# Stability of ferrous ions and iron(II) hydroxide in acidic conditions

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 $K_\mathrm{instability} \approx 10^{-37}$.

Half-reactions

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):

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).

Problem

• 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.
• States of aggregation should not be subscripted, it is not wrong, but the recommendations (Sec. 2.1.) are different. – Martin - マーチン Jun 16 '16 at 9:41
• @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). – Linear Christmas Jun 16 '16 at 14:51