For the reaction K 1.3 102 at a given temperature. If the system at equilibrium is

Consider an equilibrium mixture of four chemicals (A, B, C, and D, all gases) reacting in a closed flask according to the following equation: a. You add more A to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer. b. You have the original set-up at equilibrium, and add more D to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer

The boxes shown on the following page represent a set of initial conditions for the reaction: Draw a quantitative molecular picture that shows what this system looks like after the reactants are mixed in one of the boxes and the system reaches equilibrium. Support your answer with calculations.

For the reaction H2 I2 2HI, consider two possibilities: (a) you add 0.5 mole of each reactant, allow the system to come to equilibrium, and then add 1 mol H2, and allow the system to reach equilibrium again, or (b) you add 1.5 mol H2 and 0.5 mol I2 and allow the system to come to equilibrium. Will the final equilibrium mixture be different for the two procedures? Explain.

Given the reaction A B C D, consider the following situations: a. You have 1.3 M A and 0.8 M B initially. b. You have 1.3 M A, 0.8 M B, and 0.2 M C initially. c. You have 2.0 M A and 0.8 M B initially. Order the preceding situations in terms of increasing equilibrium concentration of D and explain your order. Give the order in terms of increasing equilibrium concentration of B and explain.

Consider the reaction A B C D. A friend asks the following: I know we have been told that if a mixture of A, B, C, and D is in equilibrium and more A is added, more C and D will form. But how can more C and D form if we do not add more B? What do you tell your friend?

Consider the following statements: Consider the reaction A(g) B(g) C(g), for which at equilibrium [A]  2 M, [B]  1 M, and [C]  4 M. To a 1-L container of the system at equilibrium you add 3 moles of B. A possible equilibrium condition is [A]  1 M, [B]  3 M, and [C]  6 M, because in both cases, K  2. Indicate everything you think is correct in these statements, and everything that is incorrect. Correct the incorrect statements, and explain.

The value of the equilibrium constant, K, is dependent on which of the following? (There may be more than one answer.) a. the initial concentrations of the reactants b. the initial concentrations of the products c. the temperature of the system d. the nature of the reactants and products Explain.

You are browsing through the Handbook of Hypothetical Chemistry when you come across a solid that is reported to have a Ksp value of zero in water at 25 C. What does this mean?

What do you suppose happens to the Ksp value of a solid as the temperature of the solution changes? Consider both increasing and decreasing temperatures, and explain your answer.

Consider an equilibrium mixture consisting of H2O(g), CO(g), H2(g), and CO2(g) reacting in a closed vessel according to the equation H2O(g) CO(g) H2(g) CO2(g) a. You add more H2O to the flask. How does the new equilibrium concentration of each chemical compare to its original equilibrium concentration after equilibrium is reestablished? Justify your answer. b. You add more H2 to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer.

Equilibrium is microscopically dynamic but macroscopically static. Explain what this means.

In Section 17.3 of your text, it is mentioned that equilibrium is reached in a closed system. What is meant by the term closed system, and why is it necessary for a system to reach equilibrium? Explain why equilibrium is not reached in an open system.

Explain why the development of a vapor pressure above a liquid in a closed container represents an equilibrium. What are the opposing processes? How do we recognize when the system has reached a state of equilibrium.

Consider the figure below in answering the following questions. a. What does a catalyst do to a chemical reaction? b. Which of the pathways in the figure is the catalyzed reaction pathway? How do you know? c. What is represented by the double-headed arrow?

Write the equilibrium expression for each of the following reactions. a. C2H6(g) Cl2(g) C2H5Cl(s) HCl(g) b. 4NH3(g) 5O2(g) 4NO(g) 6H2O(g) c. PCl5(g) PCl3(g) Cl2(g)

Write the equilibrium expression for each of the following reactions. a. H2(g) CO2(g) H2O(g) CO(g) b. 2N2O(g) O2(g) 4NO(g) c. CO(g) 2H2(g) CH3OH(g)

Write the equilibrium expression for each of the following reactions. a. NO2(g) ClNO(g) ClNO2(g) NO(g) b. Br2(g) 5F2(g) 2BrF5(g) c. 4NH3(g) 6NO(g) 5N2(g) 6H2O(g)

Write the equilibrium expression for each of the following reactions. a. CO(g) 2H2(g) CH3OH(g) b. 2NO2(g) 2NO(g) O2(g) c. P4(g) 6Br2(g) 4PBr3(g)

Suppose that for the reaction PCl5(g) PCl3(g) Cl2(g) it is determined, at a particular temperature, that the equilibrium concentrations are [PCl5(g)]  0.0711 M, [PCl3(g)]  0.0302 M, and [Cl2(g)]  0.0491 M. Calculate the value of K for the reaction at this temperature.

