chapter 15 chemical equilibrium - ksu...
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![Page 1: Chapter 15 Chemical Equilibrium - KSU Facultyfac.ksu.edu.sa/sites/default/files/Ch15_lectures_of_Chemical_equlibria.pdf · Chapter 15 Chemical Equilibrium Dr. Ayman Nafady John D](https://reader030.vdocument.in/reader030/viewer/2022040109/5e69c7f2163c1c4d1f6e2bba/html5/thumbnails/1.jpg)
Equilibrium
© 2009, Prentice-Hall, Inc.
Chapter 15
Chemical Equilibrium
Dr. Ayman Nafady
John D. Bookstaver
St. Charles Community College
Cottleville, MO
Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and
Bruce E. Bursten
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Concept of Equilibrium
Chemical equilibrium occurs when a
reaction and its reverse reaction proceed at
the same rate.
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Concept of Equilibrium
• As a system
approaches equilibrium,
both the forward and
reverse reactions are
occurring.
• At equilibrium, the
forward and reverse
reactions are
proceeding at the same
rate.
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Equilibrium
© 2009, Prentice-Hall, Inc.
A System at Equilibrium
Once equilibrium is
achieved, the
amount of each
reactant and product
remains constant.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Depicting Equilibrium
Since, in a system at equilibrium, both
the forward and reverse reactions are
being carried out, we write its equation
with a double arrow.
N2O4 (g) 2 NO2 (g)
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Equilibrium
© 2009, Prentice-Hall, Inc.
The
Equilibrium
Constant
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Equilibrium
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The Equilibrium Constant
• Forward reaction:
N2O4 (g) 2 NO2 (g)
• Rate Law:
Rate = kf [N2O4]
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Equilibrium
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The Equilibrium Constant
• Reverse reaction:
2 NO2 (g) N2O4 (g)
• Rate Law:
Rate = kr [NO2]2
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Equilibrium Constant
• Therefore, at equilibrium
Ratef = Rater
kf [N2O4] = kr [NO2]2
• Rewriting this, it becomes
kf
kr
[NO2]2
[N2O4] =
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Equilibrium
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The Equilibrium Constant
The ratio of the rate constants is a
constant at that temperature, and the
expression becomes
Keq = kf
kr
[NO2]2
[N2O4] =
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Equilibrium
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The Equilibrium Constant
• Consider the generalized reaction
• The equilibrium expression for this
reaction would be
Kc = [C]c[D]d
[A]a[B]b
aA + bB cC + dD
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Equilibrium
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The Equilibrium Constant
Since pressure is proportional to
concentration for gases in a closed
system, the equilibrium expression can
also be written
Kp = (PC
c) (PDd)
(PAa) (PB
b)
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Equilibrium
© 2009, Prentice-Hall, Inc.
Relationship Between Kc and Kp
• From the Ideal Gas Law we know that
• Rearranging it, we get
PV = nRT
P = RT n
V
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Equilibrium
© 2009, Prentice-Hall, Inc.
Relationship Between Kc and Kp
Plugging this into the expression for Kp
for each substance, the relationship
between Kc and Kp becomes
where
Kp = Kc (RT)n
n = (moles of gaseous product) - (moles of gaseous reactant)
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Equilibrium
© 2009, Prentice-Hall, Inc.
Equilibrium Can Be Reached from
Either Direction
As you can see, the ratio of [NO2]2 to [N2O4] remains
constant at this temperature no matter what the initial
concentrations of NO2 and N2O4 are.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Equilibrium Can Be Reached from
Either Direction
This is the data from
the last two trials from
the table on the
previous slide.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Equilibrium Can Be Reached from
Either Direction
It doesn’t matter whether we start with N2 and
H2 or whether we start with NH3: we will have
the same proportions of all three substances
at equilibrium.
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Equilibrium
© 2009, Prentice-Hall, Inc.
What Does the Value of K Mean?
• If K>>1, the reaction is
product-favored;
product predominates
at equilibrium.
• If K<<1, the reaction is
reactant-favored;
reactant predominates
at equilibrium.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Manipulating Equilibrium Constants
The equilibrium constant of a reaction in
the reverse reaction is the reciprocal of
the equilibrium constant of the forward
reaction.
Kc = = 0.212 at 100 C [NO2]
2
[N2O4] N2O4 (g) 2 NO2 (g)
Kc = = 4.72 at 100 C [N2O4]
[NO2]2
N2O4 (g) 2 NO2 (g)
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Equilibrium
© 2009, Prentice-Hall, Inc.
Manipulating Equilibrium Constants
The equilibrium constant of a reaction that has
been multiplied by a number is the equilibrium
constant raised to a power that is equal to that
number.
Kc = = 0.212 at 100 C
[NO2]2
[N2O4] N2O4(g) 2 NO2(g)
Kc = = (0.212)2 at 100 C [NO2]
4
[N2O4]2
2 N2O4(g) 4 NO2(g)
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Equilibrium
© 2009, Prentice-Hall, Inc.
Manipulating Equilibrium Constants
The equilibrium constant for a net
reaction made up of two or more steps
is the product of the equilibrium
constants for the individual steps.
