Final answer:
To address the equilibrium chemistry question, one constructs a table with initial, changes, and final partial pressures, writes the equilibrium constant expression (Kp), and calculates the partial pressure of gases like NO. Subsequent Kp values should be confirmed against provided equilibrium partial pressures.
Step-by-step explanation:
Equilibrium Constant Expression
To construct a table and enter the initial partial pressures, changes in partial pressures during the reaction, and final partial pressures of all substances involved in the equilibrium reaction, we must first identify the reaction. If this concerns the formation of nitrogen monoxide (NO) from nitrogen (N₂) and oxygen (O₂), the reaction is as follows:
N₂(g) + O₂(g) ⇒ 2 NO(g)
This reaction would involve calculating the changes in partial pressures (ΔP) according to stoichiometry. If represented by Δx as the change for NO, then the changes for N₂ and O₂ would be -Δx/2 due to the mole ratios. The final partial pressures would be the initial pressures plus the change for each gas at equilibrium.
Equilibrium Equation
The equilibrium constant expression for the reaction in terms of partial pressures (Kp) is:
Kp = (PNO)2 / (PN₂)(PO₂)
By substituting the values from the table into this equation, we can solve for the change in concentration (Δx).
Calculating Partial Pressures
To calculate the partial pressure of NO, we substitute the known equilibrium partial pressures and Kp into the equilibrium equation. By rearranging the equation, we can solve for the unknown partial pressure of NO. It's important to confirm that the calculated Kp values are consistent with the equilibrium partial pressures provided.