Chapter 14: Chemical Equilibrium

1
Chapter 14 Chemical Equilibrium

• Renee Y. Becker
• Valencia Community College

2
Introduction

• How far does a reaction proceed toward completion
  before it reaches a state of chemical
  equilibrium?
• 2. Chemical equilibrium
• a) The state reached when the concentrations of
  reactants and products remain constant over time
• b) A state in which the concentration of
  reactants and products no longer change (net)
• 3. Equilibrium mixture
• A mixture of reactants and products in the
  equilibrium state

3
Introduction

• What are we interested in?
• a) What is the relationship between the
  concentration of reactants and products in an
  equilibrium mixture?
• b) How can we determine equilibrium
  concentrations from initial concentrations?
• c) What factors can be exploited to alter the
  composition of an equilibrium mixture?

4
The Equilibrium State

• In previous chapters we have generally assumed
  that chemical reactions result in complete
  conversion of reactants to products
• Many reactions do not go to completion!!
• Example1

5
The Equilibrium State
6
The Equilibrium State

• The two experiments demonstrate that the
  interconversion of N2O4 and NO2 is reversible and
  that the same equilibrium state is reached
  starting from either substance.
• 1. This is why we use a ? instead of ?
• 2. Since both NO2 and N2O5 are products and
  reactants we will call the chemical on the left
  reactants and on the right products.
• 3. All chemical reactions are reversible     

7
The Equilibrium State

• We call a reaction irreversible when it proceed
  nearly to completion
• a. Equilibrium mixture contains almost all
  products and almost no reactants
• b. Reverse reaction is too slow to be detected
• 5. In an equilibrium state the reaction does not
  stop at particular concentrations of reactants
  and products, the rates of the forward and
  reverse reactions become equal.
• Important reaction does not stop

8
The Equilibrium State

• 6. Chemical equilibrium is a dynamic state in
  which forward and reverse reactions continue at
  equal rates so that there is no net conversion of
  reactants to products

9
Example 1

• Which of the following is correct?
• Some reactions are truly not reversible
• All reactants go to all products in all reactions
• All reactions are reversible to some extent
• The rates of the forward and reverse reactions
  will never be equal

10
The Equilibrium Constant, Kc

• General equation aA bB ? cC dD
•  
• Equilibrium equation Kc Cc Dd
  products
• 
  Aa Bb reactants
• The substances in the equilibrium equation must
  be gases or molecules and ions in solution, NO
  SOLIDS! NO PURE LIQUIDS!
•  
• Kc units are omitted but you must say at what
  temperature!

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The Equilibrium Constant, Kc

• a A bB ? cC dD
•  
• Kc Cc Dd
• Aa Bb
•  
• If we write the equation in the reverse
  direction
• cC dD ? aA bB
•  
• Kc Aa Bb 1
• Cc Dd kc

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Example 2

•  Write the equilibrium equation for each of the
  following reaction
•  
• a)      N2(g) 3 H2(g) ? 2 NH3(g)
• b) 2 NH3(g) ? N2(g) 3 H2(g)

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Example 3

• The oxidation of sulfur dioxide to give sulfur
  trioxide is an important step in the industrial
  process for synthesis of sulfuric acid. Write
  the equilibrium equation for each of the
  following reactions
• a)      2 SO2(g) O2(g) ? 2 SO3(g)
• b) 2 SO3(g) ? 2 SO2(g) O2(g)
• The following equilibrium concentrations were
  measured at 800 K
• SO2 3.0 x 10-3 M O2 3.5 x 10-3 M
• SO3 5.0 x 10-2 M
•  
• Calculate the equilibrium constant at 800 K a
  and b

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The Equilibrium Constant Kp

• Kp equilibrium constant with respect to partial
  pressures of reactants and products
• a A bB ? cC dD
•  
• Kp (PC)c (PD)d
• (PA)a (PB)b
•  
• Relationship between Kc and Kp
•  
• Kp Kc(RT)?n

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Example 4

• In the industrial synthesis of hydrogen, mixtures
  of CO and H2O are enriched in H2 by allowing the
  CO to react with steam. The chemical equation
  for this so-called water-gas shift reaction is
•  
• CO(g) H2O(g) ? CO2(g) H2(g)
•  
• What is the value of Kp at 700 K if the partial
  pressures in an equilibrium mixture at 700 K are
  1.31 atm of CO, 10.0 atm of H2O, 6.12 atm of
  CO2, and 20.3 atm H2?

