Question

When silver nitrate is added to the Fe/SCN equilibrium, why is the colorless intense and a precipitate forms?

Answer 1

Answer:

Here's what I get  

Explanation:

You have an equilibrium reaction between Fe³⁺/ SCN⁻ and FeSCN²⁺.

$\underbrace{\hbox{Fe ^{3+} }}_{\text{pale yellow-green}} +\underbrace{\hbox{SCN ^{-} }}_{\text{colourless}} \, \rightleftharpoons \, \underbrace{\hbox{Fe(SCN) ^{2+} }}_{\text{deep blood red}}$

When you add AgNO₃, the Ag⁺ reacts with the SCN⁻. It forms a colourless precipitate of Ag(SCN).

Ag⁺(aq) + SCN⁻(aq) ⟶ AcSCN(s)

According to Le Châtelier's Principle, when we apply a stress to a system at equilibrium, the system will respond in a way that tends to relieve the stress.

If you add Ag⁺ to the equilibrium solution, it removes the SCN⁻ [as an Ag(SCN) precipitate].

The system responds by trying to replace the missing SCN⁻:

The Fe(SCN)²⁺ dissociates to form SCN⁻, so the position of equilibrium shifts to the left,

You now have more Fe³⁺ and SCN⁻ and less of the highly coloured Fe(SCN)²⁺ at the new equilibrium.

The deep red colour becomes less intense.