Sciencemadness Discussion Board - ammonium perrhenate reduction

ammonium perrhenate reduction

Jor - 28-3-2010 at 06:34

Having bought 75 grams of ammonium perrhenate for about 45 euros (and having 5 more grams as a donation from woelen before this auction was on eBay

), I want to at least reduce half of this to high purity Re powder, as i want to make Re compounds from this.

I have seen that the best method of reduction is with hydrogen at very high temperatures, I think above 500C (and thus autoignition temperature). I think this is too much risk and I don't want to do this.

Having done many searching in google, the sciencemadness library and the internet library of the university, i cannot find much information, except reduction with borohydride, wich i do not have.
Does anyone have experience with the reduction of this salt in other ways? If going through aqeous routes i prefer a heavy sponge and not something such as palladium black wich is very fine and hard to filter.

not_important - 28-3-2010 at 07:09

This is a tough one. Many common reducing agents only take it to the +4 state. Hydrazine or Na-Hg on prolonged contact with perrhenate solutions give a mix of the metal and dioxide.

Unfortunately the books that I know of that really cover Re chemistry are rare and in copyright...

Fleaker - 28-3-2010 at 08:40

Congratulations on your recent acquisition--interesting enough that you get some so soon :-).

You have to reduce it with hydrogen--it's the best and only way to go. Do the reduction at 700 C in hydrogen atmosphere in a quartz vessel; purge the tube with hydrogen before it is at temperature. The reduction is smooth, and you can follow its progress by titration of the ammonia you will collect.

Cool under hydrogen, then turn off tank and pull a vacuum and heat to 500C to get the majority of the hydrogen off the sponge. If you do not do this, it will take fire when you pour it from the tube. Your quartz tube will be plated mirror-bright with rhenium (it always is).

I'm sure you can make quite a few rhenium compounds from perrhenate itself...

JohnWW - 28-3-2010 at 10:20

It is very difficult to reduce Re from the (VII) oxidation state, unlike Mn(VII). It was once thought that the polarimetric electrolytic reduction of the ReO4- anion, by 8 redox equivalents, reduced it to the Re- anion; but subsequent work, particularly including that by Prof. Neil F Curtis, who was the Prof. of Inorganic Chemistry at Victoria University Of Wellington, New Zealand, when I was there, showed that it was really hydrogenated to the ReH8- anion. (Or it could have been instead, or some of it was additionally, ReH9--, as DJF90 says below, but a coordination number of 9 is extremely rare, even with the small size of H atoms.).

Re is used largely as a catalyst (experimental). Because of its very high melting-point, it has been used as a substitute for W in the filaments of incandescent electric light bulbs, and as a carbide it could probably substitute for WC in tips of cutting blades and drill bits; but it is really too rare for such uses.

[Edited on 29-3-10 by JohnWW]

woelen - 28-3-2010 at 11:00

I agree with JohnWW that it is not easily reduced. Perrhenate is not a strong oxidizer at all. Using sodium hypophosphite combined with some acid and heating, however, I managed to reduce the colorless solution to a dark (almost black) turbid liquid. I'm not sure though if this contains rhenium or some oxide at lower than +7 oxidation state.

Jor - 28-3-2010 at 11:04

not_important , I do have hydrazine. Will heating it with excess hydrazine for a few days maybe reduce it all to the metal? And could i wash away the dioxide maybe with HCl?

Fleaker, thanks for that method, but I do not have any quartz tube at hand, and I think the process is too dangerous.
I think I will try some reducing agents.
-hydrazine
-hypophosphite
-zinc powder, possibly heating large excess of zinc powder with APR in a crucible, and washing out the zinc with HCl or NaOH? I'm not sure if this reaction is energetic.
-If it is soluble in ethanol, maybe Na/EtOH
-Al/Hg, then remove excess Al with NaOH, and boil away traces of mercury (milligrams at most)

But I am not really confident that any of these works, as according to one article even Li or Na in liq. NH3 gives ReO.2H2O. I do think that a dry mixture of Mg and APR works, but I don't have Mg powder, and that reaction could proceed to energetically, with the result that the Re flies away as smoke.

A ReH8- ion, that is quite interesting!

DJF90 - 28-3-2010 at 11:07

I thought it was ReH9(2-) with a tricapped trigonal prismatic structure if I remember correctly...

[Edited on 28-3-2010 by DJF90]

not_important - 28-3-2010 at 21:22

I don't think zinc takes it any lower than Re(IV), indeed if Na-Hg doesn't fully reduce it I'd be a bit surprised if Zn did.

I thought of hypophosphite, but couldn't find any mention of a state below IV for it either.

