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Ramium
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[*] posted on 7-2-2015 at 23:59
Is this still iron carbonate???


I made some iron carbonate from iron sulphate and sodium carbonate. It was a greeny whitey gray colour. I left it in the sun to dry and it turned dark red.
Is it still iron carbonate dispiite the colour change???
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[*] posted on 8-2-2015 at 05:09


Is it crystaline?



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[*] posted on 8-2-2015 at 06:55


Quote: Originally posted by Ramium  
I made some iron carbonate from iron sulphate and sodium carbonate. It was a greeny whitey gray colour. I left it in the sun to dry and it turned dark red.
Is it still iron carbonate dispiite the colour change???


No, the FeCO3 has oxidised to Fe2O3.

Fe(II) compounds are very prone to oxidation to Fe(III), especially in neutral/alkaline conditions.

4 FeCO<sub>3</sub>(s) + O<sub>2</sub>(g) + 6 H<sub>2</sub>O(l) === > 4 Fe(OH)<sub>3</sub>(s) + 4 CO<sub>2</sub>(g)

Drying in the sun then dehydrates the Fe(OH)<sub>3</sub> to iron(III) oxide.

What you observed is entirely within expectations.


[Edited on 8-2-2015 by blogfast25]




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Ramium
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[*] posted on 8-2-2015 at 10:37


Dang!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!

Is there any way i could preserve the iron carbonate if i made more???
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[*] posted on 8-2-2015 at 11:02


Keep it where oxygen can't get it.

Also, depending on conditions, it's possible you have iron (ii) hydroxide.

Quote (wikipedia):
It is also easily formed as an undesirable by-product of other reactions, a.o., in the synthesis of siderite, an iron carbonate (FeCO3), if the crystal growth conditions are poorly controlled (reagent concentrations, addition rate, addition order, pH, pCO2, T, aging time, ...).




As below, so above.
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[*] posted on 8-2-2015 at 16:26


Quote: Originally posted by Cheddite Cheese  
Also, depending on conditions, it's possible you have iron (ii) hydroxide.



If you're referring precipitating Fe(OH)2 from an Fe(II) salt solution with a soluble carbonate, in that case CO2 must be evolved, for instance:

Fe<sup>2+</sup> + Na<sub>2</sub>CO<sub>3</sub> + 2 H<sub>2</sub>O === > Fe(OH)2 + CO<sub>2</sub> + 2 Na<sup>+</sup>




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Savior
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[*] posted on 8-2-2015 at 16:35


To preserve it, keep it away from water and air. With water it forms Fe3O4, with air and water together it forms Fe2O3.
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[*] posted on 8-2-2015 at 18:49


Quote: Originally posted by Savior  
To preserve it, keep it away from water and air. With water it forms Fe3O4, with air and water together it forms Fe2O3.


Fe3O4 (Magnetite) is an Fe(II,III) oxide: FeO.Fe2O3.

If the OP started from pure Fe(II) then in the absence of an oxidiser half of the Fe(II) cannot oxidise to Fe(III) and no Magnetite could form. Water alone cannot be responsible for the oxidation to Fe(II,III).

It's also not possible to "keep it away from water": it was formed in water!




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[*] posted on 8-2-2015 at 21:16


So if i put it in an airtight bottle it should not decompose???

Also is the iron oxide i made by accident useful or should i throw it out???
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[*] posted on 8-2-2015 at 22:24


Quote: Originally posted by blogfast25  
Quote: Originally posted by Savior  
To preserve it, keep it away from water and air. With water it forms Fe3O4, with air and water together it forms Fe2O3.


Fe3O4 (Magnetite) is an Fe(II,III) oxide: FeO.Fe2O3.

If the OP started from pure Fe(II) then in the absence of an oxidiser half of the Fe(II) cannot oxidise to Fe(III) and no Magnetite could form. Water alone cannot be responsible for the oxidation to Fe(II,III).

