Question

At the start of a reaction, there are 0.0249 mol N2,

3.21 x 10-2 mol H2, and 6.42 x 10-4 mol NH3 in a
3.50 L reaction vessel at 375°C. If the equilibrium constant, K, for the reaction:
N2(g) + 3H2(g)= 2NH3(g)
is 1.2 at this temperature, decide whether the system is at equilibrium or not. If it is not, predict in which direction, the net reaction will proceed.​

Answer 1

Step-by-step explanation:

The reaction is given as:

`N_(2(g)) + 3H_(2(g)) \to 2NH_(3(g))`

The reaction quotient is:

`Q_C = ([NH_3]^2)/([N_2][H_2]^3)`

From the given information:

TO find each entity in the reaction quotient, we have:

`[NH_3] = (6.42 * 10^(-4))/(3.5)\\ \\ NH_3 = 1.834 * 10^(-4)`

`[N_2] = (0.024 )/(3.5)`

`[N_2] = 0.006857`

`[H_2] =(3.21 * 10^(-2))/(3.5)`

`[H_2] = 9.17 * 10^(-3)`

∴

`Q_c= ((1.834 * 10^(-4))^2)/((0.0711)* (9.17* 10^(-3))^3) \\ \\ Q_c = 0.6135`

However; given that:

`K_c = 1.2`

By relating
`Q_c \ \ and \ \ K_c`, we will realize that
`Q_c \ \ < \ \ K_c`

The reaction is said that it is not at equilibrium and for it to be at equilibrium, then the reaction needs to proceed in the forward direction.