There are mentions of electrolysis in fused salts to obtain the bulk metal, although that sounds no more fun than the H2 reduction.

Mentions of perrhenate oxidising HBr to produce Br2, again implied you end up with the IV state.

hmmm...

Quote:

the reduction of potassium perrhenate by potassium metal yields the nonahydridorhenate K2ReH9, containing the ReH92− anion in which the oxidation state of rhenium is actually +7

doesn't sound too encouraging for lower temperature chemistry to do the job. However it appears that most Re compounds don't start with the metal, but the +4 or +7 state so perhaps the metal is not what you want.

Jor - 29-3-2010 at 04:32

Well I at least wanted some Re metal as an element sample, so about 0,5g powder. Maybe I can do the reduction with hydrogen on a very small scale safely, although even small amounts of the gas can cause dangerous explosions, wich is highly dangerous in glass. I don't have a quartz tube. The only device I have to lead a gas in and out is a 3-neck RB flask. I could place some of the APR in that and lead a stream of H2 through, but can RB-flasks be used at temperatures of 700C?

I thought that Zn might work when you melt the zinc together with the APR, because of the high temperatures involved. I will try this on a 100mg scale when I get home, so 100mg of APR and 500mg Zn (large excess) and melt in a crucible with cover. If H2 works at 500-700C, I don't see why zinc won't at very high temperatures. And it is safer than H2 gas at autoignition temps I think.
The only problem is the H2 absorbed on the Re. I don't have vacuum, can it removed in other ways, such as simple heating to 500C in a current of CO2? Or adding some H2O to the cold sample, and then treating this suspension with something with oxidises the H2 but not the Re?
To be honest I don't know how reactive Re is. I know it reacts with HNO3, but does it react (slowly) with HCl? Because it seems it is very hard to get to the metal, wich would indicate that this metal is not very noble at all (although precious

)

If i want to make a compound of Re, I think I will first reduce and precitate an oxide, and dissolve in HCl to make a hydrated chloride.

[Edited on 29-3-2010 by Jor]

not_important - 29-3-2010 at 05:21

I think that the high temperature H2 reduction works because the gas flow removes the H2O formed.

Dry NH3 at the same temperature _might_ be an alternative, especially if first flowed over some iron at the same temperature.

You want a tube, and borosilicate glass is only rated to 500 C or so. Possibly you could salvage the fused quartz tubes in an old radiant electric heater. Or use a iron/steel tube with a ceramic/fused quartz boat to hold the perrhenate.

I suspect the H2 dissolves in the metallic Re, so a decent vacuum at high temperature would be needed to pull it back out. A flowing gas stream at somewhat reduced pressure might work, but CO2 would seem likely to serve as an oxidising agent to the metal; N2 should be OK.

Sounds like you need to track down a copy of Rhenium: dvi-manganese, the element of atomic number 75 by Gerald Druce. Also this PDF http://library.lanl.gov/cgi-bin/getfile?rc000051.pdf which states that aqueous HCl and HF do not dissolve the metal, but moist air slowly converts it to HReO4.

It looks as if wet methods of making chlorides often result in mixtures of halide or complex halide and either oxihalides or HReO4 - Re loves oxygen it seems. Also

Quote:

ReOCl4 will react with ammonia to form ReO(NH2)2Cl. This oxydiamidodichloride compound will decompose at temperatures above 400° to form Re and Re02.

woelen - 29-3-2010 at 05:22

Re metal is somewhat air-sensitive. I have 10 grams of the powdered metal and once it was a nice dry powder. Right now, it is a sticky mass which is very humid. Moisture and oxygen from the air have oxididized the metal somewhat in which HReO4 is formed, which is a hygroscopic material and this entirely wets the metal. If I had known this, then I would have stored the metal in a better container (it now is in a little glass vial with a flimsy plastic cap). The powder is a very compact black material, 10 grams is just a little bit more than a spatula full of powder.

I also have the bulk metal (little pieces of foil) and these seem to keep better. Of course, the contact area with air is much lower in this case and hence the much less pronounced severity of oxidation.

[Edited on 29-3-10 by woelen]

Jor - 29-3-2010 at 05:25

Maybe you could wash away those traces of perrhenic acid with water, wash with acetone and dry?

But thus it seems that Re is quite a reactive metal...

Thanks for the article not_important, I will look into it. Also, if NH3 works, that would be very nice, I might try it.

[Edited on 29-3-2010 by Jor]

watson.fawkes - 29-3-2010 at 07:59

Quote: Originally posted by not_important

I think that the high temperature H2 reduction works because the gas flow removes the H2O formed.