Suppose that for the reaction COCl2(g) CO(g) Cl2(g) it is determined, at a particular temperature, that the equilibrium concentrations are [COCl2(g)]  0.00103 M, [CO(g)]  0.0345 M, and [Cl2(g)]  0.0219 M. Calculate the value of K for the reaction at this temperature.

At high temperatures, elemental nitrogen and oxygen react with each other to form nitrogen monoxide. Suppose the system is analyzed at a particular temperature, and the equilibrium concentrations are found to be [N2]  0.041 M, [O2]  0.0078 M, and [NO]  4.7  104 M. Calculate the value of K for the reaction.

Suppose that for the reaction 2N2O(g) O2(g) 4NO(g) it is determined, at a particular temperature, that the equilibrium concentrations are [NO(g)]  0.00341 M, [N2O(g)]  0.0293 M, and [O2(g)]  0.0325 M. Calculate the value of K for the reaction at this temperature.

What is a homogeneous equilibrium system? Give an example of a homogeneous equilibrium reaction. What is a heterogeneous equilibrium system? Write two chemical equations that represent heterogeneous equilibria.

Explain why the position of a heterogeneous equilibrium does not depend on the amounts of pure solid or pure liquid reactants or products present.

Write the equilibrium expression for each of the following heterogeneous equilibria. a. P4(s) 6F2(g) 4PF3(g) b. Xe(g) 2F2(g) XeF4(s) c. 2SiO(s) 4Cl2(g) 2SiCl4(l) O2(g)

Write the equilibrium expression for each of the following heterogeneous equilibria. a. Fe(s) H2O(g) FeO(s) H2(g) b. 4Al(s) 3O2(g) 2Al2O3(s) c. CH4(g) 4Cl2(g) CCl4(l) 4HCl(g)

Write the equilibrium expression for each of the following heterogeneous equilibria. a. C(s) H2O(g) H2(g) CO(g) b. H2O(l) H2O(g) c. 4B(s) 3O2(g) 2B2O3(s)

Write the equilibrium expression for each of the following heterogeneous equilibria. a. CS2(g) 3Cl2(g) CCl4(l) S2Cl2(g) b. Xe(g) 3F2(g) XeF6(s) c. 4Fe(s) 3O2(g) 2Fe2O3(s)

In your own words, describe what Le Chteliers principle tells us about how we can change the position of a reaction system at equilibrium.

Consider the reaction Suppose the system is already at equilibrium, and then an additional mole of CO(g) is injected into the system at constant temperature. Does the amount of CO2(g) in the system increase or decrease? Does the value of K for the reaction change?

For an equilibrium involving gaseous substances, what effect, in general terms, is realized when the volume of the system is decreased?

What is the effect on the equilibrium position if an endothermic reaction is carried out at a higher temperature? Does the net amount of product increase or decrease? Does the value of the equilibrium constant change if the temperature is increased?

For the reaction system C(s) H2O(g) H2(g) CO(g) which has already reached a state of equilibrium, predict the effect that each of the following changes will have on the position of the equilibrium. Tell whether the equilibrium will shift to the right, will shift to the left, or will not be affected. a. The pressure of hydrogen is increased by injecting an additional mole of hydrogen gas into the reaction vessel. b. Carbon monoxide gas is removed as it forms by use of a chemical absorbent or scrubber. c. An additional amount of solid carbon is added to the reaction vessel

For the reaction system P4(s) 6F2(g) 4PF3(g) which has already reached a state of equilibrium, predict the effect that each of the following changes will have on the position of the equilibrium. Tell whether the equilibrium will shift to the right, will shift to the left, or will not be affected. a. Additional fluorine gas is added to the system. b. Additional phosphorus is added to the system. c. Additional phosphorus trifluoride is added to the system.

