Sciencemadness Discussion Board - Preparation of Ammonium iron (III) sulphate (ferric alum)

Preparation of Ammonium iron (III) sulphate (ferric alum)

chemoleo - 15-4-2006 at 18:39

Ferric alum (NH4Fe(SO4)2), of similar structure to ordinary alum, (potassium aluminium sulphate, KAl(SO4)2), is of interest because it forms (sometimes?) violet to clear crystals that are supposed to be suitable for crystallisation trials.

One established method for making it is via boiling of ammonium sulphate and ferrous sulphate (II) togther with nitric acid, which functions here as the oxidiser to convert Fe2+ to Fe3+.
A preparation is detailed on Lambdasyn (german). However, this prep uses precious HNO3, and it evolves NOx during the oxidation - neither which I want to waste.

Thinking, I am cleverer than this, I tried the same thing with H2O2, with interesting, but unexpected results:

Reaction:

K2SO4 (or (NH4)2SO4) + 2 FeSO4 + H2SO4 ----H2O2----> 2 KFe(SO4)2

I used 0.01 moles initially, all in stoichiometric amounts, including the H2SO4. I even checked the pH before and after the reaction, it changed from 0.3 to 0.6, so very little, and still damn acidic. For the H2O2, I used 30%, in a total volume of ~100 ml for each reaction.
I first dissolved all the salts, and the H2SO4. The solution was nice and clear, with a smaragd green tint (1). Adding a total of 13 ml of H2O2 (which works out as a slight stoichiometric excess of H2O2 per oxidation to be done). Already after the addition of 1 ml, the solution immediately turned yellow. After adding the whole 13 ml, the solution was bubbly, and quite intensely yellow (2), it also became handwarm. - suggesting that a reaction was certainly taking place. Thinking, let's really drive this reaction to the right, by heat, and also to remove the left-over H2O2, I heated this in the microwave until the foaming stopped, and true boiling began. By this point the solution was dark red/brown (3). In the case of the potassium salt, small yellow crystals could be seen, which covered the bottom of the beaker alongside with a brownish sludge, which was thankfully very little. When this cooled down, the solution turned yellow again, but the precipitate remained (4). Interestingly, the complex displayed heat-sensitive properties, in that the hot solution was dark-red while the cold solution was a clean yellow - an entirely reversible process. I discovered, however, that overheating this for a long time (I tried to boil down the whole solution) in the microwave yielded a brown preciptiate, which would not redissolve - suggesting it is some FeIII oxide/hydroxide. The ammonium salt seemed more stable, although I didnt push it by harshly boiling it in the MW.
Anyway, with the K-solution, I precipitated the soluble leftover (not the yellow crystals/brown sludge (4)) with ethanol, an immediate yellow precipitate formed, which turned the solution thick. Upon centrifugation, I got a nice cream-yellow voluminous pellet (5) that happily redissolved in water, forming a dark yellow solution again.
Even more interestingly, after the oxidation of the K-salt with H2O2, the yellow crystal leftover, as well as the brownish sludge, happily redissolved upon adding more water, no residue left. So clearly, This is not an insoluble hydroxide/oxide.

Right now I have this sitting in crystallisation trays, in the hope of getting some crystals from it. I don't, however, think that this is ferric alum, because the colour is wrong, and the crystals were clearly yellow, but not violet or clear (apparently super-clean ferro-alum crystals are clear, the colour derives from manganese impurities), according to the lambdasyn description. Neither were they octaedral, but rather very small and needle-like, although this is hard to judge.

The question is, what did I get? There aren't that many variables to this, but it clearly is a new salt, which is not some crappy non-dissolvable oxy-hydroxide.
Ideas?

I will be trying the same thing with HNO3, for reference and comparison, in a few days. I'll be darned if that solution turns violet!

Alternatively, I'll try and turn the FeSO4 (II) into Fe2(SO4)3 (III) by direct oxidation with H2O2 itself, and see what I get. Iron III sulphate should yield the ferro-alum once ammonium sulfate or potassium sulfate are added.

guy - 15-4-2006 at 20:09

Shouldn't the hexaaquoiron(III) ion be almost colorless? If its a different color, does that means some kind of complex is formed?