This is certain. I've been reading up on the construction of vacuum devices, and the procedure for brazing stainless steel requires a dry hydrogen atmosphere because of the formation of chromium oxides. "Dry" hydrogen here doesn't mean tank hydrogen (or worse, Kipp-generator hydrogen). It means a hydrogen source that's been passed through a drying train consisting of first, a catalytic oxidizer to convert any residual oxygen (much lower than LEL; these are impurities) to water vapor and second, a heated column of molecular sieves to trap all the water vapor. For achieving bulk reduction of rhenium, all this apparatus (likely) won't be necessary, but it does indicate that complete reduction is pretty sensitive to the steady-state concentrations of H2 vs. H2O, and that's for chromium. Rhenium is certainly much more sensitive.

Safety for heated hydrogen atmosphere reactions isn't easy to achieve, but it's pretty straightforward. The apparatus is a continuous flow system, not a sealed system. On the far end of the apparatus, there's an exhaust for the gas (see below). Before applying heat, you purge the chamber with nitrogen. Then, again before heating, you replace the nitrogen with hydrogen. At this point, you bring the furnace up to heat. When your reaction is done, you keep the hydrogen flowing while the furnace cools off. Finally, you purge with nitrogen again before opening the furnace.

The exhaust tube is cooled in a water jacket to reduce the possibility of explosion. The LEL (lower explosive limit) of any fuel gas is lower when the gas is hot. If you've got an adequate ventilation system and can dilute the hydrogen coming out quickly enough (and have no spark sources in the gas flow), it's fine to just exhaust the hydrogen. It's also possible to flare off the hydrogen by using a gas jet with an adequately small orifice and an adequately long and thin feed tube (to cool off any flashback below its propagation point). At the beginning of the hydrogen purge, you use an auxiliary flame to drive combustion. After a while, the hydrogen will burn by itself.

As for using NH3, I would guess that nitriding would be a serious problem.

not_important - 29-3-2010 at 08:39

The references I have at hand indicate that rhenium does not form a nitride, thus my suggested for using ammonia. The sulfides decompose under strong heating, so precipitating a sulfide followed by drying and then strong heating under reduced pressure might be another route.

unionised - 29-3-2010 at 12:12

Just a thought.
Here's a quote from Sidgwick's Chemical elements and their compounds;
"Rhenium compounds are easily reduced to the metal."
If it was easy in 1962 why isn't it easy now?

woelen - 29-3-2010 at 13:17

From a theoretical point of view it is easy, just pass some dry H2 over the red hot metal

But practice is another thing. Theory does not tell anything about the engineering problems, it does not suffer from safety issues (explosion and fire risk). That is a problem which I frequently encounter. Many times, text books mention how simple a certain reaction is, but if you try it in a real system, then a lot of nasty practical engineering problems jump out of the hat and you have a hard time to catch all those beasts

watson.fawkes - 30-3-2010 at 06:22

Quote: Originally posted by not_important

The references I have at hand indicate that rhenium does not form a nitride, thus my suggested for using ammonia.

OK. Practically speaking, the dryness of the ammonia could matter a lot (I don't know with any certainty one way or another), which means an anhydrous ammonia source or a drying train. The trouble with drying in this case is that molecular sieves won't distinguish between NH3 and H2O. Even tank ammonia might have too much water in it. Per Armarego, use a cold trap, chemical drying with CaO, BaO, or Na, and the following gem: "It can be rendered oxygen-free by passage through a soln of potassium in liquid ammonia".

It's not clear to me how much of this is necessary for the reduction, but it seems like more equipment than a hydrogen furnace requires. Then again, I'm pretty comfortable mitigating down fire dangers in ways that others are not.

[Added] It appears that rhenium catalyzes the decomposition of ammonia at a temperature range below that needed to reduce the oxide. See Ammonia Decomposition and Related Phenomena on Rhenium Catalysts. If I'm reading this right, it means that use of ammonia doesn't get rid of the fire hazard, but rather gives a mixed fire + ammonia fume hazard.

[Edited on 30-3-2010 by watson.fawkes]

not_important - 30-3-2010 at 07:12

Even iron will do the same, it's used when a reducing atmosphere is desired and the low volumetric density of H2 is a problem; liquid NH3 (pressure tank) has 50% more hydrogen atoms per volume than does liquid H2.

It does reduce the fire hazard, as until decomposed NH3 has a smaller flammability range than H2, isn't near so eager to leak through tiny holes and small leaks are easy to detect.

And yes it likely takes a good drying, but so would H2. Traces of water shouldn't be a problem, the reduction creates water. The sweeping out by the gas stream handles the water issue.

