Sciencemadness Discussion Board - FeBr3 from Fe

FeBr3 from Fe

Klute - 19-10-2007 at 11:28

I'm planning on trying secondary and benzylic alchols oxydations with H2O2 catalyzed by FeBr3 ( S.E.Martin, A.Garrone; Tet.Lett 44(2003) 549-552 ), and I need to make the FeBr3 myself. The FeBr3 doesn't need to be anhydrous, and I guess the reaction can tolerate impuities such as FeBr2 (which will be oxidized, but might act as FeSO4 and generate some free radicals), aswell as neutral to acid salts in small quantities. To obtain FeBr3, I'm thinking of two ways:

Either making a HBr solution from KBr and H2SO4, and slowly add enough H2O2 to form Br2. To this is then directly added Fe wool in exces, according to:

2Fe + 3Br2 --> 2FeBr3

Then the suspension is filtered from excess Fe and eventually solid impurities, and evaporated. The FeBr3 is then recrystallized, from water I guess, as nothing is mentionned in Purification of Lab. Chem. 4Th, only FeBr2. I suppose FeBr3 is simuilar to FeCl3. Impurities should be KHSO4, K2SO4, maybe FeBr2, and crap from the H2SO4. No need to handle liquid bromine, would like to avoid that.

Or, I was thinking of reacting Fe with an acid, isolating the corresponding Fe2+ salt, oxydizing it with H2O2 to convert it to Fe3+ salt, then adding NaOH to precipitate Fe(OH)3, filtering, washing, and dissolving in minimum HBr (aq), followed by evaporating and crashing out with alcohol.
I also could just buy some FeCl3 solution from the lectronics store, neutralize with NaOH to obtain the hydroxyde, and continue on from there.
The salt made this way could be a bit more pure, but it's more work.

The problem with the first reaction is that I've only seen this as written equations, which means nothing on the usefullness of the reaction. It could take days.
I'm supposing the Fe(OH)3 can easily dissolve in HBr (aq), but then again, I've got nothing backing this up apart from personal convinction.

So I would like to know your opinions on this, and if anyones tried to obtain this salt by such a way, or has some information on it.

I will post my results with the oxydations when done, aswell as the ref. (I think most of you see which paper I'm reffering to).
The KBr/TCCA in DCM (http://www.sciencemadness.org/talk/viewthread.php?action=att...) will also be tried, if amount of solvant can easily be reduced, and the yields of both oxydations compared.

Seconday and benzylic alcohols to corresponding carbonyls is a very important reaction in organic chemistry, so tried procedures and experimental data would be very usefull here, IMHO. Alot as been discussed, but little has been tried, to my knowledge.

woelen - 19-10-2007 at 12:03

The easiest way is to proceed as follows:

Dissolve KBr in an as little as possible amount of water. Take a stoichiometric amount of concentrated sulphuric acid (2 KBr for 1 H2SO4) and dilute this with 50% of its volume with water (careful: lots of heat are evolved).

Heat the solution of KBr and slowly add the hot mix of water and H2SO4. This will result in a very hot solution of KBr/H2SO4. The solution also will turn somewhat yellow/orange, but that is not a real concern at this stage. By pre-diluting the H2SO4 somewhat this effect should be minimal, however. Mix the hot solution carefully. Do not add 96% concentrated acid to the solution of KBr, that will result in excessive oxidation of bromide to bromine, with loss of SO2 from the solution.

On cooling down, lots of glittering crystals are formed. These crystals mainly are K2SO4, with some KBr and KHSO4 as contamination. Allow the liquid to cool down, and then put it in a freezer to get even more of these crystals.
The liquid then is decanted from the crystals. This gives fairly pure HBr in water, with only a minimal amount of sulfate and potassium ions. Now you have much less crap in your aqueous HBr than with your method of making aqueous HBr.

I have done this prep of raw HBr myself and I was quite content with the result. For most of my experiments the remains of potassium and sulfate were not a problem at all.

