Sciencemadness Discussion Board

Trouble with anhydrous cobalt and ferric sulfates

Altreon - 20-12-2025 at 05:53

For the past few months, I've been trying to make a sample of every anhydrous metal sulfate I can make, minus the s-block sulfates. After failing to synthesize Ti2(SO4)3 and TiOSO4 I have successfully made ≈20g of anhydrous NiSO4 and FeSO4 by simply dissolving the corresponding metal by using double the stoichiometry of somewhat concentrated H2SO4 and boiling down the solution, which at some point instantly precipitates the anhydrous sulfates, as baby green and snow white products respectively. FeSO4 worked surprisingly well for something made of the very oxidizable ferrous ion, keeping my hopes up for the (deceivingly) simple Fe2(SO4)3.

I performed the same procedure for Fe2(SO4)3 as I did for FeSO4 but with the slow addition of 50% hydrogen peroxide, which would still leave an excess of H2SO4. I got to observe the interesting red color of Fenton's reagent, which stubbornly refused to oxidize the Fe2+ without heating, and even with that there was some iron powder stuck to the magnetic stir bar. After wiping that off, I boiled down the solution at >180°C until it stopped bubbling, but nothing precipitated. I added more H2SO4, nothing precipitated. I don't remember what I did, but after some amount of sulfuric acid and some ethanol, I crystallized some yellowish fluffy "crystals" in a brown ethanolic solution, which I filtered and heated to remove "water." Whatever I boiled off smoked intensely and somehow partially melted the sulfate, even though my hotplate was nowhere near the supposed point where Fe2(SO4)3 decomposes into oleum. It eventually got to the point where I had to boil it outside, leaving a somewhat less smoky brownish product that was sticky and cooled down to become a nice light pink (still somewhat clumpy) powder that looked like the image on Wikipedia without the brown bits. I did not bother to do anything with the remaining liquid since the sulfate was clearly not well behaved, and the yield was abysmal.
I thought that the lack of online data regarding the stoichiometry and structures of ferric sulfate hydrates wasn't concerning, but in hindsight I should've expected that if ferric sulfate really exists then many people would have written about it by now. I turned all the waste into purely brown Fe2O3, meaning I did completely oxidize all the iron to Fe3+. (Complete oxidation has been problematic in my attempted synthesis of [Fe(C2O4)3]3-.)

After this, I thought that maybe the problem was just with trivalent sulfates X2(SO4)3 and that all the divalent sulfates would be nice and kind to me, so I tried synthesizing CoSO4. I started with pottery grade "CoCO3," initially starting with stoichiometric H2SO4 but I added quite a bit extra after wasting my time trying to dissolve a suspended brownish black powder that apparently commonly appears in basic "CoCO3" which I filtered off. Boiling down the solution, I expected some beautiful reddish powder to begin crashing out, but for some reason the deep red [Co(H2O)6]2+ solution became a really nice, much less opaque hot pink solution that I also saw when dripping some CoSO4 soln. in the conc. H2SO4. I hoped that result was due to dilution, but only a mL of a purple solid precipitated out of the bulk mixture. Again I don't recall what I did (something involving ethanol), before I resorted to boiling down this very conc. H2SO4, I managed to crystallize some nicer red crystals that were most likely CoSO4•(6 or 7)H2O. Heating these crystals had the same vigorous smoking effect as the Fe2(SO4)3, leaving behind some rock hard solid mix of lilac and purple which required a metal spatula to break, nothing like the supposedly red CoSO4•(0 or 1)H2O. I think I once saw some very unpromising report saying that CoSO4 dehydrates at >350°C (my hotplate goes up to 280°C) which I tested by getting a small amount of the purple solid (I wasn't gonna use the lilac solid) in a test tube and blasting it with a blowtorch. This resulted in some liquid condensing at the top of the tube and the purple solid turning lilac, but nothing further happened. I believe the liquid was H2SO4 but I do not recall fuming. I have no idea why the cobalt sulfate isn't behaving nearly as well as the divalent sulfates of the elements right next to it, along with there being multiple sources describing anhydrous cobalt sulfate.

My main suspicion for these failures is that the excess of sulfuric acid in both cases crystallized out some acid salt (similar to those described in other threads on reacting NaHSO4 with EtOH to pull out H2SO4) with corresponding to the purple and brown/pink(thermochromic?) solids. I could not have avoided this with the very hydrolyzable Fe3+, but if I try this again with Co2+ I might be able to succeed. I would like to hear the explanations the rest of you guys can give regarding these failures.



