Sciencemadness Discussion Board - Dissolving germanium metal

Dissolving germanium metal

Mixell - 20-2-2011 at 10:08

I've tried dissolving germanium powder with various acids, nothing worked so far.
The following mixtures and acid were used:
concentrated nitric acid (and later hydrogen peroxide was added).
concentrated sulfuric acid (and later hydrogen peroxide was added).
and Aqua Regia wish has proven to dissolve the germanium, but extremely slowly (the other solutions were completely useless).
The experiment was done at ~40C.

Can anyone suggest a substance that could efficiently dissolve germanium?
Maybe I can try heating the aqua regia to 60-90 C (although that will speed up the harmful gas production)?

blogfast25 - 20-2-2011 at 10:44

Trying to dissolve germanium in acids is a bit like trying to dissolve carbon or silicon in acids: as useless as trying to fit wheels to a tomato.

Germanium isn’t, despite its deceptive appearance, a metal. It’s a metalloid (‘half-metal’, if you prefer) and behaves more like a non-metal than an actual metal.

Germanium can be dissolved by fusing with alkalis, possibly also slowly by reacting with very hot, very concentrated (50 %) NaOH or KOH, resulting in the analogs of carbonates and silicates, the germanates: GeO3(2-). In the presence of water these should promptly hydrolyse to GeO2.

Wiki states that it dissolves slowly in conc. H2SO4 but I'm betting all you'll get from that is insoluble GeO2... Hot aqua regia will also produce GeO2 (at best).

Forget about making conventional salts from germanium: most compounds in the prevalent IV oxidation state are non-ionic (see halides, e.g.) and easily hydrolysed back to GeO2...

Where did you obtain the germanium?

[Edited on 20-2-2011 by blogfast25]

garage chemist - 20-2-2011 at 10:52

After you've obtained GeO2 by acidification of a germanate solution, this can be reacted with saturated aqueous HCl and GeCl4 distilled off. This is an important contrast to the homologous SiCl4, whose hydrolysis to SiO2 is irreversible.

Mixell - 20-2-2011 at 10:56

Aqua Regia is starting to show some result when heated to 50-70C (it dissolves a few miligrams in 20 mins, noticeable but useless).
I've also tried burning germanium to get some germanium oxide and dissolving it in water to get germanic acid (H4GeO4), the quantities here were also extremely tiny, I didn't find the time to test the solution for pH, maybe I will do it in a few days.
The germanium was given to me by my granfather (he was a chemistry professor at a zinc production plant (that also contained a research laboratory) in Soviet Russia.

blogfast25 - 20-2-2011 at 10:59

Quote: Originally posted by garage chemist

After you've obtained GeO2 by acidification of a germanate solution, this can be reacted with saturated aqueous HCl and GeCl4 distilled off. This is an important contrast to the homologous SiCl4, whose hydrolysis to SiO2 is irreversible.

Are you sure of this? Wiki mentions that but it also mentions how easily the tetrahalides are hydrolysed. The two statements seem highly contradictory to me…

blogfast25 - 20-2-2011 at 11:02

Quote: Originally posted by Mixell

The germanium was given to me by my granfather (he was a chemistry professor at a zinc production plant (that also contained a research laboratory) in Soviet Russia.

Lucky you! Ge is indeed the by product of several metal smelter's 'fly ash'. Some fly ashes from certain types of coal also contain Ge...

not_important - 20-2-2011 at 11:11

Aqua regia will work, as will strong hydrobromic acid with the slow addition of Br2; in either case it needs to be 5060 C, and done under reflux to avoid loss of the tetrahalide as noted by garage chemist. In analytical chemistry this was a method of separating Se, As, Ge, and sometimes Sb, from other metals. Hot sulfuric acid does work, but is quite slow.

Hot concentrated NaOH solution, or molten NaOH - add Ge __slowly__- is a better method.

garage chemist - 20-2-2011 at 11:32

Brauer gives a method for GeCl4 from GeO2 and HCl, but it is found as part of the GeO2 preparation.
GeO2 is mixed with conc. (37%) HCl and very slowly heated and distilled. If As is present as an impurity, a slow stream of chlorine has to be bubbled in to oxidise it to arsenic acid and prevent volatilization as AsCl3.
A biphasic distillate collects, the upper layer is 20% aqueous HCl. The layers are separated and the lower one is poured into a large volume of ice, shaken, and left to stand over night.
GeO2 precipitates and is filtered, from the acidic supernatant, dissolved Ge is precipitated as GeS2 with H2S.

