What happens to the colour change of an equilibrium reaction when one of the reactants begins to form a complex? (specific problem included) I'm working on an equilibrium lab/quiz for my chemistry class and I came across this reaction:

$\ce{Fe^{+3}\ \text{(pale yellow)} + SCN- &lt;=&gt; FeSCN^{+2}}\ \mathrm{(red)}$

Then the lab said a stressor was added: $\ce{Na2HPO4}$ was added to the equilibrium reaction -- which formed a complex with some of the $\ce{Fe^{3+}}$ ions.

This stressor changed the solution's color from an orange to a pale-yellow as seen in the image below:

(left: before stressor ($\ce{Na2HPO4}$) was added; Right: after stressor ($\ce{Na2HPO4}$) was added)


I was asked to explain why that after $\ce{Na2HPO4}$ is added (which actively forms a complex with the iron ion ($\ce{Fe^{3+}}$)) the system/solution turns increasingly pale-yellow. 

Basically I want to know what is happening chemically when $\ce{Na2HPO4}$ is added, and more specifically why their is the resultant color change from orange to pale-yellow.

My basic confusion: wouldn't the formation of a complex between $\ce{HPO4^{2-}}$ and $\ce{Na2HPO4}$ actively remove the $\ce{Fe^{3+}}$ ions responsible for the pale-yellow color of the solution? Making the solution less rather than more yellow?

Thanks!
Going by the right hand picture, what was formed there is the very poorly soluble and lightly yellow $\ce{FePO4.2H2O}$. A precipitate and clear supernatant liquid can clearly be seen.

If so, the reaction is:

$$\ce{FeSCN^{2+}(aq) + HPO4^{-}(aq) + 3 H2O(l) \to FePO4.2H2O(s) + SCN-(aq) + H3O+(aq)}$$

This explains why on adding the $\ce{Na2HPO4}$ the red colour of the $\ce{FeSCN^{2+}}$ complex disappears and a yellowish, flocculant precipitate appears. This is not however, as the OP suspected, a complex. It's merely an almost insoluble salt.

Some info on $\ce{FePO4}$.



A true complex would have been formed by the addition of $\ce{CN-}$, which would have led to the formation of $\ce{Fe(CN)6^{3-}}$, which has a much higher formation constant ($K_f$) than the thiocyanate complex. It's also weakly coloured and would not form a precipiate. The reaction would have been:

$$\ce{FeSCN^{2+}(aq) + 6 CN-(aq) \to Fe(CN)6^{3-}(aq) + SCN-(aq)}$$

The solution would have evolved from bordeaux to weakly amber/red.