BromicAcid - 16-4-2006 at 12:18

A little unsure as to what you formed, but although it's obvious it should be noted that Fe(II)/H2O2 is a Fenton's Reagent type of situation and it is able to oxidize aqueous NH4+ to nitrogen and water. Here is a few pages on alums that I found in my chem book 'Descriptive Inorganic Chemistry' Third edition, Rayner-Canham Overtone (you already have this though).

alums.jpg - 322kB

woelen - 17-4-2006 at 22:53

What you created is an iron (IV) species, or even an iron (V) species. There still is quite some doubt on the precise nature of the material you made, but as BromicAcid already pointed out, Fenton's reagent is the key to understanding what is going on.

On my website I have a description of what is going on, together with a link to a paper which covers this subject. Look at the following page on iron salt solutions and especially the part on iron (IV).

Also, iron (III) ions are almost colorless at such a low pH, without complexing agent. With chloride present, the color is intense yellow. This also is explained on the webpage.

http://woelen.homescience.net/science/chem/solutions/fe.html

EDIT: Changed link, so that it works again.

[Edited on 27-9-12 by woelen]

ScienceSquirrel - 7-9-2008 at 08:20

Preparation of Ammonium Iron (III) Sulphate (Ferric Alum)

Experimental

Ferrous sulphate heptahydrate (11.1g, 40mM) was weighed into a 50ml flask. Deionised water (10mL) was added and the mixture was stirred magnetically with gentle warming until the salt had dissolved. Sulphuric acid (90%, 1.8mL, 30mM) was added dropwise and the mixture was placed in a water bath. Hydrogen peroxide (35%, 2.9mL, 30mM) was added dropwise to the stirred mixture. The reaction is vigorous and the reaction mixture goes a muddy brown, clearing to a bright orange as it nears completion.
After the addition the water bath was removed and the mixture was heated for 5 - 10 minutes until effervescence ceased. A solution of ammonium sulphate (2.6g, 20mM) in deionised water (5mL) was added in one portion and the mixture was filtered. The filter paper was washed with deionised water (5mL) and the solution was cooled in an ice salt bath. Pale lilac crystals gradually formed from the deep brown solution. The crystals were filtered off, pressed dry and then allowed to dry in the air. The solid is a pale amethyst in the mass but colourless as small. single crystals. Yield 12.7g, 66%. mp 39-40.5C, Sigma Aldrich catalogue 37-41C.

Discussion

The reaction gives good yields of the alum and could easily be scaled up to produce sufficient material to grow a large single crystal.
It is an advantage to keep the mixture slightly acid to avoid the precipitation of ferric hydroxide.

[Edited on 7-9-2008 by ScienceSquirrel]

kmno4 - 7-9-2008 at 13:35

Yes, I confirm this procedure. FeSO4 + H2SO4 + H2O2 + (NH4 )2SO4 really works fine.
I used to make larger amount ( ~ 200g) of this alun.
Indeed, from deep brown solution small white crystals are formed.
On the picture - mentioned alun.
It is almost coloreless, with a little amethyst (~violet) shade.
ps. my fellow made (in similar way) beautiful, large (> 1 cm), bright violet crystals of Fe(III) sulfate ( x 9 H2O).
ps2. with aid of ~5% acidic solution of K3[Fe(CN)6] you can test presence of Fe(II) in your oxidated solution. When Fe(II) ions are present, deep blue colour/precipitate appears when you mix ferricyanide sol. with tested sol. It is very sensitive test. When only Fe(III) ions are present, mixing gives clean brown solution.