Also, methods for reclaiming Ge from residues are given.
Elementary Ge is best worked up by chlorination with Cl2 at 450-500°C followed by hydrolysis of the GeCl4.
GeO2 easily dissolves in aqueous alkali, from the acidified solution, GeS2 can be precipitated by gassing with H2S. Precipitation as GeS2 is the best method to work up dilute aqueous Ge solutions.

Eclectic - 20-2-2011 at 13:21

@Blogfast: You can get germanium from Ebay as IR lenses

UnintentionalChaos - 20-2-2011 at 13:41

Quote: Originally posted by Eclectic

@Blogfast: You can get germanium from Ebay as IR lenses

You can also get it from ebay as germanium.

Mixell - 22-2-2011 at 06:01

Is it possible to extract sodium germanate (Na2GeO3) from a solution of NaOH?
And how do I exactly extract the solid GeO2 from a germanate solution?
Slowly adding acid until the germanium dioxide precipitates out of the solution?
I assume the reaction between germanium and sodium hydroxide is as following:
2NaOH +Ge +H2O ---> Na2GeO3(aq) + 2H2 ?
Or the reaction may be more complex with varying products depending on the exact conditions?

blogfast25 - 22-2-2011 at 10:20

Quote: Originally posted by Mixell

Is it possible to extract sodium germanate (Na2GeO3) from a solution of NaOH?
And how do I exactly extract the solid GeO2 from a germanate solution?
Slowly adding acid until the germanium dioxide precipitates out of the solution?
I assume the reaction between germanium and sodium hydroxide is as following:
2NaOH +Ge +H2O ---> Na2GeO3(aq) + 2H2 ?
Or the reaction may be more complex with varying products depending on the exact conditions?

Just gently acidify the solution of germanate (with excess NaOH) until the white GeO2 drops out. Actually isolating the germanate may not be possible because it's so sensitive to hydrolysis.

The balanced reaction is correct: 2NaOH (aq) + Ge (s) + H2O (l) ---> Na2GeO3 (aq) + 2H2 (g). No other reaction products can form as far as I know...

Mixell - 22-2-2011 at 12:29

According to wikipedia, GeO2 can be dissolved in water to give germanic acid (H4GeO4).
May be it is possible to add NaOH to a solution of germanic acid (with a proper mole ratio) and get Na4GeO4? Could be very interesting to try this.

blogfast25 - 22-2-2011 at 12:45

Quote: Originally posted by Mixell

According to wikipedia, GeO2 can be dissolved in water to give germanic acid (H4GeO4).
May be it is possible to add NaOH to a solution of germanic acid (with a proper mole ratio) and get Na4GeO4? Could be very interesting to try this.

Not in my version, it doesn’t.

Mine says:

”The dioxide, GeO2 can be obtained by roasting germanium sulfide (GeS2), and is a white powder that is only slightly soluble in water but reacts with alkalis to form germanates.”

http://en.wikipedia.org/wiki/Germanium#Chemistry

Think periodic table of the elements: SiO2 is completely insoluble in water, soluble only in HF or alkalis (but with great difficulty), SnO2 is completely insoluble in water and acids, but soluble in very hot, concentrated alkalis (and can be fused with them to form stannates). What in your book should make GeO2 any different???

Mixell - 22-2-2011 at 12:51

Well, so it possible that the information given in wikipedia is wrong, but here it is:
"The rutile form of germanium dioxide is more soluble than the hexagonal form and dissolves to form germanic acid, H4GeO4 or Ge(OH)4."

http://en.wikipedia.org/wiki/GeO2

blogfast25 - 22-2-2011 at 13:00

'dissolves' is fairly non-specific: no actual solubility numbers are given. I suggest you treat even freshly precipitated GeO2 with water and you will find that 'insoluble in water' is probably an apt term...

Mixell - 22-2-2011 at 13:03

Numbers are given:
5.2 g/l (25 °C)
10.7 g/l (100 °C)

Its quite minute, but its something.

blogfast25 - 22-2-2011 at 13:40

5.2 g/l at RT that's 0.5 %. More than I would have given it credit for.