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Example 5

• When will kc kp ?
• 1. 2 SO2(g) O2(g) ? 2 SO3(g)
• CO(g) H2O(g) ? CO2(g) H2(g)
• 3. N2(g) 3 H2(g) ? 2 NH3(g)

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Example 6

• Nitric oxide reacts with oxygen to give nitrogen
  dioxide, an important reaction in the Ostwald
  process for the industrial synthesis of nitric
  acid
•  
• 2 NO(g) O2(g) ? 2 NO2(g)
•  
• If Kc 6.9 x 105 _at_ 227?C, what is the value of
  Kp _at_ 227?C?
• b) If Kp 1.3 x 10-2 _at_ 1000 K, what is the
  value of Kc _at_ 1000 K?

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Heterogeneous Equilibria

• Introduction
• 1. So far we have been talking about
  homogeneous equilibria, in which all reactants
  and products are in a single phase (gas or
  solution)
• 2. Heterogeneous equilibria are those in which
  reactants and products are present in more than
  one phase

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Example 7

• For each of the following reactions, write the
  equilibrium constant expression for Kc
•  
• a)   2 Fe(s) 3 H2O(g) ? Fe2O3(s) 3
  H2(g)
• b) 2 H2O(l) ? 2 H2(g) O2(g)
• c) SiCl4(g) 2 H2(g) ? Si(s) 4 HCl(g)
  
• d) Hg22(aq) 2 Cl-(aq) ? Hg2Cl2(s)

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Example 8

• Which of the following has a Heterogeneous
  equilibria?
• 1. 2 SO2(g) O2(g) ? 2 SO3(g)
• CO(g) H2O(g) ? CO2(g) H2(g)
• SiCl4(g) 2 H2(g) ? Si(s) 4 HCl(g)

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Using the Equilibrium Constant

• Introduction
• Knowing the value of the equilibrium constant for
  a chemical reaction lets us
• 1. Judge the extent of the reaction
• 2. Predict the direction of the reaction
• 3. Calculate the equilibrium concentrations from
  any initial concentrations

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Using the Equilibrium Constant

• The numerical value of the equilibrium constant
  for a reaction indicates the extent to which
  reactants are converted to products
• Large value for Kc gt 103 reaction proceeds
  essentially to 100 (mostly products)
• Small value for Kc lt 10-3 reaction proceeds
  hardly at all before equilibrium is reached
  (mostly reactants)
• If a reaction has an intermediate value of Kc
  103 to 10-3
• a.  Appreciable concentrations of both reactants
  and products are present in the equilibrium
  mixture

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Predicting the direction of Reaction

• Reaction Quotient Qc
•       1. Not necessarily equilibrium
  concentrations, at some time, t, snapshot of
  reaction
•       2. As time passes, Qc changes toward the
  value of Kc
•       3. When the equilibrium state is reached
  Qc Kc
•       4. Qc allows us to predict the direction
  of reaction by comparing the values of Kc and Qc
•        a) If Qclt Kc, net reaction goes
  from left to right, (reactant to
  products)
•       b) If Qc gt Kc, net reaction goes
  from right to left, (products to
  reactants)
•       c) If Qc Kc, no net reaction
  occurs

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Example 9

• The equilibrium constant for the reaction
•  
• 2 NO(g) O2(g) ? 2 NO2(g)
•  
• is 6.9 x 105 _at_ 500 K. A 5.0 L reaction vessel
  at this temperature was filled with 0.060 mol of
  NO, 1.0 mol O2, and 0.80 mol NO2.
• a) Is the reaction mixture at equilibrium? If
  not, in which direction does the net reaction
  proceed?
• b) What is the direction of the net reaction if
  the initial amounts are 5.0 x 10-3 mol of NO,
  0.20 mol of O2 and 4.0 mol of NO2?