If you want really pure HBr, then you need to distill the liquid. HBr and water go over, potassium ions and sulfate ions remain behind.

In this solution, you dissolve your Fe(OH)3. Freshly precipitated Fe(OH)3 dissolves quite well, but if it has aged, then it only dissolves very slowly. Keep in mind that FeBr3.xH2O is very hygroscopic and it will be hard to get the solid. FeCl3.6H2O also is very hygroscopic, it is deliquescent. Heating a solution of FeBr3.xH2O will result in loss of HBr and a basic ferric bromide remains behind. Probably you'll have to use an organic solvent which dissolves in water in order to precipitate the ferric bromide.

[Edited on 19-10-07 by woelen]

Klute - 19-10-2007 at 12:52

Thanks for the tip with the aq.HBr solution, that's clever. I could always distill the HBr (aq) at over 100°c, but I think it's not necessary.

As the FeBr3 will be used in aq. solutions anyway, I won't even bother crysatllizing it: I'll calculate the amount used, and estimate the concentration of my fianlly solution and use it as is.

I guess I'll get some aq. FeCl3, it will be more practical and clean. Etching FeCl3 solutions are of reasonable quality.

What do think about the Br2 + Fe reactions? Would that proceed reasonnably quickly at room temp (I bet it's an exothermic reaction though)

Eclectic - 19-10-2007 at 13:21

Wet scrap iron is (or was) used as a scrubber/concentrator to trap Br in several brine and seawater bromine extraction processes as of 1960 or so.

Klute - 19-10-2007 at 14:44

So I guess the reaction must be pretty quick then...

S.C. Wack - 19-10-2007 at 14:51

I think there has to be a little water and excess bromine present.

12AX7 - 19-10-2007 at 14:56

Isn't it the case that FeBr3, like FeCl3 and AlCl3, is a lewis acid and as such must be anhydrous to function?

Tim

Klute - 19-10-2007 at 16:15

No, it's used with 30% H2O2, surely forming hypervalent Fe species.

12AX7 - 19-10-2007 at 17:11

Ok, then in solution, how much does the Br- matter? Why not Cl, I, or even H2O? (I assume we're talking ligands now, so H2O would represent solutions of Fe2(SO4)3, Fe(NO3)3, etc.)

But then, it occurs to me that we're talking chemistry. It always makes me chuckle to hear a foolhardy chemist generalize about reactions, only to end up covering the remaining elements with 87 exceptions in the theory.

Tim

Klute - 20-10-2007 at 03:21

Actually, in the paper, they tried the same reaction with FeCl3, with hardly any results. So the Br- seams to be absolutly necessary. Don't ask me why though!
Also, KBr alone didn't produce any reaction. One thing they could have tried, would have been using different Fe3+ salts WITH Kbr in solution, to see if the presence of Br- is usefull, or if it really has to be the counter ion. I will try that perhaps.

But we are slowly drifting off topic, I will post my results of the oxidation in the org. Chem. forum, and the FeCl3 --> FeBr3 here.
I've just got some FeCl3 granules.

Klute - 21-10-2007 at 10:22

Following Woelen's advice, 10.7g (90mmol) of KBr were dissolved in 20mL of dH2O with slight heating. To this solution, 11.91g of 37% H2SO4 (45mmol) were slowly dripped in over 5 min, with slow stirring. The water clear solution turned to a pale green/yellow color. The erlenmeyer was cooled at room temp, then in a fridge, then at -20°c, where nice cristals formed.

During that time, 8.1 g (30 mmol) of yellow FeCl3.6H2O pellets were dissolved in 30mL of warm water, giving a orange-brown solution. 3.6g (30 mmol) of NaOH were dissolved in 30mL of dH2O and cooled.
The NaOH solution was then added dropwise to the FeCl3 solution with good stirring, over 10min. A dark brown precipitate formed at each drop.
Near then end, the suspension turned to a thick paste of granular precipitate, but the magnetic stirring turned the Fe(OH)3 into a very fine solid after a few minutes. The pH was stronglty basic to universal pH paper.
The oxyde was left to decant, and the surnatant removed with a pipette. 75ml of dH2O was added and everything stirred, and left to decant. The oxyde was washed 3 times this way (The last wash was still basic), and then vacuum filtered. This took a long time as the very fine oxyde plugged the filter paper immediatly. After 15min, a thick brown paste was remaining.