[Edited on 8-9-2008 by kmno4]

jj.PNG - 297kB

blogfast25 - 27-9-2012 at 04:43

I’ve made ferric alum many times, usually with useful product to show for but rarely without problems. Very variable yield, unexpected precipitates, incomplete oxidation, excessive foaming (with HNO3) are all problems I’ve encountered. And once I attempted one madcap scheme to try and synthesize this alum which involved fusing together ferric oxide, ammonium sulphate and conc. sulphuric acid, which actually gave some alum (after leaching) but at low yield. Googling for ‘ferric alum synthesis’ I came across this thread and now finally found the time to test the method cited by ScienceSquirrel, a few posts up. I followed the recipe closely with only a couple of minor modifications.

I applied it to try and make 0.25 mol of the alum (NH4Fe(SO4)2.12H2O, molar mass about 482 g/mol) and firstly combined 69.5 g FeSO4.7H2O, 63 ml of DIW and 11.3 ml of H2SO4 95 % and heated the mixture to about 50 C while stirring vigorously. The sulphate dissolved easily.

Then I put the beaker on ice bath to cool as much as possible but at about 40 C, turbidity (crystallites?) set in, so I started adding the 18 ml of prescribed H2O2 35 % (NOTE: that’s a 45 % excess with respect to stoichiometry). It was ice cold (from the fridge) and added drop by drop (about one drop per second, continuously) from a mounted separation funnel, whilst constantly stirring the reacting mix. Relatively little effervescence was noted and temperature of the liquid only increased by about 10 C, while the mixture darkened to about the colour of not-too-strong coffee. This colour is a good sign because it’s caused by quite concentrated ferric sulphate. No precipitate or turbidity developed.

Effervescence during oxidation seems to point that not only the reduction:

H2O2 + 2 H+ + e- → 2 H2O takes place but presumably also:

H2O2 → H2O + ½ O2 and this is basically lost for the desired oxidation reaction:

Fe2+ → Fe2+ + e-

The mixture was then simmered until all effervescence stopped, which only took a few minutes.

Then 16.3 g of ammonium sulphate was added as a solid (not a solution) and the mix simmered further until the amm. sulphate had dissolved completely (a few minutes). The final volume of solution was about 175 ml.

The coffee brown liquid was cooled to room temperature, a seed crystal was then added and it was refrigerated (about 5 C) overnight. A nice mass of the usual light lilac crystals developed which after recovery, washing with a bit of iced DIW and drying amounted to 108 g of NH4Fe(SO4)2.12H2O or an actual yield of 90 %, the highest I’ve ever achieved and higher than ScienceSquirrel’s yield too. Supernatant liquid solution weighed in at 70 g.

This confirms the validity of the method described by ScienceSquirrel and confirms also what I suspected based on previous experiences: with H2O2 as the oxidant care must be taken that the oxidation of FeSO4 to Fe2(SO4)3 is complete to ensure maximum and consistent yields. Slow addition in relatively cold conditions to avoid breaking down the peroxide to oxygen is essential to achieve this.

[Edited on 27-9-2012 by blogfast25]

Poppy - 27-9-2012 at 11:03

So...
You performed oxydation with H2O2 at low temperatures. This is what I was all worried about.
True is that Fe III is colorless if in the presence of an appropriate non complexating agent.

blogfast,
Your method must have achieved yields over 99% for sure.
The proof can be attained by the undiscriminate addition of hot H2O2 35% to a conc. solution of FeSO4 with a very stabilizing excess of H2SO4 equimolar to that of SO4(2-) already present in the dissolved ferrous salt (that is, enough to react forming Fe2(SO4)3 with a lot of H2SO4 left over to grant proper acidic conditions for Fe3+ to do not hydrolyse).
Add how much H2O2 as you want. A vigorous bubbling is observed. Looking at the solution against light should still show hues of green. Add more H2O2 as you please.
Then, put this to react with slight excess Ca(NO3)2 as to neutralise all sulfate ions in the solution. As reaction slowly takes place (yes, calcium sulfate will not precipitate that fast here).
Fe2(SO4)3 + H2SO4 + 4 Ca(NO3)2 --> Fe(NO3)3 + 3 CaSO4
After precipitation of calcium sulphate, the supernatant liquid should be greenish: the proof that hot addition of H2O2 does not fully oxydise Fe2+ to Fe3+.
All the Fe3+ resulting from the double exchange reaction is now colorless, but Fe2+ is still green. I think its a fairly sensitive test for a pair of good looking eyes.

