# Why would a solution of FeCl2 appear brown/yellow sort of like FeCl3?

If we prepare a solution of $\ce{FeCl2}$ by dissolving solid $\ce{FeCl2}$ powder in water, supposing the bottle of $\ce{FeCl2}$ was left open for sufficient time, is it possible for $\ce{FeCl2}$ to have oxidised to $\ce{FeCl3}$? I can't find a suitable reaction for this excecpt one which includes $\ce{HCl}$ which was definitely not present when preparing the mixture.

• Yes, it is very common. You actually have to protect Fe(II) solution from air when you do experiments if oxidation is something you want to avoid.
– Greg
Jun 6 '17 at 4:03

The yellow/brown color you are seeing in the solution of $$\ce{FeCl2}$$ is due to formation of various types of hydrated iron oxide and not iron(III) chloride. They may include both the anhydrous($$\ce{FeOOH}$$) or the monohydrate($$\ce{FeOOH.H2O}$$) often referred to as ferric hydroxide($$\ce{Fe(OH)3}$$). They are generally yellow to brown in color. The reaction is given here:

$$\ce{4FeCl2 + 6H2O + O2 → 4FeO(OH) + 8HCl}$$

Reaction between iron chloride (II), water and oxygen(moisture) to form iron metahydroxide(Iron(III) oxide-hydroxide) and hydrogen chloride. The reaction takes place at reflux.

Moreover, if you further heat iron chloride in oxygen at 450-480℃, iron oxide will form.

\begin{align} \ce{Fe^2+ &-> Fe^3+ + e-} &\quad E^\circ &= \pu{+0.771 V} \tag{R1}\\ \ce{O2 + 2 H2O + 4 e- &-> 4OH-} &\quad E^\circ &= \pu{+0.40 V} \tag{R2} \end{align}
Thus, oxidation of iron(II) to iron(III) ions is feasible with atmospheric oxygen as the overall cell potential for the reaction is $$\pu{+1.171 V}$$ and it does occur at a reasonably observable rate.