But why do you want a solution of Ge(OH)4?

Mixell - 22-2-2011 at 13:46

I would like to try to isolate some of its salts, if its possible of-course, could be quite interesting to see what properties they have (would be even better to find some information on them, but I didn't manage to do so).

blogfast25 - 22-2-2011 at 14:30

Well, more or less forget about salts in which Ge plays the tole of cation. But maybe Na2GeO3 can be isolated, who knows? Dissolve some freshly precipitated GeO2 in an exact equivalent amount (acc. the reaction equation you wrote) of quite concentrated NaOH or KOH (allow for plenty time and stirring), filter off any turbidity and gently evaporate. If the product is crystalline and redissolves in pure water it's likely to be a germanate.

Or ice the solution, instead of evaporating...

[Edited on 23-2-2011 by blogfast25]

woelen - 23-2-2011 at 23:47

I don't believe the numbers in Wikipedia. I have 50 grams of GeO2 and did quite a few experiments with this. It does not dissolve in water, at least not visibly. If it really could dissolve in water at 10 grams per liter, then I would be able to dissolve a spatula full of this solid in 100 ml of water. No, it doesn't. I am inclined to think that the numbers in the Wiki page must be in mg/liter.

In conc. HCl indeed quite some GeO2 can be dissolved. It dissolves slowly, but I managed to dissolve several 100's of mg in a test tube containing 10 ml of acid. When this solution is diluted, then the liquid becomes milky/opaque and a very fine white precipitate settles at the bottom.

blogfast25 - 24-2-2011 at 08:51

Sounds plausible, Woelen. I'm very skeptical about GeO2 being even slightly soluble in water: see SiO2 and SnO2...

[Edited on 24-2-2011 by blogfast25]

ScienceSquirrel - 24-2-2011 at 09:59

It seems that there is a soluble form and an insoluble form, also solution is achieved by boiling under reflux for some time!

http://www.krasgermanium.com/products?article=prod20

and

http://www.teck.com/DocumentViewer.aspx?elementId=115437&...

Try a search on solubility germanium dioxide.

Material scientists, right cards!

blogfast25 - 24-2-2011 at 12:41

The combined solubility data from both sources are 1 g/100 ml at RT, 4.5 g/L at 25 C and 10.7 g/L at 100 C. Hardly a testimony to solubility, in my book...

Inorganic - 24-2-2011 at 17:10

I have that same problem with Ge...

I attach two publications were trying to dissolve Ge...

Attachment: JES000492.pdf (544kB)
This file has been downloaded 1143 times

Attachment: JES000508.pdf (469kB)
This file has been downloaded 4754 times

ScienceSquirrel - 25-2-2011 at 05:11

Quote: Originally posted by blogfast25

The combined solubility data from both sources are 1 g/100 ml at RT, 4.5 g/L at 25 C and 10.7 g/L at 100 C. Hardly a testimony to solubility, in my book...

I did not say it was soluble, quite the opposite, but it seems that refluxing will get it to dissolve a little.

blogfast25 - 25-2-2011 at 07:56

Assumimg these numbers are correct, that's quite remarkable IMHO...

ScienceSquirrel - 25-2-2011 at 08:05

One of the uses for germanium dioxide is as a catalyst for making PET, it seems that making solutions is quite difficult so manufacturers supply ready made stable solutions.

http://www.polyester-technology.com/Publication/publication_...

Mixell - 21-3-2011 at 12:28

I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?
I'll post some pictures latter, if I will manage to locate the cable from the camera.

blogfast25 - 21-3-2011 at 13:43

Quote: Originally posted by Mixell

I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?
I'll post some pictures latter, if I will manage to locate the cable from the camera.

Where did you get this idea from?

Mixell - 21-3-2011 at 13:48

By comparing the reduction potentials of germanium and copper. Why? Something isn't right?

EDIT: The potentials, from Wikipedia:
Ge4+ + 4 e− Ge(s) +0.12.
Cu2+ + 2 e− Cu(s) +0.34.

And its working, solid copper particles precipitate and the solution gradually becomes transparent.