The greenish HBr solution was decanted from the formed crystals while still cold, then warmed to 40-50°c. With steady stirring, the Fe(OH)3 paste was added to the soltuion. White fumes appeared as soon as the paste was above the solution. The solution quickly took a reddish tint, and the oxyde dissolved slowly. After 15min, a yellow/white fine precipitate was left, with a orange-redish supernatant liquid. The salt must have been NaBr formed from excess NaOH. This was vacuum filtered, which took time, to afford 30mL of a red solution., containing roughly 25 mmol of FeBr3 with some salts.

Half of the solution was used as is for the oxydation (See the thread Organic Chemistry forum).

If I wanted purer FeBr3, I would have distilled the HBr, further washed the oxyde, and cristallized the FeBr3 by evaporating the water and adding some solvant to crash the salt.
Rather than making small quantities of this soltuion, I think I will take a weekend to make a largish amount, as everything is cheap. It just takes time.

woelen - 21-10-2007 at 11:45

Klute, next time try to use a more concentrated solution of H2SO4. I did this with H2SO4 of roughly 75% (by weight), made by mixing 2 volumes of concentrated acid with one volume of distilled water. Using these higher concentrations, you will have even more crystals of K2SO4 separating from the liquid, and you'll have a purer (and more concentrated) solution of HBr. You also will have a little bit more Br2, but still, that will be low-level.

Klute - 22-10-2007 at 06:43

I decided on using the 37% just because it was closer at hand and I was too lazy to dilute some 96%.

But I will do next time. Also, i will vacuum filter the crystals as alot of liq. HBr stays trapped in the crystals/sludge when decanting. Once it heats up again, it makes a thickish green syrup.
Does your HBr aq always have that beautifull green/yellow color? I don't remember reagent grade 48% HBr having that color, but it could just be the purity, or I dodn't remember well.
In any case, thanks for the tip.

BTW, it seems that using Br2 and Fe filings to make FeBr3 quantitaively isn't feasible, as a protective layer of FeBr3 forms over the the Fe and prevents further reaction, even with strong stirring and different rough conditions.... Thanks to Java and Lugh for this.

[Edited on 29-10-2007 by Klute]

Monoamine - 21-7-2021 at 13:33

Once you have some Br₂, couldn't you just dry it by washing it with concentrated sulfuric acid?

This is basically how chlorine gas is dried when making FeCl₃, by passing it through concentrated sulfuric acid.

So once you have dry Br₂, i'd imagine you can just drip it onto the iron (very slowly!) until it all reacted. And of course you want to make sure your apparatus is completely moisture free (for instance by adding some drying tubes containing a dessicant at the top of any outlet.

Monoamine - 21-7-2021 at 13:42

Quote: Originally posted by Klute

I decided on using the 37% just because it was closer at hand and I was too lazy to dilute some 96%.

How about this then: Take an flat iron tile and cover it with Br₂. Once the Br₂ has all run off and the plate is dry, take a steel brush and just scrape the off the FeBr₃ layer and collect it. Continue covering the iron plate with bromine and scraping off the top layer until you collected enough.

woelen - 22-7-2021 at 00:43

I do not expect any decent yield. Nearly all Br2 will evaporate and you just fume yourself! You also won't get pure FeBr3, but some mixed iron(II)/iron(III) compound of approximate composition Fe3Br8. This compound is very hygroscopic and you certainly won't be capable to scrape this off with a brush to get a nice dry powder.

No, your method is not useful for making FeBr3.

unionised - 22-7-2021 at 04:03

Well, the OP has had 13 years or so to get it to work.