Question

The equilibrium constant (KP) is 0.16 at a particular temperature for the reaction: N₂O₄(g) ⇌ 2NO₂(g)

Given the following sets of initial conditions, what is the net change that must occur for the reaction to reach equilibrium? Does the reaction shift left to reach equilibrium, does the reaction shift right to reach equilibrium or is the reaction at equilibrium at these initial concentrations so no net change will occur?

equilibrium, left, right? P_NO2 = 0.154 atm, P_N2O4 = 0.148 atm

equilibrium, left, right? P_NO2 = 0.210 atm, P_N2O4 = 0.138 atm

equilibrium, left, right ? P_NO2 = 0.149 atm, P_N2O4 = 0.05 atm

equilibrium, left, right? P_NO2 = 0.087 atm, P_N2O4 = 0.13 atm

equilibrium, left, right? P_NO2 = 0.068 atm, P_N2O4 = 0.058 atm

Answer 1

Kp = [NO2]²/[N2O4]
0.15[N2O4] = [NO2]²

From PV=nRT P (Pressure) is directly proportional to n (moles). As the volume is constant we can use the pressure as equivalent to moles.

So for the first one substitute the pNO2 into the equilibrium equation and solve for pN2O4
[N2O4] = 0.154²/0.16
            = 0.148
As you have 0.148 it will be in equilibrium

2. [N2O4] = 0.21²/0.16
                  = 0.275

As you have 0.138 which is less concentration the reaction will move to the N₂O₄ or to the left

3. [N2O4] = 0.149²/0.16
                  = 0.138

As you have 0.138 which is less concentration the reaction will move to the N₂O₄ or to the left

4. [N2O4] = 0.087²/0.16
                  = 0.047

As you have 0.047 which is too much the reaction will move back to the NO2 or to the right.

5. [N2O4] = 0.068²/0.16
                  = 0.0289

As you have 0.0289 which is too much the reaction will move back to the NO2 or to the right.