Neither should need to be completely free of water or oxygen, so long as the O2 level didn't bring you inside the explosive range (a fairly high level for contaminant in a gas generator) You can burn the NH3-H2-N2 mix the same way you burn straight H2, feed into a gas flame.

unionised - 30-3-2010 at 10:00

Why worry about water?
The same reference I quoted says you can reduce just about any rhenium compound with hydrogen, that would include the oxide so the reaction
ReOx +x H2 --> Re + x H2O.

If the reaction goes in that direction the it won't be interested in going in the opposite direction so water won't react with rhenium.

BTW "From a theoretical point of view it is easy, just pass some dry H2 over the red hot metal"
There's a hole in my bucket...

I think you could get some relatively inert gas like nitrogen or argon, bubble it through conc. ammonia pass the mixture over the hot Re salt , leave this gas running while it cools and get the metal. You might want to add a flux so the powdered metal is covered up and doesn't intermediately react with the air when you open the tube.

watson.fawkes - 30-3-2010 at 13:51

Quote: Originally posted by unionised

Why worry about water?
The same reference I quoted says you can reduce just about any rhenium compound with hydrogen, that would include the oxide so the reaction ReOx +x H2 --> Re + x H2O.

The amount of water determine (1) how fast the reaction proceeds and the amount of reagent gas required and (2) how far along the reaction proceeds to completion before it hits steady-state. My original point was that if you need highly pure material, you've got to worry about water content. If want a sample for an element collection, as Jor does, I would not doubt that the oxides could be reduced out adequately for that purpose. (Although if you want it nice and shiny, you might want to do more, and then ampule it.) If you have a different purpose, say you're using it as an alloying element, oxide purity can matter more.

watson.fawkes - 30-3-2010 at 14:25

Quote: Originally posted by not_important

Even iron will do the same[...]
It does reduce the fire hazard, [...]
And yes it likely takes a good drying, but so would H2.[...]
Neither should need to be completely free of water or oxygen, [...]

I agree with all this. My question, still open to me, is which of these works better in practice in a home lab context. There's the relative cost of reagent gases per gram of reduced metal. I would guess that, at small quantities of reductant, this has more to do with minimum flow through the apparatus than with molar efficiency. There's the cost of assembling and operating a drying train, if needed at all. There's the furnace and the plumbing. There's waste gas flare-off and ventilation. I don't see that one or the other reagent gases (H2 or NH3) is obviously superior; each seems to have different advantages and disadvantages.

There's one open chemistry question that could tip the balance. It's conceivable that the reduction temperature in NH3 could be significantly lower than for H2. If it's low enough to do the reduction in borosilicate glass, that would be an advantage.

Jor - 30-3-2010 at 15:19

I will only do the reaction in glass, if that's not possible, I will not do it, because i don't have the time to assemble a metal device for doing the reaction.

The other real problem is the absorption of hydrogen by rhenium making it pyrophoric. And i don't have a vaccuum.

So likely I will try the zinc route first, or maybe even with molten sodium (very small amounts). Or hydrazine/hypophosphite (although I prefer to keep the latter). Then I might try the NH3 and H2 routes but I'm really not sure yet, because of safety issues.
And finally if everything fails (wich i doubt) I will try to find someone here in the Netherlands who can do the H2 reduction for me, letting him have a few grams of the salt.

Oh and is anyone prepared to trade some of the perrhenate for ruthenium/iridium/platinum/palladium? I'm willing to trade up to 30 grams.
I'm also willing to sell about 2-3g of gold powder (99+% pure).

[Edited on 30-3-2010 by Jor]

[Edited on 30-3-2010 by Jor]

Fleaker - 31-3-2010 at 04:11

Jor,

I can trade you Re for your perrhenate.

DJF90 - 31-3-2010 at 04:58

The king of precious metal extraction speaks

Fleaker - 31-3-2010 at 16:17

King? I think not! My first reduction of ammonium perrhenate failed miserably! I nearly burnt the hell out of myself upon removing it from the quartz tube when it took fire (hence my advice above to pull a vacuum and heat it up to remove hydrogen!!!) I should mention, the [mixed] oxides of rhenium are beautiful colours!!

Furthermore, when the stuff was sent out to be ebeam melted, my shoddy and quite unsatisfactory neophyte reduction caused the operator quite some trouble. I'm sure that guy still wants a piece of me. Lesson learned--always find out how much ammonia has been produced!

watson.fawkes - 1-4-2010 at 06:20

Quote: Originally posted by Fleaker

I nearly burnt the hell out of myself upon removing it from the quartz tube when it took fire (hence my advice above to pull a vacuum and heat it up to remove hydrogen!!!)

