1. During one step in the synthesis of nitric acid, nitrogen dioxide (NO2) is in equilibrium withdinitrogen tetroxide (N204) at 60°C.N204(g) = 2N02 (g)0.350 mol of N2O4 was placed in a 2.00 L vessel. When equilibrium was achieved at 60.0 C,0.120 mol of NO2 was present. Calculate the value of the equilibrium constant at thistemperature.

We have to predict the equilibrium constant.

1. During one step in the synthesis of nitric acid, nitrogen dioxide (NO2) is in equilibrium with dinitrogen tetroxide (N204) at 60°C. N204(g) = 2NO2(g) 0.350 mol of N2O4 was placed in, a 2.00 L vessel. When equilibrium was achieved at 60.0 C, 0.120 mol of NO2 was present. Calculate the value of the equilibrium constant at this temperature.

1. During one step in the synthesis of nitric acid, nitrogen dioxide (NO2) is in equilibrium with dinitrogen tetroxide (N204) at 60°C. N204(g) = 2NO2(g) 0.350 mol of N2O4 was placed in, a 2.00 L vessel. When equilibrium was achieved at 60.0 C, 0.120 mol of NO2 was present. Calculate the value of the equilibrium constant at this temperature.

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

See similar textbooks

Similar questions

The following reaction is investigated (assume an ideal gas mixture): 2N20(g) + N2H4(g) 3N2(g) + 2H2O(g) Initially there are 0.100 mol of N20 and 0.25 mol of N2H4, in a 10.0-L container. If there are 0.059 mol of N20 at equilibrium, how many moles of N2 are present at equilibrium? O 4.1 x 10^-2 1.2 x 10^-1 6.2 x 10A-2 O 2.1 x 10^-2 O none of these
Dinitrogen tetraoxide and nitrogen dioxide are two gases that exist in equilibrium at a range of temperatures. NO2 is a reddish brown gas while N2O4 is colorless. ... At High Temperature the red color is strong. ... ... At Low Temperature the gas has less color. ... If we represent the equilibrium as:... 2 NO2(g) N2O4(g)We can conclude that: fill in the blank 1 1. This reaction is: A. Exothermic B. Endothermic C. Neutral D. More information is needed to answer this question. fill in the blank 2 2. When the temperature is decreased the equilibrium constant, K: A. Increases B. Decreases C. Remains the same D. More information is needed to answer this question. fill in the blank 3 3. When the temperature is decreased the equilibrium concentration of NO2: A. Increases B. Decreases C. Remains the same D. More information is needed to answer this question.
The following gases are in equilibrium in a sealed glass tube. 2 NO2 (g) = N2O4 (g) (brown) (colorless) Increasing the temperature makes the gas in the tube more intensely brown in color. Decreasing the temperature has the opposite effect. Based on this observation, the reaction shown must be... O neither exothermic nor endothermic O endothermic exothermic
= KINETICS AND EQUILIBRIUM Setting up a reaction table Suppose a 500. mL flask is filled with 1.7 mol of N₂ and 0.10 mol of NH3. This reaction becomes possible: N₂(g) + 3H₂(g)2NH, (g) Complete the table below, so that it lists the initial molarity of each compound, the change in molarity of each compound due to the reaction, and the equilibrium molarity of each compound after the reaction has come to equilibrium. Use x to stand for the unknown change in the molarity of N₂. You can leave out the M symbol for molarity. initial change equilibrium N₂ 0 0 H₂ 0 0 0 NH3 0 0 00 1/5 X
At a certain temperature, 0.900 mol SO3 is placed in a 4.50 L container. 2SO3(g)↽−−⇀2SO2(g)+O2(g) At equilibrium, 0.140 mol O2 is present. Calculate ?c.
The reaction N2O4(g) → 2NO2(g) has the equilibrium constant KP=0.450 at 349 K. If the partial pressures of N2O4 and NO2 in a certain mixture are 320. torr and 12 torr, respectively, which of the following statements is true? A. N2O4 and NO2 are in equilibrium. B. N2O4 and NO2 are not in equilibrium, because PNO2 < PN2O4. c. C. More N2O4 must be formed to reach equilibrium. D. N2O4 and NO2 are not in equilibrium, but additional information is needed to predict how the two gases will behave so they eventually reach equilibrium. E. Certain amount of N2O4 must decompose to reach equilibrium.
A sample of H,S(g) at 0.495 atm is placed in a container and the following reaction reaches equilibrium. What is the total pressure of all gases in the container? 2H2S(g) = 2H2(g) + S2(g) K, = 3.85 x 10-7 0.497 atm 0.215 atm 0.793 atm 0.624 atm 0.387 atm O O O
For the equilibrium reaction. 2IBr (g) I2 (g) + Br2 (g) Kc=.0085. If .025 M of IBr is introduced to an empty flask and allowed to reach equilibrium, calculate the final concentrations of all components. Consider the decomposition reaction at 555 K 4POCl3 (g) P4 (g) + 2O2 (g) + 6Cl2 (g) If .450 atm of POCl3 is introduced to an otherwise empty flask and the reaction is allowed to reach equilibrium, the final total pressure is .850 atm. Find Kp and Kc.
Nitrogen dioxide is one of the many oxides of nitrogen (often collectively called "NOx") that are of interest to atmospheric chemistry. It can react with itself to form another form of NOx, dinitrogen tetroxide. A chemical engineer studying this reaction fills a 125.L tank at 13.°C with 16.mol of nitrogen dioxide gas. He then raises the temperature considerably, and when the mixture has come to equilibrium determines that it contains 6.9mol of nitrogen dioxide gas. The engineer then adds another 8.0mol of nitrogen dioxide, and allows the mixture to come to equilibrium again. Calculate the moles of dinitrogen tetroxide after equilibrium is reached the second time. Round your answer to 2 significant digits.