It's also not possible to "keep it away from water": it was formed in water!


H+ (water) is an oxidant. It somehow can oxidize it to Fe3O4. I remember when I electrolyzed water (with NaCl in it) using Fe anode, it generated insoluble Fe(II) Hydroxide (green color). When left on air it quickly becomes yellow and red, converting to Fe(III). This change becomes visible very quickly.

But, when I left that green powder in water in a closed bottle, it became black, and stays black forever.

When, bottle is opened, black powder becomes red/yellow after longer time.

Even Wikipedia says it on some Fe article, but I'm too lazy to find it, sorry.

So, he can't make Fe(II) carbonate, and I don't see any practical reason for making it. There is Fe even in soil, and there are more stable carbonates (Ca...). The only way to make pure carbonate of Al or Fe, is to do that without water, using high temperature. That's how Al carbonate is made.

My advice is that he should throw it out. It's not useful at all.

[Edited on 9-2-2015 by Savior]
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[*] posted on 8-2-2015 at 22:43


Quote: Originally posted by Savior  
Quote: Originally posted by blogfast25  
Quote: Originally posted by Savior  
To preserve it, keep it away from water and air. With water it forms Fe3O4, with air and water together it forms Fe2O3.


Fe3O4 (Magnetite) is an Fe(II,III) oxide: FeO.Fe2O3.

If the OP started from pure Fe(II) then in the absence of an oxidiser half of the Fe(II) cannot oxidise to Fe(III) and no Magnetite could form. Water alone cannot be responsible for the oxidation to Fe(II,III).

It's also not possible to "keep it away from water": it was formed in water!


H+ (water) is an oxidant. It somehow can oxidize it to Fe3O4. I remember when I electrolyzed water (with NaCl in it) using Fe anode, it generated insoluble Fe(II) Hydroxide (green color). When left on air it quickly becomes yellow and red, converting to Fe(III). This change becomes visible very quickly.

But, when I left that green powder in water in a closed bottle, it became black, and stays black forever.

When, bottle is opened, black powder becomes red/yellow after longer time.

Even Wikipedia says it on some Fe article, but I'm too lazy to find it, sorry.

So, he can't make Fe(II) carbonate, and I don't see any practical reason for making it. There is Fe even in soil, and there are more stable carbonates (Ca...). The only way to make pure carbonate of Al or Fe, is to do that without water, using high temperature. That's how Al carbonate is made.

My advice is that he should throw it out. It's not useful at all.

[Edited on 9-2-2015 by Savior]
so iron carbonate is completely unstable and cant be stored???
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[*] posted on 9-2-2015 at 01:29


Correct! Except under special conditions and production methods.

Any Iron(II) compund is unstable, even oxide (it converts to Fe3O4 if left alone).

Do you know which colour are Iron(ll) carbonate, chloride, and most salts?
Not green, as most think! But white! That green color appears when Iron(III) starts forming. Because it is yellow/red and this one is white, when you combine that colors, it is green.

That's only transition color...temporary. Because Iron(II) is never stable.
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[*] posted on 9-2-2015 at 02:37


Quote: Originally posted by Savior  


Any Iron(II) compund is unstable, even oxide (it converts to Fe3O4 if left alone).


That's only transition color...temporary. Because Iron(II) is never stable.

And Iron 2 (ferrous) Sulphate? - seems pretty stable to me, I have had some now for about a year, still lime green.

Actually not wishing to hijack this thread, but for the sake of not starting a new thread this would be a good place to ask a couple of questions. Comparing ferrous chloride/carbonate to ferrous sulphate what forces are at work whereby Oxygen is able to oxidize the ferrous to ferric so easily, yet in sulphate not so? The atomic orbitals and the structures are not giving anything away as far as I can see with my limited understanding. I've looked at everything I know, so what should I be looking at?


