Chem 1B
Dr. White
Chapter 13: Chemical Equilibrium Outline
13.1. Chemical Equilibrium
A. Definition:
B. Consider: N2O4 (g, colorless) ⇄ 2NO2 (g, brown)
C. 3 Main Characteristics of Equilibrium
13.2-13.4. The Equilibrium Constant (K)
A. K –
1. Law of Mass Action –
2. General Description: jA (aq) + kB (aq) <-> lC (aq) + mD (aq)
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a. K –
b. Kp –
i. CO (g) + 3H2 (g) ⇄ CH4 (g) + H2O(g)
ii. Kp =
iii. K =
c. Heterogeneous Equilibrium –
i. S8 (s) + 24F2 (g) ⇄ 8SF6 (g)
3. Lecture Example: Write the equilibrium constant expression (K) for the
following reactions.
a.
N2 (g) + 3H2 (g) ⇄ 2NH3 (g)
b. 3 H2O (l) + H3PO4 (aq) ⇄ 3H3O+ (aq) + PO43-(aq)
c. AgCl (s) ⇄ Ag+ (aq) + Cl- (aq)
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4.
Lecture Example: Calculate K for the reaction in reaction a above if
the equilibrium concentrations were determined at 127°C to be: [NH3] =
3.50 M, [N2] = 1.25 M, and [H2] = 1.50 M
B. Equilibrium Positions –
Consider: N2O4 (g, colorless) ⇄ 2NO2 (g, brown)
C. More about K
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Dr. White
a. K<1
b. K>1
c. Intermediate K
D. Converting between K and Kp
General equation:
Example: The equilibrium constant K for the following reaction at 2127oC is
2.5 x 10-3. What is Kp at this temperature? N2(g) + O2(g)  2 NO(g)
Example: The Kp value for the following reaction at 25°C is 2.2 x 1012. What is K
at this temperature? 2NO (g) + O2 (g)  2NO2 (g)
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13.5. Applications of the Equilibrium Constant
A. Predicting the Reaction Direction
1. The Reaction Quotient (Q) –
Lecture Example: Write the Q expression for the following reaction:
2 SO2 (g) + O2 (g) ⇌ 2 SO3 (g)
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6 Lecture Example: For the reaction, COCl2  CO + Cl2, K = 8.3 X 10–4
(at 360 °C). If at a given time in the reaction, [COCl2] = 0.18 M, [CO]
= 0.35 M, [Cl2] = 0.12 M which direction will the reaction proceed?
Lecture Example: A 50.0 L reaction vessel at 400°C contains 1.00
mol N2, 3.00 mol H2, and 0.500 mol NH3. Will more NH3 be formed or
will it dissociate when the mixture goes to equilibrium at 400°C.
N2 (g) + 3 H2 (g)  2 NH3 (g) K = 0.500 at 400°C
B. Calculating Equilibrium Pressures and Concentrations from Initial Conditions
1. Use the “ICE Box” Method
a. I:
b. C:
c. E:
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Dr. White
2. ICE Box Lecture Examples:
a. CALCULATING K FROM INITIAL & EQUILIBRIUM CONCENTRATIONS
In a 2.00 L vessel, 3.8 mol CO are combined with 12.0 mol H2. At equilibrium, there are
0.60 mol CO. What are the concentrations of the other substances at equilibrium and
what is the equilibrium constant?
CO(g) + 3 H2(g) ⇌ CH4(g) + H2O(g)
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b. DETERMINING EQUILIBRIUM CONCENTRATIONS FROM INITIAL CONCENTRATIONS. (“Perfect
Square”)
The equilibrium constant, K, for the reaction of H2 gas with I2 gas to form HI gas is
57.0 at 700 K. If 1.00 mol of H2 is allowed to react with 1.00 mol I2 in a 10.0 L
vessel at 700 K, what are the concentrations of H2, I2, and HI at equilibrium?
c. DETERMINING EQUILIBRIUM PRESSURES FROM INITIAL PRESSURES. (“Perfect Square”)
CO gas and water vapor react to form CO2 gas and H2 gas. At 1000 °C the Kp for this
reaction is 0.58. If 0.0200 atm of CO and water vapor are placed in a vessel what are
the pressures of each substance at equilibrium?
