Introduced into a 1.50L container is 0.100mol PCl5(g). The flask is held at 227degreesC until equilibrium is established. What are the partial pressures and the total pressures of the gases in the flask at the equilibrium. PCl5(g)<-->PCl3(g)+Cl2(g)
I got a total pressure of 2.74 atm by plugging in the values into PV=nRT. THe mols I used were 0.1 for PCl5.
My partial pressures were 0.9atm for the other two gases.
Do you have a Kp or Kc or equilibrium concns of anything? The 2.74 atm is correct for the pressure of PCl5 at the start but is that the pressure at equilibrium? How did you arrive at 0.9atm for PCl3 and Cl2.
To find the partial pressures of the gases at equilibrium, you need to use the ideal gas law equation and consider the stoichiometry of the reaction.
First, let's determine the number of moles of PCl5 that will react. Since PCl5(g) decomposes into PCl3(g) and Cl2(g) in a 1:1 ratio, initially there will be 0.100 mol of PCl5.
When the reaction reaches equilibrium, let's assume that x moles of PCl5 have reacted to form x moles of PCl3 and x moles of Cl2. Therefore, the amount of PCl5 remaining will be (0.100 - x) mol.
Now, we can write the expression for the equilibrium constant (Kp) in terms of partial pressures:
Kp = (P_PCl3 * P_Cl2) / P_PCl5
By knowing the value of Kp, we can calculate the partial pressures of PCl3 and Cl2.
The total pressure of the gases at equilibrium will be the sum of the partial pressures of each gas.
To solve for the equilibrium partial pressures and total pressure, follow these steps:
Step 1: Calculate the equilibrium constant (Kp)
Since the equation is given as PCl5(g) <--> PCl3(g) + Cl2(g), the equilibrium constant expression will be:
Kp = (P_PCl3 * P_Cl2) / P_PCl5
Step 2: Calculate the equilibrium partial pressures
Let's assume x mol of PCl5 decomposes at equilibrium. Therefore, the molar concentrations of PCl3 and Cl2 will also be x mol.
Now, the partial pressure of PCl3 will be (x/V), where V is the volume (1.50L in this case).
Similarly, the partial pressure of Cl2 will also be (x/V).
Step 3: Calculate the remaining PCl5 moles at equilibrium
Initially, you had 0.100 mol of PCl5. Since x moles of PCl5 decompose, the remaining moles will be (0.100 - x).
Step 4: Calculate the total pressure at equilibrium
To find the total pressure at equilibrium, we add the partial pressures of all the gases.
Total pressure = P_PCl3 + P_Cl2 + P_PCl5
Step 5: Substitute the values into the ideal gas law equation
Use the ideal gas law equation, PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the ideal gas constant, and T is the temperature in Kelvin.
For PCl3, the pressure is (x/V), the volume is 1.50L, the number of moles is x, and the temperature is 227 degrees Celsius (convert to Kelvin).
For Cl2, the pressure is also (x/V), the volume is 1.50L, the number of moles is x, and the temperature is 227 degrees Celsius (convert to Kelvin).
For PCl5, the pressure is (0.100 - x/V), the volume is 1.50L, the number of moles is (0.100 - x), and the temperature is 227 degrees Celsius (convert to Kelvin).
Step 6: Use the ideal gas law to solve for x
Solve the ideal gas law equations for PCl3, Cl2, and PCl5.
Step 7: Calculate the partial pressures and total pressure
Substitute the value of x into the equilibrium expressions for PCl3, Cl2, and PCl5 to calculate their partial pressures. Then, add these partial pressures to find the total pressure.
By following these steps, you can determine the partial pressures of PCl3 and Cl2, as well as the total pressure of the gases in the flask at equilibrium.