[Edited on 9-2-2015 by CHRIS25]




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[*] posted on 9-2-2015 at 09:35


Quote: Originally posted by Savior  
H+ (water) is an oxidant. It somehow can oxidize it to Fe3O4. I remember when I electrolyzed water (with NaCl in it) using Fe anode, it generated insoluble Fe(II) Hydroxide (green color). When left on air it quickly becomes yellow and red, converting to Fe(III). This change becomes visible very quickly.

But, when I left that green powder in water in a closed bottle, it became black, and stays black forever.

When, bottle is opened, black powder becomes red/yellow after longer time.

Even Wikipedia says it on some Fe article, but I'm too lazy to find it, sorry.

My advice is that he should throw it out. It's not useful at all.


I believe you are wrong on this. H3O+ (not H+, which in water cannot exist) is not capable of oxidising Fe(II) to Fe(III). If it was then dissolving iron in any non-oxidising acid would yield a Fe(III) salt solution and not a Fe(II) salt solution as in reality it does. I think the oxidation half potentials for Fe(0) to Fe(II) and Fe(II) to Fe(III) bear that out (I will dig these up later).

Edit:

The electrochemical oxidation potential for:

Fe(0) to Fe(II) is 0.447 V
For:
Fe(II) to Fe(III) is -0.771 V
and the reduction potential for H(I)/H(0) is of course 0 V

These values clearly suggest that H<sub>3</sub>O<sup>+</sup> is not capable of oxidising Fe(II) to Fe(III), as demonstrated by the dissolution of iron metal in dilute non-oxidising acids like HCl or H2SO4.

%%%%%%%

A change in colour can be due to many things. If it was due to oxidation by H3O+ then hydrogen would need to evolve.

I agree that "FeCO3" isn't useful.

Quote: Originally posted by Savior  
Do you know which colour are Iron(ll) carbonate, chloride, and most salts?
Not green, as most think! But white! That green color appears when Iron(III) starts forming. Because it is yellow/red and this one is white, when you combine that colors, it is green.

That's only transition color...temporary. Because Iron(II) is never stable.


This is nonsense, I believe. Anhydrous Fe(II) salts are often white but the colour of the hydrated (complexed) Fe<sup>2+</sup> ion is caused by a small energy difference between the e type and t type part filled 3d suborbitals, allowing the absorption of VIS photons and thus colour to arise. This energy difference arises from molecular orbitals binding the ligands to the central ion to exert different electronic repulsions on the e and t type 3d suborbitals, due to their different orientations in space.

It's perfectly possible to create a solution that is Fe(II) and free of Fe(III): just dangle some iron wool in it, this reduces all Fe(III) to Fe(II) and such solutions don't react to K4Fe(CN)6 (Prussian Blue is formed with Fe<sup>3+</sup> ions).

Such solutions (of FeCl2 and FeSO4) are green.

By contrast the hydrated Fe<sup>3+</sup> is almost colourless and gains colour only at lower pH due to formation of Fe(OH)<sup>2+</sup> (possibly solvated, not sure)

Quote: Originally posted by CHRIS25  
Comparing ferrous chloride/carbonate to ferrous sulphate what forces are at work whereby Oxygen is able to oxidize the ferrous to ferric so easily, yet in sulphate not so? The atomic orbitals and the structures are not giving anything away as far as I can see with my limited understanding. I've looked at everything I know, so what should I be looking at?


To make such a statement you need to compare ferrous chloride and ferrous sulphate in IDENTICAL conditions of concentration, pH and oxygen concentration. I'm therefore not claiming you're wrong but only that you don't present enough evidence for your assertion.


[Edited on 9-2-2015 by blogfast25]




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[*] posted on 9-2-2015 at 10:40


Are all transition metal compounds unstable or is it just iron compounds
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[*] posted on 9-2-2015 at 10:48


Quote: Originally posted by Ramium  
Are all transition metal compounds unstable or is it just iron compounds


It completely depends on the specific compound and the oxidation state of the transition metal in question. Generally iron(II) compounds are seen as unstable relative to iron(III), even though iron(II) sulfate seems wonderfully stable in my experience. Silver(II), copper(I), and highly oxidized manganese compounds are more examples of generally unstable transition metal compounds, although silver(I), copper(II), and manganese(II) and (IV) are usually pretty stable.