[Edited on 20-12-2025 by Altreon]

bnull - 20-12-2025 at 06:39

Have you checked Mellor (A Comprehensive Treatise on Inorganic and Theoretical Chemistry), Brauer (Handbook of Preparative Inorganic Chemistry), or Sidgwick (The Chemical Elements and Their Compounds)? They're all in the Library. These books have methods of preparation for the salts, along with original references, and may give some clues about what happened.

Altreon - 20-12-2025 at 06:49

I checked every result in the internet archive copy of the first two books with keyword "ferric." Nothing relevant appeared in Brauer, while Mellor mentions solutions of ferric sulfate, the process of generating oleum from ferric sulfate, and the existence of a "basic ferric sulphate, say, Fe2O(SO4)2...." It also describes on page 543 that truly anhydrous ferric sulfate can be made by oxidizing ferrous sulphate solutions and evaporating them in the cold due to the formation of basic ferric sulphate when boiled. I do not see why this problem would occur in my solution with H2SO4 in large excess, even in the EtOH solution used to crash the sulfate out, and the fact that it fumed definitely does not make it look like a basic ferric sulfate.


About the scimad library: Any PDF I have downloaded from it ends up corrupted and as a ".doc" file.
Except for these books apparently. Wow.

[Edited on 20-12-2025 by Altreon]

Ferric sulfate extra details

Altreon - 20-12-2025 at 07:00

About the ferric sulphate, boiling the pure soln. actually did make a dense off-white precipitate. In fact, the whole solution along with the H2SO4 froze at once. I could not remove the ferric sulphate (acid salt?) so I redissolved it in water, requiring me to add so much water because the remaining chunks just would not redissolve, even after adding twice the soln. volume in water. I believe that boiling it down did not reproduce the same effect, and I could not get it to crystallize again without EtOH, at which point the yellow product was only a fraction of the ≈75mL of white acid crystals that initially formed. Maybe the extra water was all the Fe3+ ion needed to decide not to behave nicely again (i.e. form an inseparable complex like Cr3+ does).

[Edited on 20-12-2025 by Altreon]

bnull - 20-12-2025 at 10:47

Quote: Originally posted by Altreon  
About the scimad library: Any PDF I have downloaded from it ends up corrupted and as a ".doc" file.
Except for these books apparently. Wow.

Use another PDF reader. Brauer used to break Acrobat in my PC. As for the .doc issue, it may be related to your browser configurations (try Firefox or Edge if you've been using Chrome).

If it still persists, manually change the extension to PDF. There's also the chance you've downloaded a DjVu file, in which case you need a DjVu reader (DjVu Libre for PC, DjView for Android).

teodor - 5-2-2026 at 13:50

Sulfates except of ammonia are insoluble in anhydrous acetic acid. Ammonia sulfate is slightly soluble (less than 1:100). Chlorides are very soluble except of ammonia chloride which is slightly soluble. So, there is a possibility of getting anhydrous sulfates by metathesis reaction in acetic acid.

There could be some compilcations. For example FeCl3 reacts with AcOH with evolution of gaseous hydrogen chloride.

3FeCl3 + 7AcOH -> 6HCl + [Fe3(OAc)6]Cl3 * HOAc

But I suppose the Fe(III) complex should be as good as FeCl3 for getting Fe(III) sulfate by metathesis.

semiconductive - 8-2-2026 at 17:41

I'm trying to understand what you actually saw.
Quote:
I got to observe the interesting red color of Fenton's reagent, which stubbornly refused to oxidize the Fe2+ without heating, and even with that there was some iron powder stuck to the magnetic stir bar.


You mean the solution did not turn red before it was heated, or do you mean that it was red and turned brown after heating?

eg: I think Both a red color or a mix of brown colors are supposed indicate ferric state...

How long did you wait for the solution to change color before heating it?

semiconductive - 8-2-2026 at 18:28

I'm no expert.

But I do know ferric Iron is below hydrogen in the electroactive series.
In most electroplating situations, iron will want to stay in solution more than hydrogen will.

From standard oxidation/reduction tables,

Fe²⁺ + 2e⁻ ⇌ Fe(s) -0.44 [V]
Fe³⁺ + e⁻ ⇌ Fe²⁺ +0.771 [V],

I think this means is that weak hydrogen ions have a hard time removing hydroxide ions from ferric iron.

eg: It might be easier for the hydroxide you've supplied via hydrogen peroxide to decompose into oxygen/rust than for it to become water and 'move aside' so that a sulfate may displace the hydroxide ion that is complexing Fe³⁺ ? (A guess, with some education behind it. )

I imagine:
The structure of (Fe³⁺)₂ (SO₄²⁻)₃ has to have two different bonding types that's will crudely form a linear molecule. Each iron is bound twice to an outer sulfate, and once to an inner sulfate, for a total of three bonding electrons on each iron.

