Phosgene decomposes into carbon monoxide and chlorine gas according to the follo

Question

Phosgene decomposes into carbon monoxide and chlorine gas according to the following equation The Gibbs energy for this reaction is +82.1 kJ mol^-1 at standard temperature and pressure and the enthalpy of reaction is +107.6 kJ mol^-1. Calculate the equilibrium constant for this reaction at 773 K and determine the degree of dissociation of phosgene at 773 K at a total pressure of 1 bar. You may assume that the enthalpy of reaction is independent of temperature. Note any assumptions used in your calculations.

Explanation / Answer

At 298.15 K, deltaG= -RT lnK 298.15

lnK298.15 =-82.1*1000/(8.314*298.15)

K298.15= 4.129*10-15

we also know that

ln(K2/K298.15)= (deltaH/R)*(1/298.15-1/T2)

K2= Equilibrium constant at T2= 773K

Ln(K2/ K298.15) = (107.6*1000/8.314)*(1/298.15-1/773)=26.67

K2= 4.129*10-15*3.81*1011=0.001572

The reaction is CoCl2(g) ----àCo(g)+ Cl2(g)

Basis : 1 M of COCl2 is assumed to be initially present

Let x= drop in concentration to reach equilibrium

So at equilibrium [CoCl2] =1-x, [CO] =[Cl2]= x

K2= equilibrium constant = [CO] [Cl2]/ [COCl2] = x2/ (1-x)= 0.001572

When solved using excel, x= 0.0389

Degress of dissociation = 0.0389*100/1= 3.89%

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