Question

The reaction below is at equilibrium at a temperature T. There are four possible changes listed. Select all the changes that will shift the equilibrium so as to produce more products
MgO(s) + SO3(g) + 95 kJ <======> MgSO4(s)
I – Increase the temperature
II – Increase the volume
III – Add more MgO
IV – Remove SO3

(A) I and III
(B) I
(C) I, II and III
(D) I, II, III and IV

Answer 1

Answer:

The correct answer is B.

Explanation:

Heterogeneous equilibrium is that in which reagents and products are present in more than one phase.

When the reaction is carried out in a closed container, three equilibrium phases are present: solid magnesium oxide, solid magnesium sulfate and gaseous sulfur trioxide.

Hence, the equilibrium contant is given by:

$K=\frac{[MgSO_4]}{[MgO][SO_3]} =\frac{1}{[SO_3]}$

The concentrations in the equilibrium equation are the relationships of the real concentrations between the concentrations in the standard state. Since the standard state of a pure solid is the pure solid itself, the ratio of concentrations for a pure solid is equal to one.

Now, we analyse each statement:

I) As the reaction is endothermic (ΔH>0), increasing the temperature shifts the balance to the right because excess heat will be used to form more products.

II) Increasing the volume will decrease the concentration of SO₃, so Q>K and then this shifts the balance to the left.

III) As it is a heterogeneous balance, adding MgO will not affect the balance.

IV) Removing SO3 will decrease its concentration and therefore the reaction equilibrium will shift to the left.