SO₄ : Fe · SO₄ · Fe : SO₄.

However, the strength of sulfuric acid's ionization constant is different between the first and second ionizations. eg: Therefore, the middle SO₄ molecule will have different difficulties in displacing a first hydroxide from the left iron, and a second hydroxide from the right iron.

Isn't Hydroxide is one of the strongest potentials in an oxidation reduction table?
It's hard for me to imagine that both hydroxides on different iron atoms would be removable by one sulfate ion having two different hydrogen ionization strengths.

? You might have better luck using a reagent that isn't as strong an oxidizer as hydrogen peroxide ?

DraconicAcid - 8-2-2026 at 21:54

Quote: Originally posted by semiconductive  
I'm no expert.

But I do know ferric Iron is below hydrogen in the electroactive series.
In most electroplating situations, iron will want to stay in solution more than hydrogen will.

From standard oxidation/reduction tables,

Fe²⁺ + 2e⁻ ⇌ Fe(s) -0.44 [V]
Fe³⁺ + e⁻ ⇌ Fe²⁺ +0.771 [V],

I think this means is that weak hydrogen ions have a hard time removing hydroxide ions from ferric iron.

No, it means that iron(II) ions will not reduce H+ ions to hydrogen gas. It has nothing to do with acid-base chemistry or the solubilities of the hydroxides.
Quote:
eg: It might be easier for the hydroxide you've supplied via hydrogen peroxide to decompose into oxygen/rust than for it to become water and 'move aside' so that a sulfate may displace the hydroxide ion that is complexing Fe³⁺ ?

Hydrogen peroxide does not contain hydroxide ion.
Quote:
I imagine:
The structure of (Fe³⁺)₂ (SO₄²⁻)₃ has to have two different bonding types that's will crudely form a linear molecule. Each iron is bound twice to an outer sulfate, and once to an inner sulfate, for a total of three bonding electrons on each iron.

SO₄ : Fe · SO₄ · Fe : SO₄.


It's amazing how fluently you go through the most mathematical models of chemistry, but still haven't grasped that ionic compounds are not molecules. In anhydrous ferric sulphate, you will have iron ions surrounded by sulphate ions, and sulphate ions surrounded by iron ions. None of the sulphate ions are distinct from any other, and all the iron ions are equivalent.

In aqueous solution, you'll have hydrated iron(III) ions (some of them with hydroxo ligands except at very low pH) and sulphate ions, without any bonding or coordination between them.

I know there are texts that describe Fe2O3 as O=Fe-O-Fe=O, but these are ones that predate Pauling.
Quote:
Isn't Hydroxide is one of the strongest potentials in an oxidation reduction table?


No. Hydroxide isn't a strong reducing agent or a strong oxidizing agent.

semiconductive - 9-2-2026 at 09:46

Quote:
No, it means that iron(II) ions will not reduce H+ ions to hydrogen gas. It has nothing to do with acid-base chemistry or the solubilities of the hydroxides.


Oh, and there are no H+ ions in solution around the anhydrous sulfate only hydronium according to you. Yet here you describe H+ ions?
Quote:

Hydrogen peroxide does not contain hydroxide ion.


Hydrogen peroxide is H₂O₂. Or HOOH.
When I look up how it reacts with Fe²⁺ online, I am told it breaks into OH⁻ and OH neutral.
Fe²⁺ + HOOH ⇌ Fe³⁺ + OH⁻ + OH

There are only two possibilities, misinformation exists on the web as to what hydrogen peroxide is/does, or it dis-associates into hydroxide ions. If it's the former, then it's not my fault. If it's the latter, you're correcting someone over an issue that is ambiguous. I don't come with a way to tell good information from bad information. I do come with ways to tell bad math deductions from good math deductions.
Quote:

It's amazing how fluently you go through the most mathematical models of chemistry, but still haven't grasped that ionic compounds are not molecules. In anhydrous ferric sulphate, you will have iron ions surrounded by sulphate ions, and sulphate ions surrounded by iron ions. None of the sulphate ions are distinct from any other, and all the iron ions are equivalent.


Oh, I did say that chem lab was the only subject in College where I got below a "C" grade.

I'm glad that I entertain and amaze you. I apologized for my faults in advance when I said "I'm not an expert."