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[*] posted on 9-2-2015 at 10:55


Quote: Originally posted by Ramium  
Are all transition metal compounds unstable or is it just iron compounds


Stability/instability is a very relative and quite useless term here.

Despite protestations by Savior Fe(II) compounds (that can indeed be oxidised relatively easily) can be kept in 'good' conditions for years and more. Mohr's Salt is considered stable enough for use as a Primary Standard, for instance.

Some transition metal compounds like Nb<sub>2</sub>O<sub>5</sub> are among the most stable imaginable.

Rather than look at it your way, look at it like this. Take a compound XYZ.

If you can make this compound participate in a reaction, say:

XYZ + AB === > X + AY + BZ

and the Gibbs Free Energy change ΔG (from left to right) is negative (< 0) such a reaction can proceed and 'XYZ' could be said to be 'instable' in this specific context. And yet in many other contexts XYZ may resist attack and be considered 'stable'.

It's relative to context i.o.w.

[Edited on 9-2-2015 by blogfast25]




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[*] posted on 9-2-2015 at 12:19


Quote: Originally posted by blogfast25  

Quote: Originally posted by CHRIS25  
Comparing ferrous chloride/carbonate to ferrous sulphate what forces are at work whereby Oxygen is able to oxidize the ferrous to ferric so easily, yet in sulphate not so? The atomic orbitals and the structures are not giving anything away as far as I can see with my limited understanding. I've looked at everything I know, so what should I be looking at?


To make such a statement you need to compare ferrous chloride and ferrous sulphate in IDENTICAL conditions of concentration, pH and oxygen concentration. I'm therefore not claiming you're wrong but only that you don't present enough evidence for your assertion.


[Edited on 9-2-2015 by blogfast25]

Ooh, allow me to re-phrase then: The only evidence I have for this is that having made ferrous chloride many times and seen how easily it becomes ferric chloride in the presence of air, and then having made ferrous sulphate a few times for different purposes and seen how it does not oxidize that easily, ie high temperature and higher pH will do it quicker but then that will not lead to a fully oxidised ferrous sulphate either, I am really wanting to know why the oxygen can steal that extra electon from the Iron 2 so easily when bonded to carbonate or chloride, yet when the Iron 2 is bonded to the sulphate anion it resists oxygen's advances. The other piece of evidence is that lovely green ferrous sulphate sitting on my shelf.

[Edited on 9-2-2015 by CHRIS25]




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[*] posted on 9-2-2015 at 12:47


Quote: Originally posted by CHRIS25  
The only evidence I have for this is that having made ferrous chloride many times and seen how easily it becomes ferric chloride in the presence of air, and then having made ferrous sulphate a few times for different purposes and seen how it does not oxidize that easily, ie high temperature and higher pH will do it quicker but then that will not lead to a fully oxidised ferrous sulphate either, I am really wanting to know why the oxygen can steal that extra electon from the Iron 2 so easily when bonded to carbonate or chloride,


Remember firstly that FeCl2 and FeSO4 are ionic compounds and thus more or less completely dissociated in water:

FeCl<sub>2</sub>(s) === > Fe<sup>2+</sup>(aq) + 2 Cl<sup>-</sup>(aq)

The Fe<sup>2+</sup> ion is however solvated to [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>, a coordination complex with water as ligands.

Now it is possible (but I'm not affirming this positivistly) that some ligand exchange takes place whereby a small amount of either chloride or sulphate ions end up as ligands, replacing some of the H<sub>2</sub>O.

These new complexes may vary slightly in strength and this could very tentatively explain why the chloride appears to be easier to oxidise.

But personally I'd still prefer to confirm/infirm your conclusion experimentally, before further theorising.

