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Question

The reaction N2O4 (g) 2NO2 (g) is
carried out at 298 K and 20 bar. Five moles of each of N2O4and
NO2, are taken initially. Given: ΔfG
(N2O4) = 100 kJ mol1 and ΔfG
(NO2) = 50 kJ mol1. The reaction proceeds at an initial pressure
of 20 bar. Determine the direction in which the reaction proceeds to achieve
equilibrium. Also determine the amounts of N2O4 and NO2 when the reaction attains equilibrium.


A

= 8.333, = 1.667

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B

= 1.667, = 3.333

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C

= 3.333, = 1.667

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D

= 6.667, = 1.667

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Solution

The correct option is D

= 6.667, = 1.667


Some thermodynamic data is given.

ΔG = ΔG + RTlnQp- (1)

R = gas constant. T = temperature; Qp = reaction quotient

Qp = (PNO2)2(PN2O4) = 10210 =
10 bar

ΔG (reaction) = 2 ΔG(NO2) -
ΔG(N2O4) = 2 (50 kJ mol1) - 100 kJ
mol1 = 0

ΔG = ΔG + RTlnQp = 0 +
(8.314 J K1 mol1) (298 K) In (10) = (8.314 x 298 x 2.303)
Jmol1 = 5.7058 kJmol1

ΔG= - RT InKp = 0 (from calculation)

This implies that Kp = 1 bar; Earlier Qp =
10 bar.

As Qp>Kp the reaction will proceed in
the reverse direction. Thus we choose 2x moles to disappear from NO2 and x
moles of N2O4to appear on the other side

N2O4(g)2NO2(g)

5 mol + x 5 mol - 2x; total amount = 10 mol -x

where x is the amount of N2O4 formed at equilibrium.

PN2O4=5mol+x10molx (20 bar) and PNO2=5mol2x10molx (20 bar)

At equilibrium,

Kp = 1 = (PNO2)2 / (PN2O4) \{ here we
take the equilibrium partial pressures; not the given values\}

[5mol2x10molx(20)]2=5mol+x10molx(20)

20(5mol2x)2 = (5 mol + x) (10 mol - x)

500mol2 + 80 x2 - 400x mol = 50 mol2 + 5x mol - x2

81x2 - 405x mol + 450 mol2 = 0

x=(405 mol)±(405mol)24(81)(450 mol2)2×81

Solving x = 3.333 mol or x = 1.667 mol

Since (5 - 2x )>0

X = 1.667 mol

Number of moles N2O4 = 5 mol + x = 6.667 mol and Number of moles
of NO2= 5 mol - 2x = 1.667 mol


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