But, seriously, what is "indistinguishable" ?:

example:
In quantum mechanics, how do you know there are two electrons if they are totally indistinguishable?
How can you 'count' electrons without putting the labels '1' and '2' on them?


Therefore, what are you saying -- that there aren't "three" Sulphate Ions for every "Two" iron ions?
What exactly does "indistinguishable" mean?


You can repost the picture of sodium chloride that you showed me earlier.
I will just note that the "sodiums" can be numbered 1,2,3,4... in the picture.
It's wrong to number them 1,1,1,1,1,1 .....

2nd) The word indistinguishable implies an exclusion of isotopes of sulpher in any of the sulphate molecules. Is that guaranteed? I'm being inclusive of possibilities -- you're being exclusive.

I'm excellent at math, I'm not great at double talk that isn't numerically backed.

When you say stuff, it has to mean something.

Lord Kelvin paraphrase -- "If you can't assign a number to it, you don't really understand what you're talking about"
Quote:

In aqueous solution, you'll have hydrated iron(III) ions (some of them with hydroxo ligands except at very low pH) and sulphate ions, without any bonding or coordination between them.

I know there are texts that describe Fe2O3 as O=Fe-O-Fe=O, but these are ones that predate Pauling.


Yes, there are. No argument against your knowledge of history, here.

But my example is also about the order of the chemical reaction.
How do the first two irons in solution, combine with the first two sulfates, before becoming a tri-sulfated molecule?

It's not a simple question (to me).
This is WHY people used to think the linear O:Fe·O·Fe:O was the stable form of rust, rather than a ring structure with two capping irons.

Or are you arguing the reaction order is equal to the number of oxgens and irons that are in the final product?

To help me understand:
Define the words bond and coordination, numerically. At what value does something have a bond, at what value is it 'co-ordinated'?

You have said something here, but I don't understand what it is.
:shrug:


[Edited on 9-2-2026 by semiconductive]

bnull - 9-2-2026 at 14:20

No flame wars here, please. Wrong forum, wrong time.
Quote:
When I look up how it reacts with Fe²⁺ online, I am told it breaks into OH⁻ and OH neutral.

I suppose you mean this paper: Chemistry of Hydrogen Peroxide Formation and Elimination in Mammalian Cells, and Its Role in Various Pathologies (https://doi.org/10.3390/stresses2030019).

semiconductive - 9-2-2026 at 17:41

Quote:
I suppose you mean this paper: Chemistry of Hydrogen Peroxide Formation and Elimination in Mammalian Cells, and Its Role in Various Pathologies


That's one of many. It's not limited to life sciences.

I do not claim to understand why the OH· is neutral in some cases and OH⁻ in others.

In my easy sulfite ion thread, I mentioned the Russian differences in considering hydrated object that are perhaps not molecules vs. USA articles.

See equations (14)(15) where he begins to mention hydrogen peroxide.

https://www.researchgate.net/publication/276498338_Electroch...

No one has clarified for me the difference.

I have no way to sort chemistry information by accuracy at my level of understanding.
I do remember what people tell me and show me.

There is no word given to me to describe FeSO₄ as a ferrous thing, vs alternate and distinct stoichiometries that crystallize; eg: before crystallization, since I am talking about the point where it *first* forms in aqueous solution -- not the point where it is fully crystallized. The only word I know is 'molecule'.

I am inclined to use the same word when discussing iron phosphate (vivianite crystals) as I would iron sulfate (whether ferric- or ferrous- ). This is what I am told a molecule is:
Quote:
A molecule is the smallest particle of a substance that retains all the physical and chemical properties of that substance


There is no mention of it being different if it is charge neutral or not.

eg:
A single Na atom and Cl atom bonded together, will have the properties of a particular salt -- even though it's not a crystal structure (yet).

So I don't see anhing wrong with calling a structure of five or less atoms a "molecule" AKA: three sulfur atoms covalently bonded to oxygens and bonded in any way whatsoever to two irons ?

I'm not talking about a thousand sodium atoms bonded to a thousand chloride atoms; nor a thousand iron atoms bonded to a large number of sulfates.

Nor can properties of crystals be used to differentiate the definition in this thread.
For that would require a different name for crystals of any geometry based on the same atoms that differed in electronic properties; layer thick planar crystals have different QM properties than 3D crystals. etc.

It's not reasonable to expect me to know all the possible exceptions.
Which is why I take this as derogatory bigotry:
Quote:

It's amazing how fluently you go through the most mathematical models of chemistry, but still haven't grasped that ionic compounds are not molecules.



