For the following reaction at 600. K, the equilibrium constant, Kp is 11.5.

PCl5(g) PCl3(g) + Cl2(g)

Suppose that 3.314 g of PCl5 is placed in an evacuated 460. mL bulb, which is then heated to 600. K.
(a) What would be the pressure of PCl5 be if it did not dissociate?
1 atm
(b) What is the partial pressure of PCl5 at equilibrium?
2Check the number of significant figures. atm
(c) What is the total pressure in the bulb at equilibrium?
3Check the number of significant figures. atm
(d) What is the degree of dissociation of PCl5 at equilibrium?
4%

Would you mind going through how you obtained the answers you have? I don't see how you could come up with 1 atm for (a). Is that the right answer?

To find the answers to these questions, we need to use the equilibrium constant expression and the given information.

(a) To find the pressure of PCl5 if it did not dissociate, we assume that all of the PCl5 remains as PCl5 and none of it dissociates. Since the volume of the bulb is given as 460 mL, we can convert it to liters: 460 mL = 0.460 L.

Using the ideal gas law equation PV = nRT, we can rearrange it to find the pressure (P):
P = (nRT) / V

R is the ideal gas constant (0.0821 L·atm/mol·K) and T is the temperature in Kelvin (600 K). n is the number of moles of PCl5. To find n, we need to convert the given mass of PCl5 to moles:

moles of PCl5 = mass of PCl5 / molar mass of PCl5

The molar mass of PCl5 is:
(1 atom of P × atomic mass of P) + (5 atoms of Cl × atomic mass of Cl)

Using the atomic masses from the periodic table (P = 31.0 g/mol, Cl = 35.5 g/mol):
molar mass of PCl5 = (1 × 31.0) + (5 × 35.5) = 208.5 g/mol

Now, we can calculate the moles of PCl5 using the given mass (3.314 g) and its molar mass:
moles of PCl5 = 3.314 g / 208.5 g/mol

Now that we have the moles of PCl5, we can substitute the values into the equation for pressure to find the pressure of PCl5:
P = (nRT) / V
P = (moles of PCl5 × R × T) / V

Substitute the values:
P = (moles of PCl5 × 0.0821 L·atm/mol·K × 600 K) / 0.460 L

Solve the equation to find the pressure of PCl5 if it did not dissociate.

(b) To find the partial pressure of PCl5 at equilibrium, we need to use the equilibrium constant expression for the reaction. The equilibrium constant expression for the given reaction is:

Kp = (PCl3 × Cl2) / PCl5

The equilibrium constant (Kp) is given as 11.5. Since we want to find the partial pressure of PCl5, we can rearrange the equation:

PCl5 = (PCl3 × Cl2) / Kp

Substitute the known values into the equation to find the partial pressure of PCl5 at equilibrium.

(c) To find the total pressure in the bulb at equilibrium, we need to consider that PCl5 dissociates into two products, PCl3 and Cl2. The total pressure in the bulb is the sum of the partial pressures of all the gases at equilibrium.

Using the given equilibrium constant, we can calculate the partial pressures of PCl3 and Cl2 at equilibrium using the equilibrium constant expression and the partial pressure of PCl5 at equilibrium from part (b). Then, we can add these partial pressures together to find the total pressure.

(d) To find the degree of dissociation of PCl5 at equilibrium, we need to compare the moles of PCl5 at equilibrium with the initial moles of PCl5. The degree of dissociation (α) is given by the formula:

α = (moles dissociated at equilibrium) / (initial moles of PCl5)

We have already found the moles of PCl5 at equilibrium in part (b). To find the initial moles of PCl5, we use the same formula as before:

moles of PCl5 = mass of PCl5 / molar mass of PCl5

Substitute the given mass of PCl5 (3.314 g) into the formula and solve for the initial moles of PCl5.

Then, substitute the values into the formula for the degree of dissociation and calculate the result.

Note: Make sure to check the number of significant figures in each answer to match the given data.