You have a 1.0-L flask filled with phosphorous pentachloride (PCl_5) at a high e

Question

You have a 1.0-L flask filled with phosphorous pentachloride (PCl_5) at a high enough temperature such that it is entirely in gaseous form. You remember that PCl_5 can decompose into chlorine gas (Cl_2) and phosphorous trichloride gas (PCl_3) in a reversible reaction. You start with one atm of PCl_5 in your flask. After equilibrium is achieved at constant temperature, you notice that the new pressure inside the flask is fifty percent higher than the initial pressure. Determine the value of the equilibrium constant, K_p, for the reaction. Explain your reasoning. You are having lunch with a friend, who is also studying the same reaction in a 1. 0-L flask at the same temperature. She has combined one atm of each of the three gases in a flask, and is now wondering what will happen. Based on your own observations of this reaction, do you predict that the pressure in her flask will go up, go down, or remain the same? Explain your reasoning.

Explanation / Answer

(a)   PCl5 <==> PCl3 + Cl2 (g)

New pressure = 1atm + 1atm * 50% = 1.5 atm

At equilibrium : (1-x) + x+x = 1.5

                        or, x = 0.5 atm

Kp = P(PCL3) * P(CL2)/ PPCl5

       = x*x/1-x = 0.5 *0.5/0.5 = 0.5 atm

------------------------------------------------------------------------

(b)

reaction quotient (Q) = 1atm * 1atm/1atm = 1atm

So, Q > Kc

Pressure of the gas will be reduce to reach the equilibrium value

PCl5 PCl3 Cl2 initial 1 0 0 change -x x x equilibrium 1-x x x

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