[Edited on 21-3-2011 by Mixell]

blogfast25 - 21-3-2011 at 14:11

So the cell potential for:

2 Cu2+ (aq) + Ge(0) (s) === > 2 Cu(0) (s) + Ge4+ (aq)

… is about + 0.34 + (- 0.12) = + 0.22 V (thus ΔG < 0, reaction proceeds spontaneously), so plating out Cu with Ge metal is possible. Never occurred to me… But isolating your Ge(SO4)2 from solution to solid will be the problem, at least without hydrolysis. Ge is a half-metal, remember?

[Edited on 21-3-2011 by blogfast25]

Mixell - 21-3-2011 at 14:54

Update: I got a nice clear solution of germanium sulfate, but decided to add more copper sulfate due to the fact that a lot of unreacted germanium was left in the vessel.
The reaction took approximately 6 hours at boiling point to react about 200 mg of germanium.

Mixell - 21-3-2011 at 15:11

The hydrolysis is Ge(SO4)2 +2H2O--> GeO2 + 2H2SO4?
It would be helpful to find some information on germanium sulfate's decomposition/evaporation point. May be it is possible just to evaporate the water, or to dry with sodium hydroxide (if germanium sulfate is not too hygroscopic).

blogfast25 - 22-3-2011 at 06:16

I doubt very much if Ge(SO4)2 actually exists. In oxidation state +IV the element prefers to form anions like the germanate ion: GeO3(2-), not cations like Ge4+. In this it shows its metalloid character.

Here is an interesting Ge mineral, Schaurteite, Ca3Ge(SO4)2(OH)6•4H2O

http://webmineral.com/data/Schaurteite.shtml

… but the structural formula is deceptive: it should probably be re-written as 2CaSO4.CaGe(OH)6.4H2O; a double salt of CaSO4 and calcium germanate.

The germanates can be re-written as: GeO3(2-) + 3 H2O = Ge(OH)6(2-), so CaGe(OH)6 can be rewritten as CaGeO3.3H2O and the total formula for Schaurteite as 2CaSO4.CaGeO2.7H2O!

So despite the sulphate groups, no actual Ge(SO4)2 in sight either!

On evaporating (more or less regardless of conditions) you’ll obtain plain old boring GeO2 again…

[Edited on 22-3-2011 by blogfast25]

Mixell - 22-3-2011 at 09:06

Well, I do have a solution of some sort.
Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?
I'll check the solution for acidity tomorrow, and also try to evaporate some part of the solution and test the resulting product, then I'll try to dissolve the solid product again, and see what happens, if it will not hydrolyze immediately, I will check the solution for a presence of sulfate.

blogfast25 - 22-3-2011 at 10:05

Quote: Originally posted by Mixell

Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?

Yes. In a nutshell.

[Edited on 22-3-2011 by blogfast25]

Mixell - 22-3-2011 at 10:52

Sorry, but I lack the knowledge of the meaning of some phrases, can you explain to me what "in a nutshell" means?

And back to the topic, I isolated the solution, it has a very-very faint yellow color, almost completely clear.

blogfast25 - 22-3-2011 at 12:15

'in a nutshell' = 'in short', 'in summary', 'basically', 'essentially'...

Mixell - 22-3-2011 at 12:38

Understood, thank you =)

The WiZard is In - 22-3-2011 at 16:11

Quote: Originally posted by Mixell

I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?

I'll post some pictures latter, if I will manage to locate the cable from the camera.

Me the Analogue guy — again.

The AG goes down the hallway and removes from a bookshelf :—

Greenwood & Earnshaw
Chemistry of the Elements
Pergamon Press 1984 [There is a latter ed.]

"An ustable sulfate Ge(SO4)2 is formed in a curious reaction
when GeCl4 is heated with SO3 in a sealed tube at 160o."

GeCl4 + 6SO3 ---> Ce(SO4)2 + 2S2O5Cl2.

blogfast25 - 23-3-2011 at 05:47

Interesting WiZ. Interesting also how the germanium transmutates to cerium!

The WiZard is In - 23-3-2011 at 06:47

Quote: Originally posted by blogfast25

Interesting WiZ. Interesting also how the germanium transmutates to cerium!

Through what Einstein (1947) called spukhafte Fernwirkung
(Spooky action at a distance).

blogfast25 - 23-3-2011 at 12:49

Quote: Originally posted by The WiZard is In

Quote: Originally posted by blogfast25

Interesting WiZ. Interesting also how the germanium transmutates to cerium!

