Iron (III) Sulfate (ferric sulfate) synthesis

Quote: Originally posted by woelen

Commercially available red/brown ferric oxide (Fe2O3) is amazingly inert and the same is true for the black oxide Fe3O4. I have both oxides from a pottery supplier and they do not dissolve appreciably, not even in hot 50% H2SO4 or boiling HNO3. It does dissolve in hot concentrated HCl, but only very slowly and a very large excess amount of acid is needed to get all of it dissolved. The solution in HCl becomes bright yellow, due to formation of the FeCl4(-) complex.

Quote: Originally posted by Hexavalent

'I must start with ferric oxide' - is this a school project, or an amateur thing? How flexible are the rules? Do you stir well when you add? How pure is the iron (III) oxide? What form is it in - granules, powder etc.? Have you tried adding it to hot sulfuric acid? [Edited on 26-6-2012 by Hexavalent]

on this webpage it shows that Ferric Sulfate is light gray, not yellow
(https://www.chemkits.eu/salts/143-ironiii-sulfate-hydrate-fe...)

So which color is right for ferric sulfate? (maybe anhydrous is yellow and xH2O is gray....)

sorry for 3 post in a row, but this is mine:

Quote: Originally posted by Bedlasky

It's still probably contaminated with Fe(II). I made some ammonium ferric sulfate from old FeSO4, H2SO4 and H2O2 years ago. When concentration of Fe(III) got high, it catalyticaly decomposed H2O2 pretty quickly, I added it slowly with stirring, but when I crystallize the product, ammonium ferrous sulfate start to crystallize first. Good thing is, that ammonium ferric sulfate is very soluble, while ammonium ferrous sulfate have low solubility. So I separated them by fractional crystallization. There wasn't very much of ammonium ferrous sulfate, but I still got few grams of this salt. So my recommendation is, make ammonium ferric sulfate from your ferric sulfate. It crystallize nicely (you can make big beautiful violet octahedral crystals if you want) and can be easily separated from any ferrous contamination.

Quote: Originally posted by woelen

Commercially available red/brown ferric oxide (Fe2O3) is amazingly inert and the same is true for the black oxide Fe3O4. I have both oxides from a pottery supplier and they do not dissolve appreciably, not even in hot 50% H2SO4 or boiling HNO3. It does dissolve in hot concentrated HCl, but only very slowly and a very large excess amount of acid is needed to get all of it dissolved. The solution in HCl becomes bright yellow, due to formation of the FeCl4(-) complex.

Quote: Originally posted by MMendoza14

I suppose I should have been more specific with my intentions... I'm attempting to model the process of sulfur species as a pollutant and then ultimately sulfuric acid and their effects on mineral dust in the atmosphere. Ferric Oxide (hematite) is a common mineral dust particle. I am using 99% pure hematite powder. Due to the attempt to model actual atmospheric conditions I'm already stretching a little bit using 1M as the pH never really reaches that low with a minimum usually around 3-4. Many of my experiments have shown that hematite does seem to be inert, but a reaction seems to occur after I put the solutions in the oven to dissolve the excess water. Before the oven there's the same red color and clear acid and after the oven theres a brownish solid leftover implying heat may cause the reaction to proceed but the product is unidentifable using ATR and XRD so far.

The problem is that I'm not sure if the laboratory grade hematite (or most commercially-available hematite) is a good model for naturally-occurring hematite dust. Wikipedia lists four different crystal structures for ferric oxide (rhombohedral, cubic bixbyite, cubic spinel, orthorhombic) and it may well be that one of these structures is more reactive to acids than the others, similar to how crystalline silica does not react with aqueous alkali hydroxide solutions but amorphous silica is moderately reactive.

The brown colour can be due to a variety of things. Personally, I've found that the iron(III) oxide I try making tends to cycle between being black at very high temperatures and red at room temperature, similar to NurdRage's video (https://www.youtube.com/watch?v=8ZQ67POLiqg). I have no idea why this occurs, but depending on the temperature in your oven, it might be the case that it's somewhere between those two states. Alternatively, it could also indicate the formation of some form of hydrated iron(III) oxide such as iron(III) oxide-hydroxide, which can be brown in certain polymorphs such as the beta form. Iron(III) oxides have many differing degrees of hydration and polymorphs, so it's difficult to be sure that what you have is the result of a reaction between the iron(III) oxide and the sulphuric acid.

I'd suggest placing the brownish sample into water and stir it to see if you can dissolve anything, then filter it off. If there was any sort of reaction, you'd get a clear orange-ish solution in the filtrate; otherwise it'd be practically colourless. Note, if your solution is cloudy or otherwise unclear, some particles have gotten through the filter and you need to use a finer filter. Alternatively, you can centrifuge a sample of the solution (or just leave it out to settle) to see if the supernatant is coloured or colourless.

In the environment you're trying to model, I wonder if there's also any effect from microbiology? As I understand it, there are several sulphur-oxidising bacteria which utilise the reduced sulphur species as a food source and can form sulphuric acid as their end-product. Such microbes would be able to live on hematite dust particles and would probably have developed strategies for mining iron out of such minerals (the strongest-known iron chelator, enterobactin, is a siderophore secreted by Gram-negative bacteria to acquire iron), or to otherwise alter their environment. Microbial activity could ostensibly increase the surface area of the hematite particles and increase their dissolution by sulphuric acid.



Regarding the synthesis of iron(III) sulphate in general: in my experience oxidation with hydrogen peroxide will often be incomplete even if you add an excess of H₂O₂. Sometimes this is because the hydrogen peroxide is not as concentrated as stated on the label, but more often, the H₂O₂ decomposes to H₂O and O₂ before it oxidises the iron (observable as effervescence upon addition of the hydrogen peroxide).

Interestingly, hydrated ferric sulphate is actually somewhat soluble in ethanol while ferrous sulphate is practically impossible. I've been exploiting this property to purify my samples of impure ferrous sulphate by dissolving it into a solution acidified with a small amount of HCl, filtering off undissolved ferric oxides and basic ferric sulphate (Fe(OH)SO₄), and then adding about 1.2x the volume of ethanol to precipitate the ferrous sulphate (e.g. if I had a 200mL solution, I'd add ~240mL ethanol).

Any iron oxides which react with the hydrochloric acid would form either ferrous or ferric chloride, both of which are highly soluble in ethanol. Ferric sulphate nonahydrate has a solubility of 12.7g/100g ethanol and 440g/100g water, so it can remain quite soluble in the aqueous ethanol solution. Ferrous sulphate heptahydrate has practically no solubility in ethanol solutions and precipitates out as a blue-green mass:



The ferrous sulphate can then be filtered off and washed with ethanol and dried. The reason I bring this up is because the aqueous ethanol filtrate presumably contains ferric sulphate with some oxide/hydroxide contamination, which could later be removed by recrystallising from water. Usually, I neutralise the filtrate and dispose of it, but it might be possible to use the same method to isolate iron(III) sulphate specifically from a mixture of partially-oxidised ferrous/ferric sulphate (although I think you'd need to swap out the HCl for sulphuric acid so you don't get ferric chloride contamination). Methanol might also be a good alternative to ethanol as ferric sulphate is presumably even more soluble in aqueous methanol, although I don't know how much ferrous sulphate might remain dissolved in such a mixture.