Another problem between is the stability of Fe3+ in solution I believe it can oxydise water and go back to +2 state. I would point out the best way to prepare Fe3+ is to carry with blogfast's method with very little or just NO excess sulfuric acid, such as upon formation Fe3+ will settle as precipitate, the reason, IMHO is that solid Fe3+ is not as reactive as the aqueous counterpart.

blogfast,
Do you have means to perform this test?

Edit:
Nvm, for low concentration a "Fe2+ meter" would be best employed! It consists of an elongated acrylic parallelepiped which serves as to provide a thicker section of the liquid to be tested. The longer it better. Use a white LED to shiny upon the solution, or any other pure white light source which suits best for you, at one end, and at the other a piece of very bright white board.
I would also suggest making standards with known drops of FeSO4 in H2SO4 solution.

[Edited on 9-27-2012 by Poppy]

blogfast25 - 28-9-2012 at 05:32

Poppy:

A few points.

Hot oxidation of Fe2+ with H2O2 is inefficient, as previous runs showed (low yield, I even got some Mohr’s Salt to sit on top of the alum in one case). Also, a yellow precipitate (an insoluble basic ferric sulphate, I suspect) has been an unwelcome visitor on occasion.

The best and easiest test for Fe2+ in these conditions is K3Fe(CN)6 solution: added to a mix of Fe2+/Fe3+ it forms deep blue insoluble Prussian Blue (aka Berlin Blue, Turnbull’s Blue), due to the Fe2+ present (with only Fe3+ present the solution only darkens somewhat).

I doubt very much that Fe3+ is capable of oxidising water in any conditions: it’s in fact a very weak oxidiser. I’ve used ferric alum as a titrant for Ti(III) determinations but then Ti(III) is extremely easy to oxidise. But if you have literature (or your own empirical) evidence to that effect, I’d be very interested to see it.

White Yeti - 28-9-2012 at 12:01

Double salts are a pain to make. I recently attempted the synthesis of ferric ammonium citrate and it didn't quite work out. I should have gotten a yellow-brown powder upon evaporation but I got a substance with the consistency and colour of molasses even though I followed the instructions to the letter. I guess the synthesis of these salts is never an exact science.

blogfast25 - 28-9-2012 at 12:09

Quote: Originally posted by White Yeti

Double salts are a pain to make. I recently attempted the synthesis of ferric ammonium citrate and it didn't quite work out. I should have gotten a yellow-brown powder upon evaporation but I got a substance with the consistency and colour of molasses even though I followed the instructions to the letter. I guess the synthesis of these salts is never an exact science.

Ferric ammonium citrate is arguably far harder to synth than ferric ammonium alum. FAC doesn't even have a fixed composition. Many double salts I've made presented no problem at all: potassium alum, ammonium alum, Mohr's Salt and others are really straightforward...

[Edited on 28-9-2012 by blogfast25]

ScienceSquirrel - 5-11-2012 at 09:51

Mohr's Salt is ridiculously easy. Garden centre grade ferrous sulphate and ammonium sulphate, a small amount of drain cleaner sulphuric acid in hot water. Cool and watch it snow out in 99+% purity.
It is easily dried, very stable.

There is a preparation of Mohr's Salt here;

<del>http://www.chem.mun.ca/courseinfo/c2210/lab/Experiment%204%20Ammonium%20Iron%28II%29%20Sulfate%20Hexahydrate%20W2011.pdf</del>
Attachment: W2013 Experiment 4.pdf (74kB)
This file has been downloaded 9518 times

Most of the alums are very easy as well.
Potassium alum can be made as white crystals from used cans or aluminium foil, potassium hydroxide and sulphuric acid. Loads of recipes on the net.
Chemistry does not get much easier than this

[Edited on 5-11-2012 by ScienceSquirrel]

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: replaced broken link with attachment]

[Edited on 16.10.13 by bfesser]

Poppy - 5-11-2012 at 12:50

Yea it works, I just can't understand why they proposed an erlenmeyer flask. Because of that, crystals are going to be creep to remove. I had to use a drill bit last time, otherwise I would miss the glassware. That gotta be ajoke..