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Factors that Alter the Composition of an
Equilibrium Mixture

• Introduction
• One of the principal goals of chemical synthesis
  is to maximize the conversion of reactants to
  products while minimizing the expenditure of
  energy.
• 1. Can be achieved if the reaction goes nearly
  to completion at mild temperatures and
  pressures.
• 2. If the equilibrium mixture is high in
  reactants and poor in products, the
  experimental conditions must be changed.
• 3. Several factors can be exploited to alter
  the composition of an equilibrium mixture.
• A.       The concentration of reactants or
  products
•      B. The pressure and volume
• C. The temperature

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Le Chateliers Principle

• Le Chateliers Principle
• If a stress is applied to a reaction mixture at
  equilibrium, net reaction occurs in the direction
  that relieves the stress
•    1. Stress means a change in the
  concentration, pressure, volume, or temperature
  that disturbs the original equilibrium
•     2. Reaction then occurs to change the
  composition of the mixture until a new state of
  equilibrium is reached
• 3. The direction that the reaction takes
  (reactants to products or products to reactants)
  is the one that reduces the stress

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Altering an Equilibrium Mixture Changes in
Concentration  

• In general, when an equilibrium is disturbed by
  the addition or removal of any reactant or
  product, Le Chateliers principle predicts that
• 1.  The concentration stress of an added
  reactant or product is relieved by net reaction
  in the direction that consumes the added
  substance
• 2. The concentration stress of a removed
  reactant or product is relieved by net reaction
  in the direction that replenishes the removed
  substance

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Example 10

• Consider the equilibrium for the water-gas shift
  reaction
•  
• CO(g) H2O(g) ? CO2(g) H2(g)
•  
• Use Le Chateliers principle to predict how the
  concentration of H2 will change and what
  direction the reaction will flow when the
  equilibrium is disturbed by
•  
• 1.      Adding CO
• 2.      Adding CO2
• 3.      Removing H2O
• 4.      Removing CO2
•  

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Example 11

• In the following reaction, if I take away CO,
  which direction will the reaction proceed to
  equilibrium?
• CO2(g) H2(g) ?CO(g) H2O(g)
• Products ?
• Reactants ?

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Altering an Equilibrium Mixture Changes in
Pressure and Volume

• In general Le Chateliers Principle predicts
  that
• 1. An increase in pressure by reducing the
  volume will bring about net reaction in the
  direction that decreases the number of moles of
  gas
• 2. A decrease in pressure by enlarging the
  volume will bring about net reaction in the
  direction that increases the number of moles of
  gas.

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Example 12

• Which direction will the reaction flow if the
  following equilibria is subjected to an increase
  in pressure by decreasing the volume?
• 1.      CO(g) H2O(g) ? CO2(g) H2(g)
•  
• 2.      2 CO(g) ? C(s) CO2(g)
•  
• 3.      N2O4(g) ? 2 NO2(g)

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Example 13

• If I increase the pressure by decreasing the
  volume, which direction will the reaction flow to
  reach equilibrium?
• C(s) CO2(g) ? 2 CO(g)
• Products ?
• Reactants ?

