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Phosphorus pentachloride is introduced into an empty gas syringe which has a movable, tightly-f‌itting plunger. The gas is allowed to expand until equilibrium is reached at a temperature at which the phosphorus pentachloride partially dissociates.

$$\ce{PCl5(g) <=> PCl3(g) + Cl2(g)}$$

expanding gas plunger

Which statements are correct?

  1. The equilibrium pressure inside the syringe will be greater than atmospheric pressure.
  2. When the plunger is pushed in the equilibrium adjusts to produce more $\ce{PCl5(g)}.$
  3. The volume of gas in the syringe at equilibrium will be greater than if no dissociation had occurred.

The correct answer is 2 and 3. Option 2 can be explained using Le Chatelier's principle and I think option 1 is wrong because the pressure inside should be equal to the pressure outside when the reaction reaches equilibrium.

But why is option 3 correct? Assuming the plunger is entirely extended and held stationary until the reaction reaches equilibrium, I think the volume of gas will be smaller if no dissociation had occurred, because there will be less molecules, hence lower pressure to push against the plunger as soon as it is released. My explanation might be entirely wrong since I don't really get how the experiment is being conducted here.

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The equation $3$ is correct because in the reaction, one gaseous molecule gets dissociated and produces a total of $2$ gaseous molecules. The number of gaseous molecules increases if the dissociation reaction proceeds. At a given pressure, the volume occupied by $2$ gaseous molecules is twice the volume occupied by $1$ gaseous molecule

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