[Edited on 10-2-2026 by semiconductive]

bnull - 10-2-2026 at 03:46

Quote:
That's one of many. It's not limited to life sciences.

One of my browsers is less biased than the other and that's what came first.

teodor - 10-2-2026 at 07:11

I try to follow this thread, but it seams nobody still mentioned that it is impossible to convert Fe(III) hydrate to anhydrous salt by heating, which is obvious thing to many who is reading this thread, but probably this is exactly the question which was asked at the beginning if I would omit the other preparative details. No, started with water or H2SO4 (=H2O + SO3) or H2O2 = (H2O + O) it is impossible to get Fe(III) sulfate anhydrous no matter what you plan to do because you always will finish with more sulfur-free oxigen connected to Fe(III) than SO4(2-) and if you would ask "why" I assume just because Fe(III)-O bond is much more stronger than the 4th S(VI)-O bond.

semiconductive - 10-2-2026 at 21:53

Quote:
I try to follow this thread, but it seams nobody still mentioned that it is impossible to convert Fe(III) hydrate to anhydrous salt by heating, which is obvious thing to many who is reading this thread


I am a nobody, I guess?
I said:
Quote:
eg: It might be easier for the hydroxide you've supplied via hydrogen peroxide to decompose into oxygen/rust than for it to become water and 'move aside' so that a sulfate may displace the hydroxide ion that is complexing Fe³⁺ ?


To which draconic acid replied:
Quote:
Hydrogen peroxide does not contain hydroxide ion.


( And, technically, I never said it did contain an ion. )

Teodor, do you have any idea why the Russian/Moscow article I linked talks about OH· separate from OH⁻ ? Is this a translation issue?

I think you are on to something. Because Draconic Acid reccomended that I read "Advanced inorganic chemistry" (cotton and Wilkerson), and the question I asked D.A (privately) was about iron complexed with water molecules.

To which he replied,
Quote:
I can address your third question. The m/n oxyacid rule applies to neutral (uncharged) molecules. The hydrated iron ions are ions. Fe(III) simply has a higher charge, so it's less able to hold onto positively-charged H+ ions. If you want to use the oxyacid rule, you'd need to consider the hydroxides FeO(OH) and Fe(OH)2.


I think sufluric acid H₂SO₄ is a molecule. It's covalent, because (in general) there is no metal in it. And this is the reagent being used in the opening post's description of the experiment.

Cotton and Wilkerson say, pp. 44, section 1-7: "Historically it has been customary to treat coordination compounds as a special class separate from molecular compounds. On the basis of actual fact only, i.e. neglecting purely traditional reasons for such a distinction there is little, if indeed any, basis for continuing this dichotomy."

So, I really don't get why he's on my case.

I think using a less strong oxidizer than H₂O₂ is not guaranteed to fail. eg: chlorine, iodine, or something along those lines might be displaced by a sulfate "molecule". But, I don't know what conditions would make this a likely outcome.

teodor - 11-2-2026 at 02:56

I promise to check those publications later, I don't have enough time now for that. Still I think I can write something possible useful using only few minutes.

I don't quite understand why you are talking about the hydroxide ion only. The oxygen itself is the second after fluorine by electronegativity, and it is the reason of all troubles of getting anhydrous compounds starting from chemicals of water group. Water is just a common form of oxigen. H2SO4, H2O2, Fe2O3 are another forms of water (or oxygen), if one would think in the term of Franclin's theory of solvent systems. The oxygen is the enemy of getting anhydrous compound, and it can be not only in the form of OH- group.

Fe2(SO4)3 is a tricky compound because it is both anhydrous and water-derived (because H2SO4 itself is water-derived, as any oxygen-containing acid). There are some unique methods to get it, e.g.:

1) 2 FeS2 + 2NaCl + O2 -> Fe2(SO4)3 + Na2SO4 + Cl2

2) 2 FeSO4 + 2SO3 -> Fe2(SO4)3 + SO2

3) FeSO4 -> Fe2(SO4)3 + 2 Fe2O3 + 3SO2

All those methods use oxydation of Fe(II) to Fe(III) (by sulfur or oxygen) using excess of S atoms. S could be replaced by O but O couldn't be replaced by S in inorganic compounds, so you can go from sulfur to oxygen but never from oxygen to sulfur that's why

Fe2O3 + SO3 -> Fe2(SO4)3 + ...

probably will never work.

As I see you use quite different theoretical frame to analyse the process, so I will try to understand your way of thinking before writing anything more.

