# 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:

$\ce{Fe^{+3}\ \text{(pale yellow)} + SCN- <=> FeSCN^{+2}}\ \mathrm{(red)}$

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.

• Gert this might very well be the right answer (and thank you a ton for the quick reply), but quick clarification: I didn't suspect that the reaction between Na2HPO4 with Fe+3 forms a complex, my lab question informed me so "the HPO4-2 ion forms a complex with the Fe+3 ion." So maybe both reactions happened?.... (the insoluble salt you spoke of aswell as the ironIII complex the lab spoke of) – Charly Dec 7 '17 at 21:53
• @Charly = You can clearly see that something, the precipitate, is scattering light in the "yellow" test tube. A complex would have a "clear" (not colorless) solution as the red solution does. I'll point out that the lab almost certainly had a centrifuge. Centrifuging the yellow test tube would have separated the ppt nicely from the solution. – MaxW Dec 7 '17 at 21:57
• @Charly: sorry but your lab question was wrong. Wouldn't be the first time or the last. – Gert Dec 7 '17 at 22:01
• Please note that the iron(III) thiocyanate complex is better described as the following: $\ce{[Fe(SCN)_{$x$}(H2O)_{$6-x$}]^{3-x}}; x \in \{1,2,3\}$. It is also with itself in equilibrium. || Most of the time there is no need for bracketing subscripts within mhchem constructs, i.e. \ce{Fe^3+} is actually preferred over \ce{Fe^{3+}}. There is also a wide variety of arrows available, that better typeset than \to, e.g. ->, <=>, <-->, etc. Just for further reference. – Martin - マーチン Dec 8 '17 at 4:12