Suppose the reaction system has already reached equilibrium. Predict the effect of each of the following changes on the position of the equilibrium. Tell whether the equilibrium will shift to the right, will shift to the left, or will not be affected. a. Any liquid water present is removed from the system. b. CO2 is added to the system by dropping a chunk of dry ice into the reaction vessel. c. The reaction is performed in a metal cylinder fitted with a piston, and the piston is compressed to decrease the total volume of the system. d. Additional O2(g) is added to the system from a cylinder of pure O2.

Consider the general reaction 2A(g) B(s) C(g) 3D(g) H  115 kJ/mol which has already come to equilibrium. Predict whether the equilibrium will shift to the left, will shift to the right, or will not be affected if the changes indicated below are made to the system. a. Additional B(s) is added to the system. b. C(g) is removed from the system as it forms. c. The volume of the system is decreased by a factor of 2. d. The temperature is increased

Hydrogen gas and chlorine gas in the presence of light react explosively to form hydrogen chloride H2(g) Cl2(g) 2HCl(g) The reaction is strongly exothermic. Would an increase in temperature for the system tend to favor or disfavor the production of hydrogen chloride?

For the general reaction A(g) B(g) heat C(g) would an increase in temperature tend to favor the forward or the reverse process? Why?

The reaction is exothermic in the forward direction. Will an increase in temperature shift the position of the equilibrium toward reactants or products?

The reaction is strongly endothermic. Will an increase in temperature shift the equilibrium position toward products or toward reactants?

Plants synthesize the sugar dextrose according to the following reaction by absorbing radiant energy from the sun (photosynthesis). Will an increase in temperature tend to favor or discourage the production of C6H12O6(s)?

Consider the exothermic reaction CO(g) 2H2(g) CH3OH(l) Predict three changes that could be made to the system that would increase the yield of product over that produced by a system in which no change was made.

Suppose a reaction has the equilibrium constant K  1.3  108 . What does the magnitude of this constant tell you about the relative concentrations of products and reactants that will be present once equilibrium is reached? Is this reaction likely to be a good source of the products?

Suppose a reaction has the equilibrium constant K  1.7  108 at a particular temperature. Will there be a large or small amount of unreacted starting material present when this reaction reaches equilibrium? Is this reaction likely to be a good source of products at this temperature?

For the reaction Br2(g) 5F2(g) 2BrF5(g) the system at equilibrium at a particular temperature is analyzed, and the following concentrations are found: [BrF5(g)]  1.01  109 M, [Br2(g)]  2.41  102 M, and [F2(g)]  8.15  102 M. Calculate the value of K for the reaction at this temperature.

Consider the reaction SO2(g) NO2(g) SO3(g) NO(g) Suppose it is found at a particular temperature that the concentrations in the system at equilibrium are as follows: [SO3(g)]  4.99  105 M, [NO(g)]  6.31  107 M, [SO2(g)]  2.11  102 M, and [NO2(g)]  1.73  103 M. Calculate the value of K for the reaction at this temperature.

For the reaction it is found at equilibrium at a certain temperature that the concentrations are [CO(g)]  2.7  104 M, [O2(g)]  1.9  103 M, and [CO2(g)]  1.1  101 M. Calculate K for the reaction at this temperature.

For the reaction CO2(g) H2(g) CO(g) H2O(g) the equilibrium constant, K, has the value 5.21  103 at a particular temperature. If the system is analyzed at equilibrium at this temperature, it is found that [CO(g)]  4.73  103 M, [H2O(g)]  5.21  103 M, and [CO2(g)]  3.99  102 M. What is the equilibrium concentration of H2(g) in the system?

The equilibrium constant for the reaction has the value 2.1  103 at a particular temperature. When the system is analyzed at equilibrium at this temperature, the concentrations of both H2(g) and F2(g) are found to be 0.0021 M. What is the concentration of HF(g) in the equilibrium system under these conditions?

For the reaction K  2.4  103 at a given temperature. At equilibrium it is found that [H2O(g)]  1.1  101 M and [H2(g)]  1.9  102 M. What is the concentration of O2(g) under these conditions?

For the reaction 3O2(g) 2O3(g) The equilibrium constant, K, has the value 1.12  1054 at a particular temperature. a. What does the very small equilibrium constant indicate about the extent to which oxygen gas, O2(g), is converted to ozone gas, O3(g), at this temperature? b. If the equilibrium mixture is analyzed and [O2(g)] is found to be 3.04  102 M, what is the concentration of O3(g) in the mixture?