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Altering the Equilibrium Mixture Changes in
Temperature

• In general, the temperature dependence of the
  equilibrium constant depends on the sign of ?H?
  for the reaction
• 1.      The equilibrium constant for an
  exothermic reaction (negative ?H?) decreases as
  the temperature increases
• 2.      The equilibrium constant for an
  endothermic reaction (positive ?H?) increases as
  the temperature increases.
• 3.      ?H? standard enthalpy of reaction,
  enthalpy change measured under standard
  conditions
• 4. Standard conditions most stable form of a
  substance at 1 atm pressure and at a specified
  temperature, usually 25?C 1 M concentration for
  all substances

34
Altering the Equilibrium Mixture Changes in
Temperature

• Le Chateliers Principle says that if heat is
  added to an equilibrium mixture (increasing the
  temperature) net reaction occurs in the direction
  that relieves the stress of the added heat.
• 1.     For an endothermic reaction heat is
  absorbed by reaction in the forward direction.
  The equilibrium shifts to the right at the
  higher temperatures, Kc increases with increasing
  temperature
• 2.     For an exothermic heat is absorbed by net
  reaction in the reverse direction, so Kc
  decreases with temperature, and the reaction
  would flow to the left (reactants)

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Example 14

• When air is heated at very high temperatures in
  an automobile engine, the air pollutant nitric
  oxide is produced by the reaction
•  
• N2(g) O2(g) ? 2 NO(g) ?H? 180.5 kJ
•  
• 1. How does the equilibrium amount of NO vary
  with an increase in temperature?
• 2. What direction is the net reaction flowing?

36
The Effect of a Catalyst on Equilibrium

• A catalyst increases the rate of a chemical
  reaction by making available a new, lower-energy
  pathway for conversion of reactants to products.
• 1.  Since the forward and reverse reaction pass
  through the same transition state, a catalyst
  lowers the activation energy for both
• 2.   The rates of the forward and reverse
  reactions increase by the same factor
• 3. Catalyst accelerates the rate at which
  equilibrium is reached
• 4. Catalyst does not affect the composition of
  the equilibrium mixture

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The Effect of a Catalyst on Equilibrium
38
Example 15

• A platinum catalyst is used in automobile
  catalytic converters to hasten the oxidation of
  carbon monoxide
• 2 CO(g) O2(g) ? 2 CO2(g) ?H? -566
  kJ
•  
• Suppose that you have a reaction vessel
  containing an equilibrium mixture. Will the
  amount of CO increase, decrease, or remain the
  same when
• A platinum catalyst is added
• The temperature is increased
• The pressure is increased by decreasing the
  volume
• The pressure is increased by adding argon gas
• The pressure is increased by adding O2 gas

39
The Link Between Chemical Equilibrium and
Chemical Kinetics

• A B ? C D
•  
• Assuming that the forward and reverse reactions
  occur in a single bimolecular step, elementary
  reactions, we can write the following rate laws
•  
• Rate of forward reaction kf A B
• Rate of reverse reaction kr C D
• When t0 C D 0
• As A and B are converted to C and D the rate of
  the forward reaction decreases and the rate of
  the reverse reaction is increasing, until they
  are equal, chemical equilibrium
•  
• kf A B kr C D

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The Link Between Chemical Equilibrium and
Chemical Kinetics

• kf C D
• kr A B
• The right side of this equation is the
  equilibrium constant expression for the forward
  reaction, which equals the equilibrium constant
  Kc
•  
• Kc C D
• A B
•  
• Therefore the equilibrium constant is simply the
  ratio of the rate constants for the forward and
  reverse reactions
•  
• Kc kf
• kr

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Example 16

• Nitric oxide emitted from the engines of
  supersonic transport planes can contribute to the
  destruction of stratospheric ozone
•  NO(g) O3(g) ? NO2(g) O2(g)
•  
• This reaction is highly exothermic (?E -200
  kJ), and its equilibrium constant Kc is 3.4 x
  1034 at 300 K
• Which rate constant is larger, kf or kr?
• The value of kf at 300 K is 8.5 x 106 M-1 s-1.
  What is the value of kr at the same temperature?
• A typical temperature in the stratosphere is 230
  K. Do the values of kf, kr, and Kc increase or
  decrease when the temperature is lowered from 300
  K to 230 K?

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Example 17

• If I increase the temperature of reaction which
  way will the reaction flow to equilibrium?
• NO2(g) O2(g) ? NO(g) O3(g) ?H
  200 kJ
• Products ?
• Reactants ?