Nitrosyl bromide (NOBr) gas decomposes into nitrogen(II) oxide (NO) gas and bromine gas (Br₂) according to the following reaction at a temperature T = 25.0 °C: 2 \operatorname{NOBr}(\mathrm{g})<->2 \mathrm{NO}(\mathrm{g})+\mathrm{Br}_{2}(\mathrm{~g}) where the double arrow <-> denotes both forward and backward reactions. You may assume that the gases behave ideally at this temperature and for all pressures below. The molar masses are: NOBr (109.9 g/mol), NO (30.0 g/mol), and Br₂ (159.8 g/mol). You introduce NOBr into an evacuated vessel, seal it, and wait. After equilibration, you find that the total pressure of all of the gaseous species is 0.0600 bar and that the mass density arising from all gases in the container is 0.200 g/L. What are the equilibrium pressures of Br₂(g) and NO(g)? What is the dimensionless equilibrium constant K for this reaction at 25.0 °C? (c) Suppose that you suddenly inject 0.0200 bar of Br₂ gas to the vessel that was at equilibrium in part (a), and the temperature remains at 25°C. What is the dimensionless reaction quotient Q immediately after the injection? ) After the injection described in part (b), will the pressure of NO gas increase or decrease? Briefly justify your answer using the numerical values of Q and K.Calculate the new pressure of NO(g) after equilibration.