I was reading a chemisty textbook when I came across the following statement.

$\ce{Fe(OH)2}$ is not stable, even more so in basic conditions. [---] $\ce{Fe^2+}$ ions are also less likely to oxidise in the low pH range.

First attempts

No explanation was given. At first, I was hoping for a simple case of Le Châtelier – Braun principle. However, the equilibrium

$$\ce{4[Fe(OH2)6]^2+ (aq) + 4H3O^+ (aq) + O2 (g) <=> 4[Fe(OH2)6]^3+ (aq) + 6H2O (l)}$$

would suggest the opposite. Next, I considered a simple approach to complex stability. As a general guideline, metals in lower oxidation states favour cationic complexes. The opposite is also true. Therefore, one might conclude then that if anionic complexes are not preferred, $\ce{Fe^2+}$ ions ought to be stable.

One example would be that the hydrolysis is more hampered in acidic media. $$\ce{[Fe(OH2)6]^2+ (aq) + H2O (l) <=> [Fe(OH)(OH2)6]^+ (aq) + H3O^+ (aq)}$$ Furthermore, in very basic conditions the hexahydroxidoferrate(II) would form. $$\ce{[Fe(OH)2(OH2)4] (s) + 4OH^{-} (aq) <=> [Fe(OH)6]^4- (aq) + 4H2O (l)}$$

However, this rule breaks down for $\ce{[Fe(CN)6]^4-}$ which is rather stable with $K_\mathrm{instability} \approx 10^{-37}$.


As you might imagine, I was trying to avoid using half-reactions and the Nerst equation. Iron chemistry is rich and there are a myriad of possibilities. Before we move forward, let us note that the corresponding ion electronic structures are $$\ce{Fe^2+:[Ar]{3d^6}};\\ \ce{Fe^3+:[Ar]{3d^5}}.$$ So it seems that ferric ions are indeed more stable. Yet the reduction half-reaction $$\begin{align} \ce{Fe^3+ + e- &-> Fe^2+} & E^\circ &= +0.77~\mathrm{V} \end{align}$$ implies the opposite. (Is this due to solvation?)

I did find a useful Pourbaix diagram (Andel Früh: Pourbaix diagram of Iron; $c(\ce{Fe}) = 10^{-6}~\mathrm{mol/l}$, $T = 25~^\circ\mathrm{C}$; wikimedia.org):

Pourpaix diagram for iron (from the University of Bath and Western Oregon University, by Andel Früh

What we can tell is that ferric ions themselves are not favoured at a higher pH. In fact, the opposite is true. It just requires a strong oxidizer. This does not mean ferrous ions are not oxidized in basic conditions, it is just that the stable species are iron(III) oxide hydrates (probably iron(III)oxidehydroxide hydrates as well).


I have not really thought about this particular issue before, so I would definitely love to hear your ideas.

  • Why are ferrous ions and iron(II) hydroxide more stable at lower pH?
  • Why is the ferrous ion generally more stable, even though the electronic structure suggests otherwise?
  • This Pourbaix diagram does not take into account possible complex formation. If someone could explain the issue via complexes as well, I would greatly appreciate it.
  • 2
    $\begingroup$ States of aggregation should not be subscripted, it is not wrong, but the recommendations (Sec. 2.1.) are different. $\endgroup$ Commented Jun 16, 2016 at 9:41
  • $\begingroup$ @Martin-マーチン: Hmm, that's interesting. Never thought it should be written in any other way. Not to argue with IUPAC, but to me it seems not using subscript unnecessarily gives the states too much attention while completely omitting them would have the opposite effect. Subscript is kind of the middle ground :). Nevertheless, thank you for editing, I'll keep it in mind in the future (if possible). $\endgroup$ Commented Jun 16, 2016 at 14:51

1 Answer 1


The stability of iron(II) hydroxide J. Chem. Educ., 1957, vol. 34 pages 178-179 shows that the book isn't entirely correct.

...Iron(II) hydroxide would thus appear to be unstable in neutral solutions and to possess increasing stability as the basicity of the medium increases.

The foregoing reasoning led to the attempt to prepare iron(II) hydroxide from the chloride under such conditions that the hydroxyl ion concentration was not allowed to fall below approximately one molal until the chloride ion had been removed. The hydroxide was precipitated from strongly basic solution by the addition of dilute iron(II) chloride to a saturated solution of sodium hydroxide in an air-free system with vigorous stirring. The precipitate was repeatedly washed with molal sodium hydroxide (oxygen and carbonate-free) until a test for chloride ion with silver nitrate could no longer he obtained. Five additional one-liter washings with normal base completed the preparation of the iron(II) hydroxide used for the solubility studies. This technique led to the formation of a coarse-grained precipitate, white when viewed by reflected light and slightly greenish tinged by transmitted light. Dilutions of the material were made to one-tenth molal base without evidence of darkening when maintained under an oxygen-free atmosphere.

  • $\begingroup$ Sorry for a delayed response. So, basically, the book wrongly generalised the stability trend of ferrous ions to iron(II) hydrokside? Also, what about the inherent stability of the ions in solution? Would the situation be reversed in vacuum, as predicted by the electronic structure? Thanks for your reply! Sadly I am unable to retrieve the complete study you cited. The quote helped. $\endgroup$ Commented Jun 2, 2016 at 18:18

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