shift
to minimize the disturbance
– disturbances all involve making the system open
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Le Châtelier’s Principle
• Think of this principle in term of a
push/pull stress.
• Whatever you do to a reaction can have
one of four possible outcomes:
– Shift product
– Shift reactant
– No Shift
– Cannot be determine
Le Châtelier’s Principle
(Concentration Stress)
Type of Stress
Increase reactants or
decrease products
Decrease reactants or
increase products
Result
Shift to the product
side
Shift to the reactant
side
Le Châtelier’s Principle
(Pressure/Volume Stress)
Type of Stress
Decrease pressure or
increase volume
Increase pressure or
decrease volume
Result
Shift to the side with
more gas particles
Shift to the side with
fewer gas particles
Note: This is only for reactions involving gasses (if no
gas or equal moles of gas then no shift occurs)
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Type of Stress
Le Châtelier’s Principle
(Temperature Stress)
Increase Temperature
Decrease Temperature
Result
Shift in the endothermic
direction
Shift in the exothermic
direction
Note: You must know whether the reaction is
exothermic or endothermic (If you do not know the
energy exchange involved in a reaction then the shift
cannot be determined)
Using Le Châtelier’s Principle
Example
• The reaction 2 SO 2(g) + O 2(g) ⇔ 2 SO 3(g) with ΔH° = -198
kJ is at equilibrium.
– adding more O 2 to the container
– condensing and removing SO 3
– compressing the gases
– cooling the container
– doubling the volume of the container
– warming the mixture
– adding the inert gas helium to the container
– adding a catalyst to the mixture
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