For the reaction the equilibrium constant K has the value 8.1  103 at a particular temperature. If the concentration of NO2(g) is found to be 0.0021 M in the equilibrium system, what is the concentration of N2O4(g) under these conditions?

Explain how the dissolving of an ionic solute in water represents an equilibrium process.

What is the special name given to the equilibrium constant for the dissolving of an ionic solute in water?

Why does the amount of excess solid solute present in a solution not affect the amount of solute that ultimately dissolves in a given amount of solvent?

Which of the following will affect the total amount of solute that can dissolve in a given amount of solvent? a. The solution is stirred. b. The solute is ground to fine particles before dissolving. c. The temperature changes.

Write the balanced chemical equation describing the dissolving of each of the following sparingly soluble salts in water. Write the expression for Ksp for each process. a. AgIO3(s) c. Zn3(PO4)2(s) b. Sn(OH)2(s) d. BaF2(s)

Write the balanced chemical equation describing the dissolving of each of the following sparingly soluble salts in water. Write the expression for Ksp for each process. a. Bi2S3(s) c. Co(OH)3(s) b. Ca(OH)2(s) d. Cu2S(s)

Ksp for copper(II) hydroxide, Cu(OH)2, has a value 2.2  1020 at 25 C. Calculate the solubility of copper(II) hydroxide in mol/L and g/L at 25 C.

Ksp for magnesium carbonate, MgCO3, has a value 3.5  108 at 25 C. Calculate the solubility of magnesium carbonate in mol/L and g/L at 25 C.

A saturated solution of nickel(II) sulfide contains approximately 3.6  104 g of dissolved NiS per liter at 20 C. Calculate the solubility product Ksp for NiS at 20 C.

Most hydroxides are not very soluble in water. For example, Ksp for nickel(II) hydroxide, Ni(OH)2, is 2.0  1015 at 25 C. How many grams of nickel(II) hydroxide dissolve per liter at 25 C?

The solubility product constant, Ksp, for calcium carbonate at room temperature is approximately 3.0  109 . Calculate the solubility of CaCO3 in grams per liter under these conditions.

Calcium sulfate, CaSO4, is only soluble in water to the extent of approximately 2.05 g/L at 25 C. Calculate Ksp for calcium sulfate at 25C.

Approximately 1.5  103 g of iron(II) hydroxide, Fe(OH)2(s), dissolves per liter of water at 18 C. Calculate Ksp for Fe(OH)2(s) at this temperature.

Chromium(III) hydroxide dissolves in water only to the extent of 8.21  105 M at 25 C. Calculate Ksp for Cr(OH)3 at this temperature.

Magnesium fluoride dissolves in water to the extent of 8.0  102 g/L at 25 C. Calculate the solubility of MgF2(s) in moles per liter, and calculate Ksp for MgF2 at 25 C.

Lead(II) chloride, PbCl2(s), dissolves in water to the extent of approximately 3.6  102 M at 20 C. Calculate Ksp for PbCl2(s), and calculate its solubility in grams per liter.

Mercury(I) chloride, Hg2Cl2, was formerly administered orally as a purgative. Although we usually think of mercury compounds as highly toxic, the Ksp of mercury(I) chloride is small enough (1.3  1018) that the amount of mercury that dissolves and enters the bloodstream is tiny. Calculate the concentration of mercury(I) ion present in a saturated solution of Hg2Cl2.

The solubility product of iron(III) hydroxide is very small: Ksp  4  1038 at 25 C. A classical method of analysis for unknown samples containing iron is to add NaOH or NH3. This precipitates Fe(OH)3, which can then be filtered and weighed. To demonstrate that the concentration of iron remaining in solution in such a sample is very small, calculate the solubility of Fe(OH)3 in moles per liter and in grams per liter.

Before two molecules can react, chemists envision that the molecules must first collide with one another. Is collision among molecules the only consideration for the molecules to react with one another?

Why does an increase in temperature favor an increase in the speed of a reaction?

The minimum energy required for molecules to react with each other is called the energy.

A(n) speeds up a reaction without being consumed.

Equilibrium may be defined as the of two processes, one of which is the opposite of the other.