Through what Einstein (1947) called spukhafte Fernwirkung
(Spooky action at a distance).

While maintaining the oxidation state too! Spookier than a direct hit on an MI6 safe house!

Mixell - 25-3-2011 at 13:33

Well, I'll post that in this topic, silicon and germanium have similar properties.
Anyway, I got my hands on about 50g of ultra-pure silicon (the kind used to make computer chips), any suggestion what to do with it? Its a pretty inert element, but I think it still can be used for something interesting.

blogfast25 - 25-3-2011 at 13:46

For a start, dissolve it in strong alkali: it forms silicates.

Mixell - 24-4-2011 at 07:59

I dissolved some germanium in nitric acid, and I got a very faint green solution (with some germanium left at the bottom).
I added some hydrogen peroxide which immediately turned the solution from very faint green to brown-yellow (urine colored).
At the moment the germanium at the bottom is giving off a good amount of bubbles, but I think its just the hydrogen peroxide decomposing (the decomposition rate gets bigger with time).
The interesting thing, is that my germanium sulfate solution is in the same color as this one, maybe its the color of the Ge4+ cation, so the hydrogen peroxide must have oxidized the former germanium cation (presumingly Ge2+) to Ge4+ ?
Or it formed some peroxo complex with a similar color to the Ge4+ cation?

blogfast25 - 24-4-2011 at 10:34

Dissolving Ge in nitric acid should, as far as I know, lead only to insoluble white GeO2 being formed (like its family member Sn, but Ge is even less inclined to make water soluble compounds).

The green, later brown yellow colour? Check the quality of your nitric acid: these colours are likely due to a contamination.

sternman318 - 24-4-2011 at 11:01

Quote: Originally posted by blogfast25

Quote: Originally posted by Mixell

Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?

Yes. In a nutshell.

[Edited on 22-3-2011 by blogfast25]

I may not follow or understand exactly, but if its in some equilibrium, couldn't you just add sulfuric acid to shift the reaction towards germanium sulfate?

Mixell - 24-4-2011 at 11:28

The hydrogen peroxide and the nitric acid (as everything else) are laboratory grade (CP). And the germanium is clean too. So I don't know what, if any impurities could of caused that.I noticed some companies are selling germanium nitrate, so it must exist...
And about adding sulfuric acid to shift the equilibrium, germanium sulfate only exist in solution, so I'll just get a solution of sulfuric acid and germanium sulfate.

Mixell - 24-4-2011 at 11:59

Ok, the solution is back to its original, slightly green coloar, after adding a few drops of hydrogen peroxide it became yellow again. Maybe some unstable peroxo complex of germanium is formed?

blogfast25 - 24-4-2011 at 12:24

Never heard of one. That type of complex usually requires dangling d or f electrons (see transition metals and rare earths).

Mixell - 24-4-2011 at 12:55

Well, so I have no idea what is causing that yellow color...
Could it be Ge4+? But why does the color disappear?

thethule - 24-4-2011 at 16:03

Quote: Originally posted by Eclectic

@Blogfast: You can get germanium from Ebay as IR lenses

That seems like a very expensive way of getting some Ge. Would it even be pure ge? Even more expensive than getting it from element sales websites.

IrC - 24-4-2011 at 18:18

Maybe I missed it in 3 pages but I am unsure what compound of Ge Mixell is going for. Back in my youth with no internet and no information (or very little) on Ge chemistry I was trying to alloy Ge. OK, so I thought it would act like a metal, and if I could liquefy it I was going to melt other metals into it. Heating a 50 gm chunk of pure Ge in a ceramic boat (no flame touching Ge), it began burning in air. Much like a Mg fire I might add. Clouds of smoke later I had a powder with both yellow/orange and white appearance. Reading about toxicity of at least one Ge - Oxygen compound I left the room, open for a half day. Still alive all these years later. Anyway, was wondering if this carried out in a fume hood to give an oxide of Ge would not be a simple, fast starting point to then further react the powder to form other salts of Ge? Much of the Ge went out the door as oxide smoke so possibly routing this smoke through another vessel with the reactants needed for some other transition might also be a viable procedure?

The Jpeg's below show the only information I possessed back then about Ge chemistry, might still be of use here.

[Edited on 4-25-2011 by IrC]