[Edited on 11-5-2012 by Poppy]

blogfast25 - 12-10-2013 at 09:14

I’ve tried yet another route to ferric ammonium alum, this one to try and further reduce the effervescence (O2 evolution) observed during the oxidation step with H2O2, and seemingly with success.

70 g FeSO4.7H2O (0.25 mole) were dissolved in 63 g of water in a 500 ml beaker, by stirring and heating on an electrical plate. 15 ml of 33 w% (d = 0.88) NH3 solution (that is approx. 0.25 mole NH3) was then added, causing a dark blue black precipitate (a basic ferrous sulphate I think) to form. The slurry was fairly viscous but still easily stirred and was cooled to about RT on an ice bath. It did not smell of ammonia at all.

As in the previous run, 18 ml of ice cold 35 % H2O2 was placed in dropper funnel and added drop by drop (about 1 drop / s) to the slurry, still on ice bath and while stirring intensely (but manually). The oxidation took place immediately and with very little effervescence with the blackish slurry changing colour to reddish/brown. After a while, with some of the peroxide left in the dropper funnel, I concluded the oxidation was complete. 6.5 ml of 35 % H2O2 was left in the dropper funnel.

To this slurry (Fe(OH)3.nH2O + ammonium sulphate) was then added 15 ml H2SO4 (96 %), a slight excess. Most of the ferric hydroxide dissolved immediately, and after about 5 minutes of simmering the solution was completely clear (but very dark in colour). The final volume was just over 100 ml. The hot solution was transferred to a PP beaker for cooling and refrigeration. Yield and colour of the product will be determined tomorrow. Theoretical yield should be 120.5 g of NH4Fe(SO4)2.12H2O.

Another variant could be to avoid the peroxide altogether and use air oxygen as oxidant but that would require the right conditions of blowing air through the slurry, heating and time.

[Edited on 12-10-2013 by blogfast25]

MrHomeScientist - 15-10-2013 at 20:20

I actually have some ferric ammonium sulfate at work, purchased from sigma something like 15 years ago. It came as beautiful lavender chunks. I wanted to try recrystallizing it to shoot for a big, transparent single crystal, but it didn't behave like I thought it would.

According to the wiki, its solubility is 1240 g/L. Accordingly, I tried dissolving 62g in 50mL of room temperature water. This is neat to watch as the purple solid dissolves into a very dark brown solution. After stirring overnight, I was only able to actually dissolve about 39g(!).
I took the remaining 23g and was able to dissolve that in only 20mL of hot water.
I left both solutions out to cool/evaporate and hopefully precipitate crystals, but after about 4 days there is zero crystallization.

It seems like this compound is indeed very soluble in water, but needs a bit of heat before it will dissolve. Once in solution, it seems reluctant to come back out.

blogfast25 - 16-10-2013 at 04:39

Mr HS:

Solubility of this alum is extremely high. So high that at about 40 C it melts in its own crystal water! I recrystallize this material by taking about 2.5 part of the alum, then add 1 part of water, then heat gently until everything is dissolved, there's no need to boil. On cooling overnight the crystals then appear as one large mass. The double pyramid structure can clearly be seen in the larger crystals.

How one would go about making these massive mono-crystals, I'm not sure though. I think without knowledge of the entire solubility-temperature curve it would be hard to do. They are in any case grown from multiple, repeated deposits onto a seed crystal.

[Edited on 16-10-2013 by blogfast25]

woelen - 16-10-2013 at 04:44

Are your solutions perfectly clear? Ferric ion is very prone to hydrolysis and in neutral solution, the free aquated Fe(3+) ion cannot exist. Aqueous Fe(3+) is nearly colorless (and as you know, solid ferric ammonium sulfate or solid ferric nitrate has a very pale color), the brown color of solutions of iron(III) are due to hydrolysis:

Simplified: Fe(3+) + H2O --> [Fe(OH)](2+) + H(+)

If you add a few drops of dilute sulphuric acid, then the liquid should become much lighter and then you should be capable of producing nice crystals. If you do not have a slight amount of acid in the solution, then you may get a brown contamination of your solid, these are small hydroxide particles encapsulated in the crystalline solid.

