Equilibrium
Read pages 539-540 and 545-548 and complete the following chart.
Definition / ExampleReversible
Reaction
Forward
Reaction
Reverse
Reaction
Equilibrium
Keq
What does it mean if the Keq is greater than 1?
Assignment 1: Equilibrium Constants:
1. Write the expression for the equilibrium constant for the following reaction:
carbon monoxide + oxygen <--> carbon dioxide
2. Write the expression for the equilibrium constant for the production of HI gas from
hydrogen gas and iodine gas.
3. For the equilibrium:
sulfur dioxide + oxygen <--> sulfur trioxide
at a certain temperature, [SO2] = 0.20 M, [O2] = 0.20 M, and [SO3] = 0.40 M
Calculate Keq.
4. If in problem #3, [SO2] = 0.10 M and [O2] = 0.10 M, what is the concentration of SO3?
(Keq was calculated in problem 3)
Assignment 2: Equilibrium Constants:
Two equilibrium experiments were done involving the following reaction at 35oC:
2 NF2 (g) à N2F4 (g)
Trial / Initial Concentrations / Equilibrium Concentrations[NF2] / [N2F4] / [NF2] / [N2F4]
1 / 1.00 / 1.00 / 0.50 / 1.24
2 / 0.25 / 0.75 / 0.36 / 0.64
1. Calculate Keq for both trials to verify it is constant. Show work.
Trial 1 Keq = Trial 2 =
Complete the following statements:
a. In trial 1 the reaction proceeds predominantly in the direction
to make ______.
b. In trial 2 the reaction proceeds predominantly in the direction
to make ______.
3. In a third trial at the same temperature [NF2] at equilibrium is determined to be
0.100M. What must [N2F4] be?
2. At a different temperature, 0.500M N2F4 is mixed with no NF2 initially present.
The equilibrium concentration of N2F4 is 0.200M. Calculate the new Keq value.
Assignment 3: Equilibrium Constants and Concentrations:
1. Consider the following reaction: 2 HCl (g) ßà H2 (g) Cl2 (g)
The initial concentration of HCl is 2.0M and there is no H2 or Cl2 present. After
equilibrium conditions have been established, the concentration of Cl2 is 0.10M.
What the equilibrium concentrations for HCl and H2?
Calculate Keq .
2. The reaction, A + B ß à C + D has a Keq of 6.57 x 10-3. Determine the
final equilibrium concentrations of all substances if 0.200M C and 0.200M D are
mixed. No A or B are initially present.
3. For the reaction, N2 (g) + O2 (g) ß à 2 NO (g), the equilibrium constant is
found to be 1.0 at room temperature. If 1.00M N2 is mixed with 1.00M O2, find
the eventual equilibrium concentration of all substances.
4. If at a higher temperature, Keq changes to 0.50 for the reaction in #3, find the
eventual equilibrium concentration of all substances if 0.25moles of NO are
placed in a 4.0L container. (No other substances are present initially.)
5. If in the reaction H2 (g) + Cl2 (g) ß à 2 HCl (g) + 53 kcal, the equilibrium concentrations are: H2 = [1.00 x 10-4], Cl2 = [1.00 x 10-2], and HCl = [9.00 x 10-1] Find the value of K.
6. In the same reaction at a different temperature: initially we put 1.00 mole of hydrogen and an undetermined quantity of chlorine into a 1.00L container. The Keq was determined to be 1.00 x 102. At equilibrium, it is found that the HCl concentration is [0.500]. What are the equilibrium concentrations of the hydrogen and chlorine? What was the initial concentration of chlorine?
7. In the same reaction as # 6, where the Keq is 1.00 x 102 M and the initial concentration of both hydrogen and chlorine are 1.0 M, what are the equilibrium concentrations of all three species?
8. In the reaction CO (g) + H2O (g) ßà CO2 (g) + H2 (g) (the Keq is 5.0), if the initial concentrations for all species is [1.00] find the equilibrium concentrations of each substance.
Assignment 4: LaChatelier’s Principle
La Chatelier’s Principle states that:
Stress / EffectChange
in
Concentration
Change
in
Pressure
Change
in
Temperature
Addition
of a
Catalyst
Use the following equation to answer the following questions:
N2 (g) + 3H2 (g) 2NH3 (g) + heat
1. Write the equilibrium constant expression
2. Calculate the value of Keq if at equilibrium nitrogen is 2.0M, hydrogen is 1.0M, and ammonia is 0.50M.
3. Will there be more reactant or product?
4. Will the reactants or products be favored if:
a. the pressure is decreased.
b. the reaction is cooled.
c. a catalyst is added.
d. NH3 is added.
e. H2 is added.
f. some nitrogen is removed.
5. How will the following change the value of Keq?
a. Ammonia is added.
b. The reaction is heated.
c. The pressure is increased.
d. A catalyst is added.
Assignment 5: LaChatelier’s Principle
For the reaction: 2 SO2 (g) + O2 (g) ß à 2 SO3 (g) + heat
Eq. Concentrations [3.0] [1.0] [2.0]
1. Write the equilibrium constant expression.
2. Calculate the value of Keq.
3. Are the products or the reactants favored?
4. Will the reactants or products be favored if:
a. the pressure is decreased.
b. the reaction is cooled.
c. a catalyst is added
d. SO2 is added.
e. SO3 is added.
f. some SO3 is removed.