What temperature was it (approximately) when you removed it?

Fleaker - 1-4-2010 at 12:29

Room temperature--it had cooled under a stream of H2.

The rule of never cooling a PGM in hydrogen is one that should be observed, especially if it is a fine powder (black). I don't think platinum black will actually "burn"-- I suspect it just promotes combustion. Palladium black does turn to PdO to some extent ( I know this because it does not all dissolve in aqua regia, and there is always a mass gain).

watson.fawkes - 1-4-2010 at 14:22

Quote: Originally posted by Fleaker

Room temperature--it had cooled under a stream of H2.

That's both neat and dangerous. So my next question: While vacuum would certainly work to extract H2, wouldn't an inert sweep gas also work? Assuming Raoult's law holds at the metal surface, I'd say so, but I have a nagging doubt in my assurance that it does. I ask because, practically speaking, it's a little easier to set up a second gas under positive pressure than to exhaust an uncondensable, inflammable gas with a pump. As for a gas, at room temperature I'd imagine that Ar, N2, and CO2 would all work, just nothing too reactive.

12AX7 - 1-4-2010 at 15:41

If you let in air slowly, will it catalytically oxidize the adsorbed hydrogen? Do it slowly so it doesn't sinter in the center or anything.

Tim

Jor - 2-4-2010 at 12:13

Ok, I just came home and rushed in the lab ASAP!

I put 100mg of ammonium perrhenate, together with about 450mg of 300 mesh zinc powder (I think it's pure but I'm not 100% sure) in a porcelain crucible, and heated for about 1 minute (maybe a little bit more) at red heat. I then left it to cool down. The mixture formed a cement like layer, wich however not that hard to remove. It was placed in a beaker and briefly boiled with some 6M HCl, until fizzing stopped. I was left with a small amount of black powder. The supernatant was decanted, the precitpate washed with water, and conc. HNO3 was added. The precitpate dissolved, with formation of NO2. the HNO3 was hot was the beaker was still hot. There was a formation of a orange color. This must be some NOCl (didn't rinse all HCl off), as perrhenate is produced when rhenium is oxidised by HNO3, wich is colorless.

I will just have to make sure that my zinc powder does completely dissolve in HCl, to rule out the possibility that it contains an impurity, but I think it is pretty pure.

So now the only question remains:
Is this rhenium metal or some oxide?

How can I test this? Do the oxides of rhenium dissolve in hot HCl? If they do, I certainly have Re powder here!

woelen - 3-4-2010 at 03:52

Solid-solid reactions usually give impure products, so even if you have Re-metal expect it to contain quite some contamination (e.g. encapsulated NH4ReO4, maybe zinc, amalgated with the Re-metal).

In order to assess the purity of your material in terms of zinc-content, dissolve some in HNO3 and then slowly add dilute NaOH, drop by drop. If the solid material contains amalgated zinc, then you'll get a white precipitate of Zn(OH)2, while perrhenic acid does not give a precipitate. Do not add a lot of NaOH at once, because Zn(OH)2 redissolves in excess solution of NaOH.

I'm not sure whether you have Re-metal though. ReO2 is a black solid (like MnO2), but it is much less reactive and the heat might have made it somewhat inert. ReO2 certainly will not dissolve in conc. HCl giving Cl2. If it dissolves then I expect that some chloro-complex is formed, but I do not know about a specific complex.

[Edited on 3-4-10 by woelen]

Jor - 3-4-2010 at 07:53

I think this was not a solid solid reaction, as the zinc likely melted, but i'm not sure. I don't want to take the cover from the crucible, as i'm not sure if the zinc will ignite at those temperatures.
The only way I can think of now to see if I'm making the metal or oxide is by precise weighing, but my analytical balance broke, so I can't do this...
But don't you think that boiling with HCl simply removes all the zinc from the rhenium (oxides) ? Just like when you heat a Ni/Al alloy in NaOH, you remove the Al (leaving raney nickel).

[Edited on 3-4-2010 by Jor]

woelen - 3-4-2010 at 12:52

It depends on how finely divided the materials are and whether particles of zinc can be encapsulated by a non-reactive layer of Re or Re-oxide. Encapsulation of solid materials by other solid materials is a common problem in this kind of reactions. If e.g. a slag is formed of Re and/or Re-oxide then this could lump up and form a protective layer around small globules of zinc and ammonium perrhenate. I'm not sure though about the melting points of all involved chemicals and about the temperature of your mix. These are just considerations, based on general knowledge.