A chemist is studying the following equilibirum, which has the given equilibrium constant at a certain temperature: 30₂(g) 203 (g) K₂=1.x 10-6 He fills a reaction vessel at this temperature with 10. atm of oxygen gas. Use this data to answer the questions in the table below. Can you predict the equilibrium pressure of O3, using only the tools available to you within ALEKS? If you said yes, then enter the equilibrium pressure of O3 at right. Round your answer to 1 significant digit. yes no atm X Ś

A chemist is studying the following equilibirum, which has the given equilibrium constant at a certain temperature: 30₂(g) 203 (g) K₂=1.x 10-6 He fills a reaction vessel at this temperature with 10. atm of oxygen gas. Use this data to answer the questions in the table below. Can you predict the equilibrium pressure of O3, using only the tools available to you within ALEKS? If you said yes, then enter the equilibrium pressure of O3 at right. Round your answer to 1 significant digit. yes no atm X Ś

Introductory Chemistry: A Foundation

9th Edition

ISBN:9781337399425

Author:Steven S. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Donald J. DeCoste

Chapter17: Equilibrium

Section: Chapter Questions

Problem 51QAP: . For the reaction 3O2(g)2O3(g)The equilibrium constant, K, has the value 1.121054at a particular...

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Because carbonic acid undergoes a second ionization, the student in Exercise 12.39 is concerned that the hydrogen ion concentration she calculated is not correct. She looks up the equilibrium constant for the reaction HCO,-(aq) «=* H+(aq) + COf'(aq) Upon finding that the equilibrium constant for this reaction is 4.8 X 10“H, she decides that her answer in Exercise 12.39 is correct. Explain her reasoning. A student is simulating the carbonic acid—hydrogen carbonate equilibrium in a lake: H,CO,(aq) 5=6 H+(aq) + HCO,'(aq) K = 4.4 X 10'7She starts with 0.1000 A1 carbonic acid. W hat are the concentrations of all species at equilibrium?
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?
Chemical Equilibrium II Magnesium hydroxide. Mg(OH)2, is a white, partially soluble solid that is used in many antacids. The chemical equation for the dissolving of Mg(OH)2(s) in water is Mg(OH)2(s)Mg2+(aq)+2OH(aq) a Describe a simple experimental procedure that you could use to study this solubility equilibrium. In your experiment, how would you determine when the solution process has attained equilibrium? b Write the equilibrium-constant expression for this dissolving of magnesium hydroxide. c Suppose equilibrium has been established in a container of magnesium hydroxide in water, and you decide to add more solid Mg(OH)2. What would you expect to observe? What effect will this addition of Mg(OH)2 have on the concentrations of Mg2+(aq) and OH(aq)? d Say you haw prepared an equilibrium solution of Mg(OH)2 by adding pure solid Mg(OH)2 to water. If you know the concentration of OH(aq), can you determine the concentration of Mg2+(aq)? If not, what information do you need that will allow you to determine the answer? e You slowly add OH from another source (say, NaOH) to an equilibrium mixture of Mg(OH)2 and water. How do you expect the concentration of the Mg2+(aq) to change? What might you be able to observe happening to the Mg(OH)2(s) as you add the OH? f Next you remove some, but not all, of the Mg(OH)2(s) from the mixture. How will this affect the concentrations of the Mg2+(aq) and OH(aq)? g If someone hands you a container of Mg(OH)2(aq) and there is no solid Mg(OH)2 present, is this solution at equilibrium? If it is not at equilibrium, what could you add to or remove from the container that would give an equilibrium system? h Consider an individual OH(aq) ion in an Mg(OH)2 solution at equilibrium. If you could follow this ion over a long period of time, would you expect it always to remain as an OH(aq) ion, or could it change in some way?
The equilibrium between nitrogen monoxide, oxygen, and nitrogen dioxide may be expressed in the equation 2NO(g)+O2(g)2NO(g). Write the equilibrium constant expression for this equation. Then express the same equilibrium in at least two other ways, and write the equilibrium constant expression for each. Are the constants numerically equal? Cite some evidence to support your answer. Nitrogen monoxide and oxygen, both colorless gases, react to form reddish-brown nitrogen dioxide.
Methanol, a common laboratory solvent, poses a threat of blindness or death if consumed in sufficient amounts. Once in the body, the substance is oxidized to produce formaldehyde (embalming fluid) and eventually formic acid. Both of these substances are also toxic in varying levels. The equilibrium between methanol and formaldehyde can be described as follows: CH3OH(aq)H2CO(aq)+H2(aq) Assuming the value of K for this reaction is 3.7 1010, what are the equilibrium concentrations of each species if you start with a 1.24 M solution of methanol? What will happen to the concentration of methanol as the formaldehyde is further converted to formic acid?
At 2300 K the equilibrium constant for the formation of NO(g) is 1.7 103. N2(g) + O2(g) 2 NO(g) (a) Analysis shows that the concentrations of N2 and O2 are both 0.25 M, and that of NO is 0.0042 M under certain conditions. Is the system at equilibrium? (b) If the system is not at equilibrium, in which direction does the reaction proceed? (c) When the system is at equilibrium, what are the equilibrium concentrations?
Consider the following equilibrium: COBr2(g) CO(g) + Br2(g)Kc = 0.190 at 73 C (a) A 0.50 mol sample of COBr2 is transferred to a 9.50-L flask and heated until equilibrium is attained. Calculate the equilibrium concentrations of each species. (b) The volume of the container is decreased to 4.5 L and the system allowed to return to equilibrium. Calculate the new equilibrium concentrations. (Hint: The calculation will be easier if you view this as a new problem with 0.5 mol of COBr2 transferred to a 4.5-L flask.) (c) What is the effect of decreasing the container volume from 9.50 L to 4.50 L?
How does equilibrium represent the balancing of opposing processes? Give an example of an “equilibrium” encountered in everyday life, showing how the processes involved oppose each other.
Consider 0.200 mol phosphorus pentachloride sealed in a 2.0-L container at 620 K. The equilibrium constant, Kc, is 0.60 for PCl5(g) PCl3(g) + Cl2(g) Calculate the concentrations of all species after equilibrium has been reached.
At room temperature, the equilibrium constant Kc for the reaction 2 NO(g) ⇌ N2(g) + O2(g) is 1.4 × 1030. Is this reaction product-favored or reactant-favored? Explain your answer. In the atmosphere at room temperature the concentration of N2 is 0.33 mol/L, and the concentration of O2 is about 25% of that value. Calculate the equilibrium concentration of NO in the atmosphere produced by the reaction of N2 and O2. How does this affect your answer to Question 11?
A mixture of N2, H2, and NH3 is at equilibrium [according to the equationN2(g)+3H2(g)2NH3(g)] as depicted below: The volume is suddenly decreased (by increasing the external pressure) and a new equilibrium is established as depicted below: a. If the volume of the final equilibrium mixture is 1.00 L, determine the value of the equilibrium constant, K. for the reaction. Assume temperature is constant. b. Determine the volume of the initial equilibrium mixture assuming a final equilibrium volume of 1.00 L and assuming a constant temperature.
. For the reaction 3O2(g)2O3(g)The equilibrium constant, K, has the value 1.121054at 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 3.04102M, what is the concentration of O3(g) in the mixture’?