5. How will the following stress change the value of Keq?
a. SO2 is added.
b. Reaction is heated.
c. Pressure is increased.
d. A catalyst is added.
Assignment 6: La Chatelier’s Principle
1. Write the equilibrium constant expression for each of the following reactions:
H2 (g) + Cl2 (g) ß à 2 HCl (g) + 55 kcal
N2 (g) + 3 H2 (g) ß à 2 NH3 (g) + 22 kcal
H2CO3 (aq) ß à CO2 (aq) + H2O (l)
NO2 (g) + CO (g) ß à CO2 (g) + NO (g) + 12 kcal
15 kcal + NH4Cl (s) ß à NH4+1 (aq) + Cl-1 (aq)
2. If, in each of the above, after equilibrium is established, an additional quantity of
the underlined substance is added, how would the concentration of each of the other reagents be affected? What happens to the value of Keq?
a. Keq
b. Keq
c. Keq
d. Keq
e. Keq
3. If each of the above reactions, at equilibrium, were heated, would the new
equilibrium favor the reactants or the products? What happens to the value of Keq?
a. Keq
b. Keq
c. Keq
d. Keq
e. Keq
4. Which would be shifted by an increase in pressure? What would be
favored? What happens to the value of Keq
a. Keq
b. Keq
c. Keq
d. Keq
e. Keq
Assignment 7: Equilibrium Constants and La Chatelier’s Principle
1. Initially 4.0M H2 and 2.0M N2 are placed in a container at a certain temperature. After the equilibrium has been established 1.0M NH3 has been produced.
What are the equilibrium concentrations of H2 and N2?
Calculate the Keq at that temperature.
2. For the following reaction: N2O4 (g) + 58.9 kJ ß à 2 NO2 (g)
A 1.0L flask at 55oC is found to contain 3.6 moles of N2O4 and 1.75 moles of NO2 at equilibrium. What is the value of Keq?
If the reaction is heated, in what direction would the equilibrium shift?
Would the Keq change? If so, would it increase or decrease?
If NO2 is removed from the flask, in what direction would the equilibrium shift?
Would the Keq change? If so, would it increase or decrease?
What effect would increasing the pressure have on this equilibrium?
3. If in the reaction H2 (g) + Cl2 (g) ß à 2 HCl (g) + 53 kcal the equilibrium
concentrations are:
H2 = 1.00 x 10-4 M
Cl2 = 1.00 x 10-2 M
HCl = 9.00 x 10-1 M
Find the numerical value of the equilibrium constant (Keq)
4. In the reaction for the question above, if the Keq at a different temperature is 1.00 x 102and the initial concentrations of both H2 and Cl2 are 1.0M, what are the equilibrium concentrations for all three species?
Was the temperature increased or decreased from the first question to the second
5. In the same reaction, initially we put 1.00 mole of H2 and an undetermined amount of Cl2 into a 1.00L container. At equilibrium, it is found that the HCl concentration is 0.50M.
What is the initial concentration of Cl2? (Kreq = 1.00 x 102)
In what direction would the equilibrium shift if more Cl2 was added?
In what direction would the equilibrium shift if a catalyst was added?
Assignment 8: Equilibrium Constants and La Chatelier’s Principle:
1. Initially 4.0M H2 and 2.0M N2 are placed in a container at a temperature of 500K. After the equilibrium has been established 1.0M NH3 has been produced.
What are the equilibrium concentrations of H2 and N2?
Calculate the Keq at that temperature.
2. At a temperature of 800K, 1.0M H2 and 1.0M N2 and 6.0M NH3 are mixed in a container. At equilibrium, the concentration of N2 is determined to be 3.0M.
What are the equilibrium concentrations of H2 and NH3?
What is the Keq at 800K?
Is the reaction exothermic or endothermic? Explain.
Under what conditions could you maximize the production of NH3 (ammonia)?
Assignment 9: Equilibrium Review:
Describe what is happening as a system goes toward equilibrium. What is true at equilibrium and what happens at each shift shown below. (Use the terms: reversible reaction, forward reaction, reverse reaction, rate, products, and reactants.)
Describe La Chatelier’s Principle:
Fill in the table below for the reversible exothermic reaction. A plus sign indicates that the concentration of that substance was increased. A minus sign indicates that the concentration of that substance was decreased.
3 O2(g) + 2 H2S (g) ß à 2 H2O (g) + 2 SO2 (g) + heat
Stress / [O2] / [H2S] / DirectionOf Shift / [H2O] / [SO2]
Add
H2S / +
Remove
H2O / -
Increase
Temperature
Decrease
Pressure
In which trial(s) would the value of the equilibrium constant (Keq ) change? Would the value increase or decrease?
What would you do to maximize the production of SO2?
Consider this reaction at equilibrium:
2 HCl (g) ß à H2 (g) + Cl2 (g)
a. Write the equilibrium constant expression.
b. Calculate the numerical value for Keq if an analysis of a 1L flask at equilibrium gives the following results: [HCl] = 0.30 moles, [H2] = 1.2 moles, and [Cl2] = 0.60 moles.
c. Calculate the Keq at a different temperature if 0.40 moles of H2 is added to 0.40 moles of Cl2. There is no HCl present initially. At equilibrium, only the concentration of H2 could be determined. [H2] = 0.30
d. At the same temperature as problem c above, a chemist combines 0.100 M concentrations of all the species in the reaction. Will the reaction proceed predominantly to the left or right to reach equilibrium? Calculate the equilibrium concentrations of all species.