When a chemical system has reached equilibrium, the concentrations of all reactants and products remain with time.

What does it mean to say that all chemical reactions are, to one extent or another, reversible?

What does it mean to say that chemical equilibrium is a dynamic process?

At the point of chemical equilibrium, the rate of the forward reaction the rate of the reverse reaction.

Equilibria involving reactants or products in more than one state are said to be .

According to Le Chteliers principle, when a large excess of a gaseous reactant is added to a reaction system at equilibrium, the amounts of products .

Addition of an inert substance (one that does not participate in the reaction) does not change the of an equilibrium.

When the volume of a vessel containing a gaseous equilibrium system is decreased, the of the gaseous substances present is initially increased.

Why does increasing the temperature for an exothermic process tend to favor the conversion of products back to reactants?

What is meant by the solubility product for a sparingly soluble salt? Choose a sparingly soluble salt and show how the salt ionizes when dissolved in water, and write the expression for its solubility product.

For a given reaction at a given temperature, the special ratio of products to reactants defined by the equilibrium constant is always equal to the same number. Explain why this is true, no matter what initial concentrations of reactants (or products) may have been taken in setting up an experiment.

Many sugars undergo a process called mutarotation, in which the sugar molecules interconvert between two isomeric forms, finally reaching an equilibrium between them. This is true for the simple sugar glucose, C6H12O6, which exists in solution in isomeric forms called alpha-glucose and beta-glucose. If a solution of glucose at a certain temperature is analyzed, and it is found that the concentration of alpha-glucose is twice the concentration of beta-glucose, what is the value of K for the interconversion reaction?

Suppose K  4.5  103 at a certain temperature for the reaction If it is found that the concentration of PCl5 is twice the concentration of PCl3, what must be the concentration of Cl2 under these conditions?

For the reaction the equilibrium constant K has the form K  [CO2]. Using a handbook to find density information about CaCO3(s) and CaO(s), show that the concentrations of the two solids (the number of moles contained in 1 L of volume) are constant.

As you know from Chapter 7, most metal carbonate salts are sparingly soluble in water. Below are listed several metal carbonates along with their solubility products, Ksp. For each salt, write the equation showing the ionization of the salt in water, and calculate the solubility of the salt in mol/L. Salt Ksp BaCO3 5.1  109 CdCO3 5.2  1012 CaCO3 2.8  109 CoCO3 1.5  1013

Teeth and bones are composed, to a first approximation, of calcium phosphate, Ca3(PO4)2(s). The Ksp for this salt is 1.3  1032 at 25 C. Calculate the concentration of calcium ion in a saturated solution of Ca3(PO4)2.

Under what circumstances can we compare the solubilities of two salts by directly comparing the values of their solubility products?

How does the collision model account for the fact that a reaction proceeds faster when the concentrations of the reactants are increased?

How does an increase in temperature result in an increase in the number of successful collisions between reactant molecules? What does an increase in temperature mean on a molecular basis?

Explain why the development of a vapor pressure above a liquid in a closed container represents an equilibrium. What are the opposing processes? How do we recognize when the system has reached a state of equilibrium?

Write the equilibrium expression for each of the following reactions. a. H2(g) Br2(g) 2HBr(g) b. 2H2(g) S2(g) 2H2S(g) c. H2(g) C2N2(g) 2HCN(g)

Write the equilibrium expression for each of the following reactions. a. 2O3(g) 3O2(g) b. CH4(g) 2O2(g) CO2(g) 2H2O(g) c. C2H4(g) Cl2(g) C2H4Cl2(g)

At high temperatures, elemental bromine, Br2, dissociates into individual bromine atoms. Suppose that in an experiment at 2000 C, it is found that [Br2]  0.97 M and [Br]  0.034 M at equilibrium. Calculate the value of K.

Gaseous phosphorus pentachloride decomposes according to the reaction The equilibrium system was analyzed at a particular temperature, and the concentrations of the substances present were determined to be [PCl5]  1.1  102 M, [PCl3]  0.325 M, and [Cl2]  3.9  103 M. Calculate the value of K for the reaction.

