C h a p t e r 13
Chemical Equilibrium
The Equilibrium State
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01
Chemical Equilibrium: A state achieved when the rates of
the forward and reverse reactions are equal and the
concentrations of the reactants and products remain
constant.
1. What is the relationship between the concentrations
of reactants and products in an equilibrium mixture?
2. Determine equilibrium concentrations from initial
concentrations.
3. How to alter the composition of an equilibrium
mixture (concentration, pressure, temp, catalyst).
The Equilibrium State
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02
Thus far, we have assumed complete conversion from
reactants to products.
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Many reactions do not go to completion
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Concentrations do not reach constant values because the
reaction stops, but because the rates of the forward and
reverse reactions become equal.
The Equilibrium State
Reversible reactions (arrows in both directions)
“reactants”
“products”
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The Equilibrium State
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Graphs of reactant and product concentrations
change with time as shown below.
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The Equilibrium State
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Rate of the forward reaction (N2O4  2NO2) decreases as
concentration of N2O4 decreases, while the rate of the reverse reaction
(N2O4  2NO2) increases as the concentration of NO2increases
Equilibrium Constant Kc
06
What is the relationship between the concentrations of
reactants and products in an equilibrium mixture?
aA + bB  cC + dD
Equilibrium Constant:
Kc = [C]c[D]d  products
[A]a[B]b  reactants
Kc is independent of concentration changes, but dependent
on the temperature.
Equilibrium Constant Kc
Kc values are reported without units
Thermodynamic state: [ ] = 1M
N2O4  2NO2
Equilibrium Constant:
Kc = [NO2]2 = (0.0125M/1M)2
[N2O4] (0.0337M/1M)
= 4.64 x 10-3 at 25˚C
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Equilibrium Constant Kc’
cC + dD  aA + bB
Equilibrium Constant:
Kc’ = [A]a[B]b  products
[C]c[D]d  reactants
Kc = 1/ Kc’
* important
to specify the form of the balanced equation
08
Example
What is the equilibrium equation?
(a) N2(g) + 3H2(g)  2NH3(g)
(b) 2NH3(g)  N2(g) + 3H2(g)
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Equilibrium Constant Kp
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10
Equilibrium equations for gas-phase reactions
(partial pressures)
aA(g) + bB(g)  cC(g) + dD(g)
Equilibrium Constant: Kp = (PC)c(PD)d
(PA)a(PB)b
 products
 reactants
Equilibrium Constant Kp
Convert between Kc and Kp using PV = nRT
PAV = nART PA = nART = [A]RT
V
Kp = Kc (RT) ∆n
∆n = moles gas products – moles of gas reactants
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Equilibrium Constant
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12
Homogeneous Equilibrium: When all reacting
species are in the same phase, all reactants and
products are included in the expression.
•
Amounts of components are given as molarity or
partial pressure of a gas.
]2
[NO2
Kc = [
N2O4 ]
Kp =
2
PNO2
PN2O4
Example
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13
The following pictures represent mixtures of A molecules
(red) and B molecules (blue), which interconvert according
to the equation A  B. If Mixture (1) is at equilibrium, which
of the other mixtures is also at equilibrium?
Examples
14
1. Write the Kp and Kc expressions for:
2 N2O5(g)  4 NO2(g) + O2(g)
2. The equilibrium concentrations for the reaction
between CO and Cl2 to form carbonyl chloride
(phosgene gas) CO(g) + Cl2(g)  COCl2(g) at 74°C
are: [CO] = 1.2 x 10–2 M, [Cl2] = 0.054 M, and
[COCl2] = 0.14 M. Calculate Kc and Kp.
Example
15
Methane (CH4) reacts with hydrogen sulfide to yield H2 and
carbon disulfide, a solvent used in manufacturing. What is
the value of Kp at 1000 K if the partial pressures in an
equilibrium mixture at 1000 K are 0.20 atm of CH4, 0.25 atm
of H2S, 0.52 atm of CS2, and 0.10 atm of H2?
Equilibrium Constant
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16
Heterogeneous Equilibrium: When reacting
species are in different phases, solid and liquid
phases are excluded from the expression because
their concentrations “do not change.”
CaCO3(s)  CaO(s) + CO2(g)
Kc = [CO2] because CaCO3 and CaO are solids.
Equilibrium Constant
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Examples
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Write the equilibrium equation for each of the
following reactions:
(a) CO2(g) + C(s)  2 CO(g)
(b) Hg(l) + Hg2+(aq)  Hg22+(aq)
(c) 2 Fe(s) + 3 H2O(g)  Fe2O3(s) + 3 H2(g)
(d) 2 H2O(l)  2 H2(g) + O2(g)
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Using Equilibrium Constants
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We can make the following generalizations
concerning the composition of equilibrium mixtures:
If Kc > 103, products predominate over reactants. If Kc is
very large, the reaction is said to proceed to completion.
If Kc is in the range 10–3 to 103, appreciable
concentrations of both reactants and products are present.
If Kc < 10–3, reactants predominate over products. If Kc is
very small, the reaction proceeds hardly at all.
Predicting Reaction direction
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The reaction quotient (Qc) is obtained by
substituting initial concentrations into the
equilibrium constant.
H2(g) + I2(g)  2HI(g)
Qc =
Qc > K c
Qc = K c
Qc < K c
[HI]t2
[H2]t[I2]t
System proceeds to form reactants.
System is at equilibrium.
System proceeds to form products.
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Predicting Reaction direction
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Example
22
The equilibrium constant (Kc) for the formation of nitrosyl
chloride, from nitric oxide and chlorine gas:
2 NO(g) + Cl2(g)  2 NOCl(g) is 6.5 x 104 at 35°C.
In an experiment, 2.0 x 10–2 moles of NO, 8.3 x 10–3 moles
of Cl2, and 6.8 moles of NOCl are mixed in a 2.0-L flask. In
which direction will the system proceed to reach equilibrium?