Home >
Community >
Why would a solution of FeCl2 appear brown/yellow sort of like FeCl3?
Upvote
23
Downvote
+ Solutions
+ Chemistry
+ Redox
Posted by
Pamela Smith
Why would a solution of FeCl2 appear brown/yellow sort of like FeCl3?
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.
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 formiron metahydroxide(Iron(III) oxide-hydroxide) and hydrogen chloride. The reaction takes place atreflux.
But no solid is formed during the oxidation process... Wouldnt the hydrated iron (III) oxide-hydroxide give an insoluble precipitate. From my own observations in the laboratory, I have never seen any precipitate in a solution of yellow iron (III) chloride solution. How would you explain the absence of a ppt?More
@TanYongBoon iron(III) hydroxide is indeed insoluble in neutral and basic pH and precipitates in the bottom of the solution(according to 1st link). However, the ferrous chloride solution present in the laboratory has somewhat an acidic pH for which the the hydrated iron oxide or iron(III) hydroxide solubilize in the solution and thus not get any precipitate.More
Thats interesting... I guess what Ive learnt in school is only for the Cambridge Advanced Levels, which is probably the simplified explanation. Your explanation seems credible and convincing as well.More
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.
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.
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:
$$\ce{4FeCl2 + 6H2O + O2 → 4FeO(OH) + 8HCl}$$
Moreover, if you further heat iron chloride in oxygen at 450-480℃, iron oxide will form.
References(all pdf links)
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:
$$\ce{4FeCl2 + 6H2O + O2 → 4FeO(OH) + 8HCl}$$
Moreover, if you further heat iron chloride in oxygen at 450-480℃, iron oxide will form.
References(all pdf links)
More
VOTE
VOTE
VOTE
VOTE
$$ \begin{align} \ce{Fe^2+ &-> Fe^3+ + e-} &\quad E^\circ &= \pu{+0.771 V} \tag{R1}\\ \ce{O2 + 2 H2O + 4 e- &-> 4OH-} &\quad E^\circ &= \pu{+0.40 V} \tag{R2} \end{align} $$
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.
$$\begin{align}\ce{Fe^2+ &-> Fe^3+ + e-} &\quad E^\circ &= \pu{+0.771 V} \tag{R1}\\\ce{O2 + 2 H2O + 4 e- &-> 4OH-} &\quad E^\circ &= \pu{+0.40 V} \tag{R2}\end{align}$$
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.
More
VOTE