Download Chapter 13

Chapter 13
Chemical Equilibrium
The state where the concentrations of all
reactants and products remain constant with
time.
On the molecular level, there is frantic
activity. Equilibrium is not static, but is a
highly dynamic situation.
Figure 13.1 A Molecular Representation of the Reaction
2NO2(g) 2O4(g) Over Time in a Closed Vessel
Figure 13.2 Changes in Concentrations
H2O(g) + CO(g)  H2(g) + CO2(g)
Figure 13.4 The Changes with Time in the Rates of Forward and Reverse Reactions
H2O(g) + CO(g)  H2(g) + CO2(g)
Figure 13.5 The Ammonia Synthesis Equilibrium
N2(g) + 3H2(g)  2NH3(g)
The Law of Mass Action
jA + kB  lC + mD
where A, B, C, and D represents chemical species
and j, k, l, and m are their coefficient in the
balanced equation.
The law of mass action is represented by the
equilibrium expression:
l
m
C D
K
A j Bk
The square brackets indicate the concentrations of
the chemical species at equilibrium, and K is a
constant called the equilibrium constant.
Equilibrium Expression
Write the equilibrium expression for the following
reaction:
4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g)
Applying the law of mass action gives,
4
6
NO2 H 2O
K
4
7
NH3 O2
Superscript 4, 6, 4, and 7 are the coefficients of NO2,
H2O, NH3, and O2 respectively. The value of the
equilibrium constant at a given temperature can be
calculated if we know the equilibrium concentrations of
the reaction components.
Notes on Equilibrium Expressions
(EE)
 The
Equilibrium Expression for a reaction is
the reciprocal of that for the reaction written
in reverse.
 When the equation for a reaction is multiplied
by n, EEnew = (EEoriginal)n
 The units for K depend on the reaction being
considered. K values are customarily written
without units.
Equilibrium Expressions Involving Pressures:
Equilibria involving gases can be described in terms of
pressures.
 n
PV = nRT  P =   RT = CRT
Ideal gas equation:
 v
where, C equals n/v or number of moles n per unit volume
V. Thus C represents the molar concentration of the gas.
N2(g) + 3H2(g)  2NH3(g)
2
2
[NH3]
CNH3
K=

3  Kc
3
[N2][H2] (CN2)(CH2)
In terms of the equilibrium partial pressures of the gasses
2
PNH3
Kp =
3
(PN2)(PH2 )
K v. Kp
For
jA + kB  lC + mD
Kp = K(RT)n
n = sum of coefficients of gaseous products minus sum of
coefficients of gaseous reactants.
l
m
l m
( PC )( PD) (Cc  RT ) (CD  RT )
Kp 
j
k 
( PA )( PB ) (CA  RT ) j (CB  RT )k
m
l
n
(l  m)  ( j  k )
(Cc)(CD) ( RT ) l  m


 K ( RT )
 K ( RT )
j
k
j

k
(CA)(CB )
( RT )
Heterogeneous Equilibria
. . . are equilibria that involve more than one
phase.
CaCO3(s)  CaO(s) + CO2(g)
K = [CO2]
The position of a heterogeneous equilibrium
does not depend on the amounts of pure
solids or liquids present.
Figure 13.6 CaCO3(s)  CaO(s) + CO2(g)
Examples
The decomposition of liquid water to gaseous hydrogen
and oxygen,
2H2O(l)
2H2(g) + O2(g)
K = [H2]2[O2] and Kp=(P2H2)(PO2)
Water is not included in either equilibrium expression
because it is a pure liquid. However, if water is a gas
rather than a liquid,
2H2O(g)
2H2(g) + O2(g)
2
2
[ H 2 ] [ O2 ]
( PH 2)( PO 2)
2
K
and Kp 
2
[ H 2O]
PH 2 O
because the concentration or pressure of water vapor can
change.
Reaction Quotient
. . . helps to determine the direction of the move toward
equilibrium.
The reaction quotient is obtained by applying the law of
mass action using initial concentrations instead of
equilibrium concentrations.
H2(g) + F2(g)  2HF(g)
Q
HF
H2
0
0
2
F2
0
where the subscript zeros indicate initial concentrations.
Direction of reaction
To determine in which direction a system will
shift to reach equilibrium, we compare the values
of Q and K. There are three possible cases:
1.Q is equal to K. The system is at equilibrium; no
shift will occur.
2.Q is greater than K. The system shifts to the left,
consuming products and forming reactants, until
equilibrium is achieved.
3.Q is less than K. The system shifts to the right,
consuming reactants and forming products, to
attain equilibrium.
Solving Equilibrium Problems
1.
2.
3.
4.
5.
6.
7.
Balance the equation.
Write the equilibrium expression.
List the initial concentrations.
Calculate Q and determine the shift to
equilibrium.
Define equilibrium concentrations.
Substitute equilibrium concentrations into
equilibrium expression and solve.
Check calculated concentrations by
calculating K.
Le Châtelier’s Principle
. . . if a change is imposed on a
system at equilibrium, the position
of the equilibrium will shift in a
direction that tends to reduce that
change.
Effects of Changes on the System
1. Concentration: The system will shift away
from the added component.
2. Temperature: K will change depending upon
the temperature (treat the energy change as a
reactant).
3. Pressure:
a. Addition of inert gas does not affect the
equilibrium position.
b. Decreasing the volume shifts the
equilibrium toward the side with fewer moles.
Figure 13.8 A Mixture of N2, H2, and NH3
Figure 13.9 The Effect of Decreased Volume on the Ammonia Synthesis Equilibrium