P.S. I think the difference is that you are give too much attention to the coordination as an obstacle of getting the anhydrous salt. It is not a coordination. Those bonds are not so strong, and normally should gone with heating. This is oxygen which makes the things impossible. Combine this with the S-O replacement rule I mentioned which prevents to make compound free of unnecessary oxygen. Heating hydrates you will get Fe2O(SO4)2 + F2O3 + H2SO4 vapours. O would be the last atom to leave Fe(III) compound.

In the potential process of getting sulfate in AcOH oxygen is locked by by dehydrating power of Ac2O. Still, Fe(III) could be more eager for oxygen, but that means the reaction still worth to study (if you will not get anhydrous sulfate you will get Ac2O or AcCl which is not a bad result if you think about it).

P.P.S. Probably you can succeed only using proper S to O ratio in a starting combination. For that reason, if you like H2O2 I would replace H-O-O-H with [O3S-O-O-SO3]2- ion (it is quite similar, isn't it) and explore some methods based on e.g.

FeSO4 (anhidrous) + K2S2O8 -> ?

Theoretically K2S2O7 could work better because it is even more S per O, but practically getting pure K2S2O7 (free from water and the acid) is a very big challenge.


[Edited on 11-2-2026 by teodor]

semiconductive - 12-2-2026 at 11:59

Quote:
I don't quite understand why you are talking about the hydroxide ion only. The oxygen itself is the second after fluorine by electronegativity, and it is the reason of all troubles of getting anhydrous compounds starting from chemicals of water group. Water is just a common form of oxigen. H2SO4, H2O2, Fe2O3 are another forms of water (or oxygen), if one would think in the term of Franclin's theory of solvent systems. The oxygen is the enemy of getting anhydrous compound, and it can be not only in the form of OH- group.


Yes. Elemental oxygen O, is the second most electronegative.
Nitrogen and Chlorine tie-ing for third place.

I am not familiar with Frranclin's theory, and it's not in the theoretical part of Cotton and Wilkerson which I've just read on Draconic Acid's recommendation. My not thinking about oxygen electronegativity by itself probably has more to do with how often certain ideas show up in U.S. vs. (translated) Russian literature that I can find.

For example, in US literature, I see people talk about diatomic oxygen O₂ and hydrogen H₂ dissolving in water. They seldom (never?) talk about hydrated "oxygen" or hydrated "hydrogen" atoms that are dis-associated. But, in the Russian article that I linked, he starts out by talk about H₃O (uncharged) and OH (uncharged).

eg: Equation (5) is about hydrated mon-atomic hydrogen vs. hydronum ion as distinct species that exist in solution.

Likewise, I know the covalent superoxide bond O₂, is not very strong and is easily broken in spite of the electronegativity of oxygen. But, if there was a lone oxygen atom in water becasue a diatomic oxygen dis-associated, then hydrating it would mean we get: H₂O·O (see text after equation 15) this would have the same empirical formula as hydrogen peroxide.

But, I don't know if it's H-O-O-H or [H-O₂-H]⁺ where the extra oxygen is in the O⁺ state, like happens with fluorine.

Notice, Draconic acid when talking about water hydrated ferric ions, said that FeO(OH) is a hydroxide. But I don't know how the oxygens are bonded to each-other, or not. ( He also didn't show Iron's oxidation state. ) But, He's just shown a formula where O(OH) is called a hydroxide.

I think the easiest molecular changing is to gain or loose a hydrogen. From hydrogen peroxide, then:

H₂O₂ + H₂O ⇋ H⁺·(H₂O) + O₂H⁻

So, I think:
H₂O₂ + [Fe·SO₄]⁺(aq) ⇋ H⁺(aq) + Fe·SO₄·O₂H

Which makes a negatively charged hydroxide ligand (according to D.A's explanation if I try to balance it) or a super-oxide-hydroxide.

But, it's also possible since the superoxide bond is supposedly weak:
H₂O₂ ⇋ HO + OH

That D.A. meant that the molecule is ferric monoxide, and that a charge netural hydroxide attaches to it.

In such a case, the question becomes how likely is the electronegativity of two oxygens near to each other going to cause the hydroxide (neutral) molecule to dis-associate.

I do not have any numbers to estimate the relative liklihood of these two dis-associations. Simple ignorance on my part.

Since I do not know what is more likely, and have no experience, I am simply stating things which obey the rules that I do know. I don't guarantee they will work, but I also am not going to claim that all such reactions are impossible.

eg: I am including possible reactions but not guaranteeing them, while draconic is excluding without clear explanations.

