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) enter image description here

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?



1 Answer 1


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.

  • $\begingroup$ 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) $\endgroup$
    – Charly
    Commented Dec 7, 2017 at 21:53
  • $\begingroup$ @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. $\endgroup$
    – MaxW
    Commented Dec 7, 2017 at 21:57
  • 1
    $\begingroup$ @Charly: sorry but your lab question was wrong. Wouldn't be the first time or the last. $\endgroup$
    – Gert
    Commented Dec 7, 2017 at 22:01
  • $\begingroup$ 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. $\endgroup$ Commented Dec 8, 2017 at 4:12

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