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Equilibrium
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Sample Exercise 15.1 Writing Equilibrium-Constant Expressions
Write the equilibrium expression for Kc for the following reactions:
Solution
Practice Exercise
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Equilibrium
© 2009, Prentice-Hall, Inc.
Sample Exercise 15.2 Converting between Kc and Kp
In the synthesis of ammonia from nitrogen and hydrogen,
Kc = 9.60 at 300 °C. Calculate Kp for this reaction at this
temperature.
Solution Analyze: We are given Kc for a reaction and asked to calculate Kp.
Plan: The relationship between Kc and Kp is given by Equation 15.14. To apply that
equation, we must determine Δn by comparing the number of moles of product with
the number of moles of reactants (Equation 15.15).
Solve: There are two moles of gaseous products (2 NH3) and four moles of gaseous
reactants (1 N2 + 3 H2). Therefore, Δn = 2 – 4 = –2. (Remember that Δ functions are
always based on products minus reactants.) The temperature, T, is 273 + 300 = 573
K. The value for the ideal-gas constant, R, is 0.0821 L-atm/mol-K. Using Kc = 9.60, we
therefore have
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Equilibrium
© 2009, Prentice-Hall, Inc.
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Equilibrium
© 2009, Prentice-Hall, Inc.
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Equilibrium
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Heterogeneous
Equilibrium
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Equilibrium
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The Concentrations of Solids and
Liquids Are Essentially Constant
Both can be obtained by multiplying the
density of the substance by its molar
mass — and both of these are constants
at constant temperature.
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Equilibrium
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The Concentrations of Solids and
Liquids Are Essentially Constant
Therefore, the concentrations of solids
and liquids do not appear in the
equilibrium expression.
Kc = [Pb2+] [Cl-]2
PbCl2 (s) Pb2+ (aq) + 2 Cl-(aq)
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Equilibrium
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As long as some CaCO3 or CaO remain in
the system, the amount of CO2 above the
solid will remain the same.
CaCO3 (s) CO2 (g) + CaO(s)
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Equilibrium
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Equilibrium
Calculations
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Equilibrium
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An Equilibrium Problem
A closed system initially containing 1.000
x 10-3 M H2 and 2.000 x 10-3 M I2 at
448 C is allowed to reach
equilibrium. Analysis of the equilibrium
mixture shows that the concentration of
HI is 1.87 x 10-3 M. Calculate Kc at 448
C for the reaction taking place, which is
H2 (g) + I2 (s) 2 HI (g)
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Equilibrium
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What Do We Know?
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change
At equilibrium 1.87 x 10-3
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Equilibrium
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[HI] Increases by 1.87 x 10-3 M
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change +1.87 x 10-3
At equilibrium 1.87 x 10-3
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Equilibrium
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Stoichiometry tells us [H2] and [I2] decrease
by half as much.
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change -9.35 x 10-4 -9.35 x 10-4 +1.87 x 10-3
At equilibrium 1.87 x 10-3
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Equilibrium
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We can now calculate the equilibrium
concentrations of all three compounds…
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change -9.35 x 10-4 -9.35 x 10-4 +1.87 x 10-3
At equilibrium 6.5 x 10-5 1.065 x 10-3 1.87 x 10-3
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Equilibrium
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…and, therefore, the equilibrium constant.
Kc = [HI]2
[H2] [I2]
= 51
= (1.87 x 10-3)2
(6.5 x 10-5)(1.065 x 10-3)
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Equilibrium
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The Reaction Quotient (Q)
• Q gives the same ratio the equilibrium
expression gives, but for a system that is
not at equilibrium.
• To calculate Q, one substitutes the initial
concentrations on reactants and products
into the equilibrium expression.
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Equilibrium
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If Q = K,
the system is at equilibrium.
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Equilibrium
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If Q > K, there is too much product, and the
equilibrium shifts to the left.
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Equilibrium
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If Q < K, there is too much reactant, and the
equilibrium shifts to the right.
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Equilibrium
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Le Châtelier’s
Principle
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Equilibrium
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Le Châtelier’s Principle
“If a system at equilibrium is disturbed
by a change in temperature, pressure,
or the concentration of one of the
components, the system will shift its
equilibrium position so as to counteract
the effect of the disturbance.”
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Haber Process
The transformation of nitrogen and hydrogen into
ammonia (NH3) is of tremendous significance in
agriculture, where ammonia-based fertilizers are of
utmost importance.
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Haber Process
If H2 is added to the
system, N2 will be
consumed and the
two reagents will
form more NH3.
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Haber Process
This apparatus
helps push the
equilibrium to the
right by removing
the ammonia (NH3)
from the system as
a liquid.
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Equilibrium
© 2009, Prentice-Hall, Inc.
The Effect of Changes in
Temperature Co(H2O)6
2+(aq) + 4 Cl(aq) CoCl4 (aq) + 6 H2O (l)
QuickTime™ and aPhoto - JPEG decompressor
are needed to see this picture.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Catalysts
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Equilibrium
© 2009, Prentice-Hall, Inc.
Catalysts
Catalysts increase
the rate of both the
forward and reverse
reactions.
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Equilibrium
© 2009, Prentice-Hall, Inc.
Catalysts
When one uses a
catalyst, equilibrium
is achieved faster,
but the equilibrium
composition remains
unaltered.