I will have to look at the rest of what you wrote, and think about it for a while.
Thanks for your quick answer, Teodor.



[Edited on 12-2-2026 by semiconductive]

teodor - 12-2-2026 at 14:22

I believe N is slightly more electronegative than Cl, according to Pauling’s study in 1932, but I never checked up-to-date publications about that.

Edward Curtis Franklin, “The nitrogen system of compounds”. One of the most important chemical books published in 20 century. In a nutshell it shows how N would behave in place of O, for this purpose liquid NH3 is used as a solvent and the difference of reactions with water chemistry is shown. Then the idea of “solvent family” is introduced and the common properties were shown. The examples of water group solvents are : H2O, alcohols; of ammonia group: liquid NH3, amines. By this theory KNH2 is analogue
of KOH, K3N - K2O etc. There are different acids and bases - for example NH4Cl is the analogue of HCl solution in water. During 20 century this theory gave many new synthetic methods. The hydrate is the result of reaction in water, so if you have to avoid hydration those methods suggest choosing another type of solvent for the reaction.

Cotton & Wilkinson is a very good book but it doesn’t cover all topics of the inorganic chemistry being only a very small window in a house. I believe Draconic Acid pointed there because there is description of ionic crystal structure.
If you can find it, you can also look into A.F. Wells “Structural Inorganic Chemistry”. This is more detailed description of the topic. You will see that the set of possible compounds in a solid form is a result of different parameters like ionic radii, bond lengths, inter bond angles etc. By some of those reason if you try to pack SO4(2-) ions with very small and charged Fe(III) you have to avoid O-, otherwise it always find the place between them (the result of their charge and the geometric forms of other ions). Wells is really interesting books which shows how crystal chemistry works.

There is a publication https://doi.org/10.1021/ic50193a021 about crystalline structure of Fe2(SO4)3. By the way, if you need some method of getting anhydrous salt you can always check publication about its crystal structure, they give the method of preparation combined with the crystallographic proof. So, the authors use FeSO4*7H2O reflux in H2SO4. H2SO4 works both as an oxidizer and a solvent to get well-formed crystals. This is also the result of applying of “solvent-system” theory I mentioned before.

H3O and OH are uncharged radicals. OH, for example, was discovered in 1924 after studying spectrum of some distant objects in space. The Russian article studies the microstructure of water caused by electric fluctuations when charges of different neighbour molecules could be different. I think this article is not applicable here because a macroscopic result is the result of thermodynamic equilibrium, which is not dependent on electron fluctuations the same way like the movement of a crowd is not dependent on a movement of a single man no matter how he tries to move in this crowd.

FeO(OH) is probably gamma-Fe2O3 * H2O. I have some book about Fe oxides and can find more information about the exact structure if you have interest, but also you can try to search “hydrated Fe(III) oxides” in google.

There are also Fe peroxides where probably oxygen can be linked to oxygen like in H2O2 (in the other case why they are named “peroxides”?).

When Fe2O3 hydrate is dissolved in some solvent (and you can dissolve it in water with addition of e.g. EDTA or tartaric acid) it doesn’t keep the crystal structure and no one compound keeps the crystal structure in dissolved form. So, the structure of salt and the structure of solvate specie is not the same.

H2O2 isn’t stable in the presence of transitional metal ions, it decomposes through catalytic process of oxidation-reduction which is not well studied yet I believe, or, which is almost the same, the topic has so many publications that nobody made a summary of what we already definitely know here. But we definitely know that few micro (nano?) grams of transition metal can cause decomposition of several liters of H2O2 (for this purpose commercial H2O2 is stabilized by additives). Based on this I wouldn’t agree with your proposed formula for FeSO4 / H2O2 interaction.

There is no role of OH radicals in FeSO4 reactions, no powerful magic, at least nothing comparable to the level of “AJKOER” spell.

There is (hexa?)aqua complex of Fe(III) in water solution as an ion and the oxygen from this complex is too close to Fe(III), and this oxygen goes to crystals in the form of water and then when you try to dehydrate by heating H2SO4 leaves the crystal leaving oxygen on Fe(III).

That's why the experimentator observed the white smoke below the decomposition point of Fe(III) sulfate.

P.S. I was not right! I've just checked the behaviour of Fe(2+) and Fe(3+) aqua complexes in the book "Mechanisms of inorganic reactions. A study of metal complexes in solution" Basolo & Pearson:

1. Initial step of the reaction between Fe2+ and H2O2 is the formation of OH radical!
2. Fe(II) aqua ion may act as a reducing agent by transferring a hydrogen atom from its hydration shell to a substrate while Fe(III) aqua as an oxidising agent by transferring hydroxyl radical to a substrate.

