Consider the following reaction at equilibrium at total pressure that we will call P1: NO2(g) + CO(g) = NO(g) + Suppose the volume of the system is compressed to 1/2 its initial volume and the equilibrium is reestablished at the same temperature. The new equilibrium total CO2(g) pressure will be: two times P1- four times P1- eight times P1. none of the above.

Consider the following reaction at equilibrium at total pressure that we will call P1: NO2(g) + CO(g) = NO(g) + Suppose the volume of the system is compressed to 1/2 its initial volume and the equilibrium is reestablished at the same temperature. The new equilibrium total CO2(g) pressure will be: two times P1- four times P1- eight times P1. none of the above.

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

Suppose a 500.mL flask is filled with 1.0mol of Cl2 , 1.3mol of CHCl3 and 0.50mol of CCl4 . The following reaction becomes possible: +Cl2gCHCl3g +HClgCCl4g The equilibrium constant K for this reaction is 1.32 at the temperature of the flask. Calculate the equilibrium molarity of CHCl3 . Round your answer to two decimal places.
Consider the following equilibrium: NH3(g) + HCl(g) NH4CI(s) Suppose a vessel containing NH4Cl(s), NH3(g) and HCl(g) is at equilibrium. If the volume of the vessel is instantaneously doubled while keeping the temperature constant, when a new equilibrium is reached, which of the following statements is incorrect? The partial pressures of NH3 and HCI in the vessel remain unchanged. The total pressure is halved. O The value of Kp remains unchanged. O The number of moles of NH4Cl decreases. O The amount of NH3 and HCI doubles.
Suppose a 250. mL flask is filled with 0.70 mol of NO₂, 0.20 mol of NO and 1.9 mol of CO₂. The following reaction becomes possible: NO₂ (g) + CO (g) NO(g) + CO₂(g) The equilibrium constant K for this reaction is 0.715 at the temperature of the flask. Calculate the equilibrium molarity of NO₂. Round your answer to two decimal places.
Consider the following equilibrium reaction:H2(g) + I2(g) ↔ 2HI(g)If the following equilibrium concentrations are observed: [H2] = 4.10 x 10-1M, [I2] = 5.54 x 10-2M, [HI] = 3.50 x 10-2M; what is the equilibrium constant
Suppose a 500. mL flask is filled with 0.30 mol of SO, and 1.2 mol of SO3. This reaction becomes possible: 2SO, (g) +O,(g) 2SO3 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 O,. You can leave out the M symbol for molarity. so, SO2 O2 initial change equilibrium
Write the equilibrium constant expression. Solid carbon and carbon dioxide gas react and form carbon monoxide: C(s) + CO₂(g) = 2CO(g)
Nitrogen and hydrogen react to form ammonia, like this: N,(g) + 3 H,(g) → 2 NH,(g) Also, a chemist finds that at a certain temperature the equilibrium mixture of nitrogen, hydrogen, and ammonia has the following composition: compound pressure at equilibrium N2 82.3 atm H2 38.7 atm NH3 73.0 atm Calculate the value of the equilibrium constant K for this reaction. Round your answer to 2 significant digits.
N2(g) and O2(g) can exist in equilibrium with NO(g), as shown below. The equilibrium constant at 25.0°C is 4.8 x 10-31. If initially there are 1.35 mol of nitrogen and 0.60 mol of oxygen in a 2.00 L vessel, find the equilibrium concentrations of each species. N2(g) + O2(g) → 2NO(g)
Consider the following reaction at equilibrium. What effect will removing some SO2 have on the system? SO2(g) + NO2(g) = SO3(g) + NO(g) The pressure of SO3 will increase The equilibrium constant will decrease. No change will occur since SO2 is not included in the equilibrium expression. The pressure of NO2 will increase. The reaction will shift to decrease the pressure.
When nitrogen gas reacts with oxygen gas, nitrogen monoxide gas forms: N2(g) + O2(g) ⇋2NO(g) Initially, 0.40 mol of each reactant is placed into a 2.0 L container and equilibrium is established. If Keq=50.1 for the reaction, what are the equilibrium concentrations of all reactants and products? What two conditions are required, if we want to move reaction forward.
Consider the following exothermic reaction system at equilibrium: N2(g) + 3H2(g) ⇒ 2NH3(g) Predict how the following changes affect the number of moles of each component of the system after equilibrium is reestablished by completing the table below: Note: Use I for increase, D for decrease, NC for no change Add N2(g) Remove H2(g) Add NH3(9) Add Ne(g) Increase the temperature (constant P) Decrease the volume (constant T) Add a catalyst Submit NH3 N₂ H2