Do not add hydrochloric acid, because that forms the intensely yellow colored complex FeCl4(-).

blogfast25 - 16-10-2013 at 08:45

Quote: Originally posted by woelen

Are your solutions perfectly clear?

Yes. During synth I always have a slight excess H2SO4 present..

But for recrystallization purposes, pure water is good enough. I think the acidity of the ammonium ions is enough to prevent ferric hydroxide to form.

The other remaining mystery with ferric alum is the colour. I've obtained both lavender and off white versions in seemingly identical conditions. Wiki shows the off white version. 'kmno4' seems also to have obtained an off white product. A commercial sample I bought years ago was lavender. And Mr HS's Sigma Aldrich product was also lavender.

And my latest batch (post) above was also lavender. So what causes this colour difference which is quite striking?

[Edited on 16-10-2013 by blogfast25]

MrHomeScientist - 16-10-2013 at 10:53

Yep my solutions are clear as well, no precipitate at all. They are just very dark reddish brown like picture 3 in the OP. They come up as pH of 1 with pH paper, with no addition of acid. My initial crystals are very lavender, very much like cloudy amethyst.

[Edited on 10-16-2013 by MrHomeScientist]

ScienceSquirrel - 17-10-2013 at 13:41

The account of the reaction is not quoted but my own.
The experiment is written up in my best JACS's style :-)
The solution is a very dark brown but it is not cloudy.
The crystals are a light lavender and they have a melting point within the range reported by Sigma Aldrich et al.
I reckon the product is pretty pure within the limits of home science.

blogfast25 - 18-10-2013 at 04:36

Quote: Originally posted by ScienceSquirrel

The account of the reaction is not quoted but my own.

Not sure what you mean by this. Higher up I gave you full credit for the acid/peroxide recipe.

The yield with the ammonia/peroxide recipe was just over 70 %. Not spectacular either.

[Edited on 19-10-2013 by blogfast25]

ScienceSquirrel - 18-10-2013 at 07:43

Sorry, my mistake.

Poppy - 22-10-2013 at 15:58

A very large (1 inch) single crystal of the alum was prepared by adding aluminium foil to ferrous ammonium sulphate with a little sulfuric until the aluminium foil shifted into solution finally forming a massive sample of the so told Fe (III) alum. The solution color was not that red but rather supposedly a 50/50 mixture of Fe3+and Fe2+ at best. I did not repeat this, but the foil was rolled as to fit about 7 - 10 full spires in a 250mL beaker.

blogfast25 - 23-10-2013 at 05:08

Poppy: was it an actual monocrystal or just one large mass of crystals, clumped together?

Poppy - 25-10-2013 at 16:06

It was really a single crystal, a picture is located somewhere along the subject of iron III!
I just cant believe that, the conditions were harsh, and I tought there had been put a stone in the flask at first sight, then woohoops! A big purple single crystal!
Careful control of this might yield the fastest greatest crystal one can control AT RT!!!!

MrHomeScientist - 12-12-2013 at 15:12

Got some photos of my crystals for your viewing pleasure!

My ferric ammonium sulfate was purchased from Sigma Aldrich something like 15 years ago. It appears in the bottle as lavender chunks. (Photo 1)

I dissolved some in water as detailed in my post above back in October and left it in the fume hood to see what it would do. Just last week I checked on it and gorgeous large crystals had grown! I harvested the largest from the liquor and some were very well formed. The largest was unfortunately plagued with lots of parasitic small crystals. (Photo 2)

I placed three of the better formed crystals back in the solution to continue growth. You can see the rich brown color of the solution, completely clear with no iron hydrolysis evident. We'll see how it goes! (Photo 3)

blogfast25 - 13-12-2013 at 11:05

Very nice indeed, MrHS.