So, bravo to your intuition. There is a place to OH in the oxidation mechanism. Still, it has no relation to the complexities of getting anhydrous compound from water solution.

[Edited on 13-2-2026 by teodor]

semiconductive - 12-2-2026 at 18:40

Quote:
I think this article is not applicable here because a macroscopic result is the result of thermodynamic equilibrium, which is not dependent on electron fluctuations the same way like the movement of a crowd is not dependent on a movement of a single man no matter how he tries to move in this crowd.


I can understand why you would think that; but it's not so.

Semiconductor physics use the exact same equations, and yet they are used to predict non-thermal equilibrium effects in active devices such as transistors that operate millions of times a second. There isn't time for thermal equilibrium to be reached. None the less: Diffusion lengths of single electron carriers is routinely used to predict non-equilibrium effects.

I have a simulator, that accurately predicts semiconductor behavior in non-equilibrium conditions.
I would love to be able to apply it to chemistry, but I don't quite have enough information.

eg:
The paper I linked is not sufficient for me to do a full simulation of water.
But, it just happens that the present thread touches on ideas that I've been trying to understand; so I made a comment.

Note: It's the hydroxide radical that I'm remembering has the extreme reduction voltage.

OH· vs. the S.H.E. is +2.8 Volts. The question I asked earlier, rhetorically, was me remembering wrong. OH⁻ has a potential of less than a volt.
Quote:
I believe Draconic Acid pointed there because there is description of ionic crystal structure.


maybe, but it's Ironic, then, that Cotton and Wilkerson say that there is no good reason to call ligand or coordination chemistry non-molecular.

Metal ions can be complexed by ligands, can't they?

Al's stated initial conditions of the reagants are as sulfuric acid (therefore molecules) and hydrogen peroxide (therefore molcules) and iron metal (some of which sticks to a stir bar, possibly a crystal -- but not ionic, just metallic.)

The reaction to make the first ferric ion, has to happen from a sequence of molecule interactions.

Why would D.A. want me to focus on ionic crystals, when those only form after the reaction has completed?

Further, why bring up the "all of them are identical" when that's patently not true.

If someone has two atoms, these atoms are distinguishable by the fact that there are two of them. If they were completely indistinguishable, you couldn't count them. ( You couldn't even write down TWO quantum mechanical equations for them. )

The same is true no matter how you group the atoms in an ionic crystal. There are co-ordinates (cartesian) that are used to describe crystal cells that distinguish the number of atoms (whether ion or not) and their geometry.

I get that two sulfur sodium atoms might behave chemically in the same way.
But, just two things are "like" each other in one way, doesn't mean that they are identical in every possible way. To claim they are, is to make a religious generalization (a Bohr religion) out of partially analyzed experimental data.

It doesn't make sense to me, scientifically.

Do all sulfur atoms and iron atoms in water suddenly and simultaneously move to their positions in a crystal and then bond in parallel (exactly the same) -- or do they randomly, and singly, interact with one iron here and then one iron there dissolved in water, building up chains of ions ?

I think it has to the be a sequential event, because high order reactions with a 100 irons simultaneously combining with 300 sulfuric acid are said not to happen due to entropy. At least, that's what my thermodynamics teacher insisted was the case 30 years ago.



[Edited on 13-2-2026 by semiconductive]

teodor - 13-2-2026 at 01:15

Look the chapter 2 of Basolo & Pearson, it describes the existing theories of coordination bond which could be not ionic and not molecular at the same time. There are a lot of mathematical and quantum physics details there and the real picture doesn't correspond a simplified model used in most assumptions. I would omit discussion of this in this thread because it has no relation to the original question which is actually only the question about removing a "hydration shell" during the reaction. But I can send you an electronic copy of the book / chapter if you would like to study this topic and doesn't have the way to find it.

As a side note I'd like to say that H2O2 could be not a good oxidiser for Fe2+ because it starts a chain reaction Fe2+ <-> Fe3+. I was not right the there is no good summary of the theory of Fenton's reaction. Those 2 publications could be considered as the recent review of the theory:

https://doi.org/10.1179/135100001101536373
https://doi.org/10.1179/135100002125000190

The first article is named "The Haber-Weiss cycle—70 years later" and the second "The Haber-Weiss cycle—70 years later: an alternative view" (by different autors) suggesting that still there is no agreement on some topics.

[Edited on 13-2-2026 by teodor]