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9 d. DETERMINING EQUILIBRIUM CONCENTRATIONS FROM INITIAL CONCENTRATIONS. (Starting
with reactants and products)
What are the equilibrium concentrations for the reaction below when you start with
0.100 mol each of CO2 and H2 and 0.0600 mol each of CO and H2O in 1.00 L vessel?
CO2(g) + H2(g) ⇌ CO(g) + H2O(g) K = 0.246
e. DETERMINING EQUILIBRIUM CONCENTRATIONS FROM INITIAL CONCENTRATIONS.
(“Quadratic”)
The equilibrium constant K for the reaction of H2 gas with I2 gas to form HI gas is 49.7
at 458°C. If 1.00 mol of H2 is allowed to react with 2.00 mol I2 in a 1.00 L vessel at 700
K, what are the concentrations of H2, I2, and HI at equilibrium?
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Dr. White
10 f. DETERMINING EQUILIBRIUM CONCENTRATIONS FROM INITIAL CONCENTRATIONS.
(“Simplification”)
Calculate the equilibrium concentration of silver ion in a 1.00 L of solution with 0.010 mol
AgCl (s) and 0.010 mol Cl- in a solution with the equilibrium reaction of:
AgCl (s)  Ag+ (aq) + Cl- (aq)
K = 1.8 x 10-10
Chem 1B
Dr. White
11 13.7 Le Chatelier’s Principle –
A. Concentration Changes
Example: Consider the following reaction:
CH3OH(g) + O2(g)  HCOOH(g) + H2O(g)
a. What direction does the equilibrium shift if more oxygen is added?
b. What direction does the equilibrium shift is water is removed?
Example: Consider the following reaction: AgCl(s)  Ag+(aq) + Cl–(aq)
a. What happens to the concentration of the ions if AgCl is added?
b. What happens to the concentration of the ions if Ag+ is added?
B. Pressure Changes
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Dr. White
12 Lecture Example: How does a decrease in volume affect the concentration
of the first reactant in the following reactions?
a. C2H4(g) + H2(g)  C2H6(g)
b. XeF6(g) Xe(g) + 3 F2(g)
c. C(s) + 2 F2(g)  CF4(g)
d. H2S(g) + Hg(l)  HgS(s) + H2(g)
D. Temperature Changes
1. Endothermic:
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Dr. White
13 2. Exothermic:
Lecture Example: For the following reactions, how does the concentration of
the first reactant change if the temperature decreases?
a. 2 NOCl (g)  2 NO (g) + Cl2 (g)
(endothermic)
b. C2H4 (g) + H2 (g)  C2H6 (g)
(ΔH<0)
13.2 (Revisited) The Equilibrium Constant
A. Manipulating K
1. Direction of reactions
a. A ⇄ B
Ka =
b. B ⇄ A
Kb =
c. Therefore,
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14 2. Reactions as a sum of steps
a.
A⇄B
Ka =
b.
B⇄C
Kb =
c.
A⇄C
Kc =
d. Koverall =
3. Coefficients multiplied by n
a . aA + bB ⇄ cC + dD
Ka =
b. n (aA + bB ⇄ cC + dD) Kb =
c. Therefore,
d. SO2 (g) + ½ O2 (g) ⇄ SO3 (g) K = 56 at 900K
2SO2 (g) + O2 (g) ⇄ 2SO3 (g) K =
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Dr. White
4. Lecture Examples:
Given: CO (g) + 3 H2 (g)  CH4 (g) + H2O (g) K = 3.92 at 1200 K
CH4 (g) + 2 H2S (g)  CS2 (g) + 4 H2 (g) K = 3.3 x 104 at 1200 K
a. Determine K at 1200 K for:
CO (g) + 2H2S (g)  CS2 (g) + H2O (g) + H2(g)
b. Determine Kc at 1200 K for:
1/3CO (g) +
H2 (g)  1/3 CH4 (g) + 1/3 H2O (g)
c. Determine Kc at 1200 K for:
2 CS2 (g) + 8 H2 (g)  2 CH4 (g) + 4 H2S (g)
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