One factor that affects the oxidation rate of Fe(II)to Fe(III) very significantly is the pH of the solution.

The formation of Fe(III) complexes like Fe(OH)<sup>2+</sup>, Fe(OH)<sub>2</sub><sup>+</sup> and Fe(OH)<sub>3</sub> are all favoured at higher than lower pH and they 'pull' the Fe(II)/Fe(III) equilibrium to the right. Comparing homemade ferrous chloride and sulphate may be difficult to achieve at equal pH. At a very minimum recrystallized materials would have to be used and maybe pH buffering.

[Edited on 10-2-2015 by blogfast25]




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[*] posted on 13-2-2015 at 01:06


Quote: Originally posted by blogfast25  


If the OP started from pure Fe(II) then in the absence of an oxidiser half of the Fe(II) cannot oxidise to Fe(III) and no Magnetite could form. Water alone cannot be responsible for the oxidation to Fe (II,III).
...
I believe you are wrong on this. H3O+ (not H+, which in water cannot exist) is not capable of oxidising Fe(II) to Fe(III). If it was then dissolving iron in any non-oxidising acid would yield a Fe(III) salt solution and not a Fe(II) salt solution as in reality it does. I think the oxidation half potentials for Fe(0) to Fe(II) and Fe(II) to Fe(III) bear that out (I will dig these up later). Edit: The electrochemical oxidation potential for: Fe(0) to Fe(II) is 0.447 V For: Fe(II) to Fe(III) is -0.771 V and the reduction potential for H(I)/H(0) is of course 0 V These values clearly suggest that H3O+ is not capable of oxidising Fe(II) to Fe(III), as demonstrated by the dissolution of iron metal in dilute non-oxidising acids like HCl or H2SO4. %%%%%%% A change in colour can be due to many things. If it was due to oxidation by H3O+ then hydrogen would need to evolve.


http://wikipedia.org/Iron
Under anaerobic conditions, the ferrous hydroxide
(Fe(OH)2 ) can be oxidized by the protons of water to form
magnetite and molecular hydrogen. This process is described by
the Schikorr reaction:

3 Fe(OH)2 → Fe3O4 + 2 H2O + H2
ferrous hydroxide → magnetite + water + hydrogen

The well crystallized magnetite (Fe3O4) is thermodynamically
more stable than the ferrous hydroxide (Fe(OH)2 ).
This process occurs during the anaerobic corrosion of iron and
steel in oxygen-free groundwater and in reducing soils below the
water table.
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[*] posted on 13-2-2015 at 10:29


Quote: Originally posted by Savior  
http://wikipedia.org/Iron
Under anaerobic conditions, the ferrous hydroxide
(Fe(OH)2 ) can be oxidized by the protons of water to form
magnetite and molecular hydrogen. This process is described by
the Schikorr reaction:

3 Fe(OH)2 → Fe3O4 + 2 H2O + H2
ferrous hydroxide → magnetite + water + hydrogen

The well crystallized magnetite (Fe3O4) is thermodynamically
more stable than the ferrous hydroxide (Fe(OH)2 ).
This process occurs during the anaerobic corrosion of iron and
steel in oxygen-free groundwater and in reducing soils below the
water table.


OK. I have no problem accepting that on thermodynamical grounds.

It would be useful and easy to establish first what precipitate occurs when a water soluble carbonate is added to a ferrous salt solution:

Fe<sup>2+</sup>(aq) + CO<sub>3</sub><sup>2-</sup>(aq) ===> FeCO<sub>3</sub>(s)

Or:

Fe<sup>2+</sup>(aq) + CO<sub>3</sub><sup>2-</sup>(aq) + H<sub>2</sub>O(l) ====> Fe(OH)<sub>2</sub>(s) + CO<sub>2</sub>(g)

The change in Gibbs Free Energy for the oxidation to Magnetite is obviously affected by the starting product.

[Edited on 13-2-2015 by blogfast25]




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