Years ago I read that growers of really macro crystals (of alums), like 2" or more, actually cheat by gently filing (or sanding) off the parasitic ones! Sound plausible to me: once a crystal fault line has developed there's no cure, well... other than filing off the side growth.

Re. hydrolysis, the brown colour is due to Fe(OH)2+, which would be the first step to Fe(OH)3. But here acidity and concentration of Fe(III) prevent it from going all the way.

[Edited on 13-12-2013 by blogfast25]

DraconicAcid - 13-12-2013 at 11:42

Impressive. I can never get nice crystals like that, even from something as simple as alum.

MrHomeScientist - 16-12-2013 at 07:28

Thanks! I had actually forgotten about them for a while, growing in the back of the hood. Apparently it's the perfect environment for slow crystal growth!

Thinking about it this weekend, I had an idea: What about growing my large crystal on a teflon stage? I.e. put a disk of teflon in the beaker that the large crystal will sit on. Would teflon resist formation of new crystals?

If not teflon, what about a hydrophobic surface? I have some Never Wet that I've been playing with. If the solution can't get to the surface, maybe that would prevent new crystals from growing.

Any other ideas to prevent these little crystals from growing where I don't want them?

Poppy - 16-12-2013 at 11:13

Mr Home Scientist, I think those big crystals are nothing but a chunk of ammonium sulphate that has went out separately!
You can compare sizes and color or even redissolve those if you please and tell us if anything might have ended bad again!

blogfast25 - 16-12-2013 at 14:42

Quote: Originally posted by Poppy

Mr Home Scientist, I think those big crystals are nothing but a chunk of ammonium sulphate that has went out separately!
You can compare sizes and color or even redissolve those if you please and tell us if anything might have ended bad again!

Utter nonsense. If you start from ferric ammonium alum you end up with ferric ammonium alum.

blogfast25 - 16-12-2013 at 14:47

Quote: Originally posted by MrHomeScientist

Thinking about it this weekend, I had an idea: What about growing my large crystal on a teflon stage? I.e. put a disk of teflon in the beaker that the large crystal will sit on. Would teflon resist formation of new crystals?

If not teflon, what about a hydrophobic surface? I have some Never Wet that I've been playing with. If the solution can't get to the surface, maybe that would prevent new crystals from growing.

I don't really see how sitting them on a hydrophobic surface would prevent solution from getting to the [crystal] surface. Even if you did achieve the latter, you wouldn't get any crystal growth at all there.

MrHomeScientist - 17-12-2013 at 11:43

Quote: Originally posted by blogfast25

Well that's the idea though. Place a hydrophobic or teflon surface in a beaker, place one of my large crystals on that, and fill the beaker with ferric alum solution. This way the large crystal will grow, but no small crystals will start growing elsewhere on the surface since there'd be no other nucleation sites. That's the thought, at least.

blogfast25 - 17-12-2013 at 12:40

Hydrophobic really means 'high surface tension' but even the highest surface tension is overcome by modest hydrostatic pressure, for instance a Teflon beaker still fills with water.

I personally don't see how a hydrophobic resting surface for the crystals could prevent side growths.

How about weightlessness? All these balls of ferric ammonium sulphate solution with crystals in them, floating about in a space station? The sky is the limit, no?

[Edited on 17-12-2013 by blogfast25]

MrHomeScientist - 17-12-2013 at 14:22

I see, thanks for the responses.

If anyone knows of any methods to prevent these extra crystals from growing I'd be interested to hear about them. Space station would be ideal!

Barring that, I know some people suspend the crystal in the solution to avoid parasitic growths. But I can't think of how to do that without ruining the formation process around whatever string holds it up, and then trapping the string inside the crystal for all to see. I'd think even fishing line would be visible, unless it happens to match the material's index of refraction perfectly.

Poppy - 19-12-2013 at 21:16

Quote: Originally posted by blogfast25

Quote: Originally posted by Poppy

Utter nonsense. If you start from ferric ammonium alum you end up with ferric ammonium alum.

