Help Please!!! I don't know how to start this problem! An equilibrium mixture contains N2O4, (P= 0.28 atm ) and NO2 (P= 1.1 atm ) at?

N₂O₄(g) ⇌ 2NO₂(g)

At original equilibrium state :
P(N₂O₄) = 0.28 atm
P(NO₂) = 1.1 atm

Under constant temperature, P₁V₁ = P₂V₂
When the volume is doubled, the total pressure and all partial pressure is halved.

According to Le Chatelier's principle, when total pressure decreases, the equilibrium position will shift to the right in order to increase the number of gaseous molecules.

When new equilibrium is established :
P(N₂O₄) = [(0.28/2) - y] atm = (0.14 - y) atm
P(NO₂) = [(1.1/2) + 2y] atm = (0.55 + 2y) atm

Kp = (P(NO₂))² / P(N₂O₄)

Under constant temperature, the equilibrium constant keeps unchanged. The equilibrium constant, Kp :
1.1² / 0.28 = (0.55 + 2y)² / (0.14 - y)
1.21 / 0.28 = (0.3025 + 2.2y + 4y²) / (0.14 - y)
0.28 × (0.3025 + 2.2y + 4y²) = 1.21 × (0.14 - y)
0.0847 + 0.616y + 1.12y² = 0.1694 - 1.21y
1.12y² + 1.826y - 0.0847 = 0

Solve the equation, y = 0.045 or y = -1.68 (rejected)

When the system reaches a new equilibrium :
Equilibrium pressure of N₂O₄ = (0.14 - 0.045) atm = 0.095 atm
Equilibrium pressure of NO₂ = (0.55 + 2 × 0.045) = 0.64 atm