What is the pH of a $\ce{0.1 M}$ solution of $\ce{Fe(CH_3COO)_3}$? I'm not sure how to do this, as it would involve complicated mass balance equations, but I think I can apply the approximation $\ce{pH = \frac{1}{2}*(pKa(Fe^{3+}) + pKa(CH_3COOH)) = \frac{1}{2}*(4.76 + 2.22) = 3.48}$.

Is this approximation valid in this situation? If not, how would I correctly calculate the pH of the solution without solving complex mass balance equations?

  • $\begingroup$ Well the acetate will protonate making the pH go up. The iron can form hydroxides making the pH go down. Assume salt completely ionizes, and first assume that iron can be ignored. If iron is now above its solubility limit then you have to solve iron and acetate simultaneously. $\endgroup$
    – MaxW
    Jul 9 '19 at 21:38
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    $\begingroup$ Does your approximation completely ignore the concentration of the solution? $\endgroup$ Jul 9 '19 at 21:39
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    $\begingroup$ Compare ammonium acetate, chemistry.stackexchange.com/questions/75047/… $\endgroup$ Jul 9 '19 at 21:52
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    $\begingroup$ Well, not iron but its aquo complex to be precise... $\endgroup$
    – Mithoron
    Jul 9 '19 at 22:21
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    $\begingroup$ Like $\ce{ [Fe(H2O)6]^3+ + H2O <=> [Fe(H2O)5(OH)]^2+ + H3O+}$ $\endgroup$
    – Poutnik
    Jul 9 '19 at 23:47

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