I just don't know but my method yielded much purplish samples, so how about no?

blogfast25 - 20-12-2013 at 05:41

Poppy:

We know ferric ammonium alum comes in two colour variants but we don't really know why (it's been discussed elsewhere here). The colourless crystals in MrHS sample are FAA too, not ammonium sulphate.

[Edited on 20-12-2013 by blogfast25]

MrHomeScientist - 20-12-2013 at 12:09

Update

I took my crystals out of the growing solution Friday of last week, since I wouldn't be around to check on them over the weekend. I returned Wednesday and found they had discolored - turning significantly brown!

I put the nicer, large crystals back into solution and covered the beakers with Parafilm. I didn't want any more growth, I just wanted to see if "rehydrating" them would solve the problem. To my relief, it did. Two days later, my large crystals have returned to their previous transparent lavender color. Had me worried for a minute there! I wasn't aware these crystals would be air-sensitive.

blogfast25 - 21-12-2013 at 06:19

Years ago I saw a sample of a large FAA crystal (about 2"), in a museum. Years of exposure to dry air had eroded the surface considerably, even though it had been varnished to prevent that!

Just keep your samples in a closed container and they should preserve indefinitely.

rey31tjia - 31-10-2017 at 02:53

Quote: Originally posted by kmno4

Yes, I confirm this procedure. FeSO4 + H2SO4 + H2O2 + (NH4 )2SO4 really works fine.
I used to make larger amount ( ~ 200g) of this alun.
Indeed, from deep brown solution small white crystals are formed.
On the picture - mentioned alun.
It is almost coloreless, with a little amethyst (~violet) shade.
ps. my fellow made (in similar way) beautiful, large (> 1 cm), bright violet crystals of Fe(III) sulfate ( x 9 H2O).
ps2. with aid of ~5% acidic solution of K3[Fe(CN)6] you can test presence of Fe(II) in your oxidated solution. When Fe(II) ions are present, deep blue colour/precipitate appears when you mix ferricyanide sol. with tested sol. It is very sensitive test. When only Fe(III) ions are present, mixing gives clean brown solution.

[Edited on 8-9-2008 by kmno4]

How did you crystallize it? Did you refrigerate the solution? Please help me..
Thank you..

jamit - 28-7-2023 at 17:43

so I followed the synthesis of blogfast25 and I obtain violet crystals but they oxidize so quickly to light brown with in 10 hrs and they are sticky? did i do something wrong? has anyone also had this problem making ferric alum.

I dissolved ferrous sulfate and acidified it with sulfuric acid. Then I oxidized it with peroxide until the effervescence ceased. I then added solution of ammonium sulfate and mixed the two solutions together.

Then I filtered and heating to remove some liquid and allowed it to evaporate until lilac crystals form. but the ferric alum oxidizes back to light brown color within 6-10 hrs while drying and they are really sticky. help!!

Fery - 29-7-2023 at 10:28

jamit you did very nice experiment
I did the same as you did
I remember that I used slight excess of H2SO4
I think the color change into orange it is not any oxidation, but it is losing of crystalline water or absorbing air moisture
pure crystals are almost colorless, very very very pale violet, I have a vial with them which I synthesized 30 years ago and they are still good quality (but the cap is also sealed with wide rubber band)
water solution and crystals with lost crystalline water have orange color
very interesting how almost colorless crystals are formed from such orange colored solution

Bedlasky - 29-7-2023 at 12:01

Jamit: The colour change is due to loss of water, not due to oxidation. Ammonium ferric sulfate is violet because it contains violet [Fe(H2O)6]3+ ions in it's crystalline structure. If some water is lost, sulfate coordinates on to Fe(III) an these complexes are brown in colour.

To prevent water loss, you can store crystal in saturated solution of ammonium ferric sulfate or in the paraffin oil. I keep mine in paraffin oil and it is still violet after several years. It doesn't look exactly like freshly crystallized crystal, it's paler, but it is still well preserved.

For crystal growing tips, this site is quite useful.

https://en.crystalls.info/Iron(III)-ammonium_sulfate