Write the equilibrium expression for each of the following heterogeneous equilibria. a. 4Al(s) 3O2(g) 2Al2O3(s) b. NH3(g) HCl(g) NH4Cl(s) c. 2Mg(s) O2(g) 2MgO(s)

Write the equilibrium expression for each of the following heterogeneous equilibria. a. P4(s) 5O2(g) P4O10(s) b. CO2(g) 2NaOH(s) Na2CO3(s) H2O(g) c. NH4NO3(s) N2O(g) 2H2O(g)

What is the effect on the position of a reaction system at equilibrium when an exothermic reaction is performed at a higher temperature? Does the value of the equilibrium constant change in this situation?

Suppose the reaction system has already reached equilibrium. Predict the effect of each of the following changes on the position of the equilibrium. Tell whether the equilibrium will shift to the right, will shift to the left, or will not be affected. a. Additional oxygen is injected into the system. b. NO2 is removed from the reaction vessel. c. 1.0 mole of helium is injected into the system.

The reaction PCl3(l) Cl2(g) 34 PCl5(s) liberates 124 kJ of energy per mole of PCl3 reacted. Will an increase in temperature shift the equilibrium position toward products or toward reactants?

For the process it is found that the equilibrium concentrations at a particular temperature are [H2]  1.4 M, [CO2]  1.3 M, [CO]  0.71 M, and [H2O]  0.66 M. Calculate the equilibrium constant K for the reaction under these conditions.

For the reaction K  1.3  102 at a given temperature. If the system at equilibrium is analyzed and the concentrations of both N2 and H2 are found to be 0.10 M, what is the concentration of NH3 in the system?

The equilibrium constant for the reaction has the value 9.2  106 at a particular temperature. The system is analyzed at equilibrium, and it is found that the concentrations of NOCl(g) and NO(g) are 0.44 M and 1.5  103 M, respectively. What is the concentration of Cl2(g) in the equilibrium system under these conditions?

As you learned in Chapter 7, most metal hydroxides are sparingly soluble in water. Write balanced chemical equations describing the dissolving of the following metal hydroxides in water. Write the expression for Ksp for each process. a. Cu(OH)2(s) c. Ba(OH)2(s) b. Cr(OH)3(s) d. Sn(OH)2(s)

The three common silver halides (AgCl, AgBr, and AgI) are all sparingly soluble salts. Given the values for Ksp for these salts below, calculate the concentration of silver ion, in mol/L, in a saturated solution of each salt. Silver Halide Ksp AgCl 1.8  1010 AgBr 5.0  1013 AgI 8.3  1017

Approximately 9.0  104 g of silver chloride, AgCl(s), dissolves per liter of water at 10 C. Calculate Ksp for AgCl(s) at this temperature.

Mercuric sulfide, HgS, is one of the least soluble salts known, with Ksp  1.6  1054 at 25 C. Calculate the solubility of HgS in moles per liter and in grams per liter

Approximately 0.14 g of nickel(II) hydroxide, Ni(OH)2(s), dissolves per liter of water at 20 C. Calculate Ksp for Ni(OH)2(s) at this temperature.

For the reaction N2(g) 3H2(g) S 2NH3(g), list the types of bonds that must be broken and the type of bonds that must form for the chemical reaction to take place.

What does the activation energy for a reaction represent? How is the activation energy related to whether a collision between molecules is successful?

What are the catalysts in living cells called? Why are these biological catalysts necessary?

When a reaction system has reached chemical equilibrium, the concentrations of the reactants and products no longer change with time. Why does the amount of product no longer increase, even though large concentrations of the reactants may still be present?

Ammonia, a very important industrial chemical, is produced by the direct combination of the elements under carefully controlled conditions. Suppose, in an experiment, that the reaction mixture is analyzed after equilibrium is reached and it is found, at a particular temperature, that [NH3(g)]  0.34 M, [H2(g)]  2.1  103 M, and [N2(g)]  4.9  104 M. Calculate the value of K at this temperature.

Write the equilibrium expression for each of the following heterogeneous equilibria. a. 2LiHCO3(s) Li2CO3(s) H2O(g) CO2(g) b. PbCO3(s) PbO(s) CO2(g) c. 4Al(s) 3O2(g) 2Al2O3(s)

Suppose a reaction has the equilibrium constant K  4.5  106 at a particular temperature. If an experiment is set up with this reaction, will there be large relative concentrations of products present at equilibrium? Is this reaction useful as a means of producing the products? How might the reaction be made more useful?