1.30 mol each of CO} and Cl2 are introduced into an evacuated 1.80 L flask, and the following equilibrium is established at 668 K.

CO + Cl2 +++> COCl2 kp = 22.5

Calculate partial pressure of COCl2. i get about 30.9. but its wrong. i solved for x, COCL2/CO*Cl2, using ICE formula, converted everything to molarity, then to moles, then using ideal gas law. still cant get it.

If I work this I don't get 30.9 but 38.28 atm. I wonder why you converted to molarity since Kp is given and they want the answer in partial pressure COCl2?

I used PV = nRT. Plug in the numbers and obtain PCO and PCl2. Something like 39.6 atm or so. Then set up the ICE chart using atm and it becomes
x/(39.6-x)^2 = 22.5 and solve for x.

To solve for the partial pressure of COCl2, you need to use the equilibrium constant expression and the given equilibrium constant value (kp). Here's a step-by-step guide to solve this problem correctly:

Step 1: Write the balanced chemical equation:
CO + Cl2 ↔ COCl2

Step 2: Set up the ICE (Initial-Change-Equilibrium) table:
CO + Cl2 ↔ COCl2
Initial: 1.30 mol 1.30 mol 0 mol
Change: -x -x +x
Equilibrium: 1.30 - x 1.30 - x x

Step 3: Write the equilibrium constant expression:
kp = [COCl2] / ([CO] * [Cl2])

Step 4: Substitute the equilibrium concentrations into the equilibrium constant expression:
22.5 = x / ((1.30 - x) * (1.30 - x))

Step 5: Solve for x:
22.5 = x / ((1.69 - x)^2)
22.5(1.69 - x)^2 = x
x = 0.07085 mol

Step 6: Calculate the partial pressure of COCl2 using the ideal gas law:
PV = nRT

Since the volume is given (1.80 L) and the temperature is provided (668 K), we can rearrange the equation to solve for pressure (P):
P = (nRT) / V

P = (0.07085 mol * 0.0821 L·atm/(mol·K) * 668 K) / 1.80 L
P ≈ 18.65 atm

Therefore, the partial pressure of COCl2 is approximately 18.65 atm, which is the correct answer.

To calculate the partial pressure of COCl2, you can follow these steps:

1. Start by writing down the balanced chemical equation for the reaction: CO + Cl2 → COCl2

2. Let's assume that initially, before the reaction takes place, there are n COCl2, n CO, and n Cl2 moles in the flask.

3. According to the stoichiometry of the reaction, for every 1 mol of COCl2 formed, 1 mol of CO and 1 mol of Cl2 react. Therefore, at equilibrium, the number of moles of CO and Cl2 that have reacted will be equal to the number of moles of COCl2 formed.

4. Use the ideal gas law equation PV = nRT to relate the number of moles (n) to the partial pressure (P), volume (V), and temperature (T) of the system. Rearrange the equation to solve for n: n = PV / RT.

5. Now, let's focus on COCl2. At equilibrium, the number of moles of COCl2 formed will be equal to n COCl2. Therefore, the partial pressure of COCl2 can be expressed as P COCl2 = (n COCl2 RT) / V.

6. To solve for n COCl2, you can use the value of Kp, which is the equilibrium constant for the reaction. The equilibrium expression for the reaction is: Kp = P COCl2 / (P CO * P Cl2).

7. Rearrange the equation to solve for n COCl2: n COCl2 = Kp * P CO * P Cl2.

8. Since you have the initial number of moles for CO and Cl2 (both are 1.3 mol) and the equilibrium constant (Kp = 22.5), substitute these values into the equation n COCl2 = Kp * P CO * P Cl2.

9. Convert the initial number of moles of CO and Cl2 to their respective initial partial pressures. Use the ideal gas law equation to determine their partial pressures: P CO = (n CO RT) / V and P Cl2 = (n Cl2 RT) / V.

10. Substitute the calculated values of P CO, P Cl2, and Kp into the equation from step 8 to find n COCl2.

11. Finally, substitute the value of n COCl2 and the given volume into the expression for P COCl2 from step 5 to find the partial pressure of COCl2.

By following these steps, you should be able to calculate the correct partial pressure of COCl2.