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dann2
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Antimony and Tin Nitrates
Hello,
I need some Antimony and Tin Nitrate. To my surprise neither of the compounds are listed in my local chemical suppliers catalogue. Googling doesent
help. It seems Tin Nitrate is an elusive compound.
Can anyone direct me to where I can purchase or make these's compounds.
Cheers,
Dann2
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not_important
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This is because they are only stable in strong nitric acid. Diluting the acid causes hydrolysis and results in the precipitation of hydrated oxides.
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Rosco Bodine
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Tin fishing weights are pure tin and can be easily dissolved in nitric acid , or in heated HCl .
The split shot form which comes in the little ziplock
plastic bags is the most economical form .
Antimony is more difficult to find , but is used for
casting ornamental small jewelry boxes and some music boxes also , IIRC . I think it is also used to harden lead used for bullet casting so you might
find ingots through a black powder weapons sporting goods supplier .
http://www.theantimonyman.com/price.htm
[Edited on 20-4-2007 by Rosco Bodine]
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not_important
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Quoting Mellor's Inorganic "Dilute nitric acid has scarcely any action, but it possibly forms an unstable antimony nitrate. Concentrated nitric acid
does not dissolve the metal but rather oxidises it to insoluble oxides.
Antimony oxide can be purchased at pottery supply places. It will dissolve in strong HCl forming the trichloride or strong HNO3 forming the nitrate.
Adding much water to the chloride results in the formation of the oxychloride, moisture will cause the formation of the basic nitrate making the
solution cloudy. See bismuth nitrate for more details, bismuth is more metallic than antimony and its salt hydrolyse a bit less readily.
Tin metal with nitric acid - highly concentrated HNO3 has little acid on tin, moderate strength (sg 1.24) reacts to form the hydrated Sn(IV) oxide,
sometimes called metastannic acid, cold dilute acid will form stannous nitrate with some stannic nitrate, but this soon give a precipitate of
metastanic acid. All simple tin salt tend to hydrolyse easily, at best they can be crystallised from acid solutions as hydrate. You can make them,
but they don't keep.
When suppliers don't list a compound, and the reference texts at best list it as decomposing in cold or hot water, you are being informed that it is
difficult to make, purify, and keep around. If the application requires a not-too-acid solution of these metals' nitrates, then it isn't going to
happen.
Quoting from http://dissertations.ub.rug.nl/FILES/faculties/science/2004/...
| Quote: | | For example, tin nitrate is a rather exotic compound, while tin acetate is easily available. The unavailability of certain nitrates limits the number
of ceramic phases that can be grown by the deposition technique described in this work. It seems, however, likely that replacing tin nitrate by tin
acetate in the recipe can yield precursor solutions for tin containing ceramics, such as indium tin oxide (ITO). |
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Magpie
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Antimony and tin are available off eBay > Business & Industrial > Metal & Alloys.
Rosco, I wasn't aware that pure tin was available as fishing weight -that would be a nice OTC source. 
I see that bismuth fishing weights are also available.
[Edited on by Magpie]
The single most important condition for a successful synthesis is good mixing - Nicodem
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Xenoid
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Antimony Nitrate
| Quote: | Originally posted by not_important
Antimony oxide can be purchased at pottery supply places. It will dissolve in strong HCl forming the trichloride or strong HNO3 forming the nitrate.
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Yes! My abridged Mellors states antimony trioxide (Sb2O3) will dissolve in cold, concentrated nitric acid to form antimony nitrate.
I have thus gone ahead and tried this with pottery Sb2O3 and 68% nitric acid. There is no apparent reaction, nothing is happening, just a yoghurt-like
sludge at the bottom of the beaker. I mixed stoichiometric amounts (actually slightly more of my valuable nitric acid). What have I done wrong.
Is it possible to make antimony citrate, acetate, oxalate or tartrate. I need a soluble antimony compound that decomposes below about 320 oC.
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The_Davster
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I have in the past purchased a solder of 95% tin, and 5% Sb. When used as the anode for electrolysis(10%HCl electrolyte), Sn is dissolved and
instantly plates out on the cathode as pure tin. Sb falls off the anode as a black powder. Beware Sb build up on the anode, as if it completly
covers the electrode surface, it will oxidize to Sb2O3.
Or, ventilation permitting, such an alloy can be dissolved in hot HCl with some loss of antimony as toxic stibine gas. You get a solution of SnCl2
and black Sb powder remaining.
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not_important
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The citrate and tartrate complexes of antimony are soluble and reasonable stable in solution.
The ceramics grade oxide can be a bit difficult to dissolve. I've had best luck by making the chloride from it, distilling the chloride to leave iron
and such behind, dissolving the chloride in roughly 10% HCl, and pouring the solution into aqueous ammonia to get freshly precipitated hydrated oxide.
Wash that with dilute ammonia to remove chloride, then dissolve in the acid of choice.
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Xenoid
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| Quote: | Originally posted by not_important
The citrate and tartrate complexes of antimony are soluble and reasonable stable in solution.
The ceramics grade oxide can be a bit difficult to dissolve. |
@ not_important
Why is the pottery oxide proving difficult to dissolve in nitric, any ideas!
Any ideas on making the citrate and/or tartrate, I have both acids of course!
Oh! I've just looked in my abridged Mellors and it says Sb2O3 forms a solution with tartaric acid, hard to believe when it doesn't want to dissolve in
nitric.
It's also supposed to dissolve in hot, conc. sulphuric acid forming the sulphate.
Please don't tell me I have to go through some long torturous process of re-precipitating the Sb2O3 before it will dissolve!
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woelen
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I also have done quite a few experiments with pottery grade Sb2O3. I found it to be fairly pure, much better than I expected. It, however, only
dissolves well in concentrated HCl, any other acid seems to be not dissolving it. Such a solution in 30% HCl is perfectly colorless, indicating that
at least it does not contains lots of transition metals like iron or copper.
When such a solution in concentrated HCl is diluted, then it becomes milky and hydrated oxychloride precipitates. This hydrated oxychloride can be
rinsed, and then dissolves quite well in a solution of tartaric acid. The resulting solution contains the tartrate complex of antimony, and besides
that, it also contains quite some chloride as well. If you want purer material, then you indeed have to go the way with ammonia. The tartrate seems to
be so stable that even in almost neutral solution it does not precipitate and it is possible to isolate the tartrate compound. The potassium salt of
this complex is listed in the catalogues of many chemical suppliers.
So, yes, Xenoid, you have to put some effort in this, if you want reasonably pure antimony compounds besides the oxide. Antimony is not an easy
element in aqueous solution, it suffers sooooo much from hydrolysis.
The reason why Sb2O3 dissolves so well in conc. HCl and not in the other common mineral acids of similar concentration is that with HCl it easily
forms a complex, SbCl4(-), which it cannot form with nitrate or sulfate. For this reason, even the rather inert pottery grade stuff dissolves easily
in conc. HCl.
Btw, a nice experiment is to bubble some H2S through a solution of Sb2O3 in conc. HCl. That is rather surprising. It also works if some Na2S is
dissolved in 5% HCl and this is added to the solution of Sb2O3 in 30% HCl. A bright yellow, or even somewhat orange precipitate is formed.
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Xenoid
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| Quote: | Originally posted by woelen
When such a solution in concentrated HCl is diluted, then it becomes milky and hydrated oxychloride precipitates. This hydrated oxychloride can be
rinsed, and then dissolves quite well in a solution of tartaric acid. The resulting solution contains the tartrate complex of antimony, and besides
that, it also contains quite some chloride as well. If you want purer material, then you indeed have to go the way with ammonia. The tartrate seems to
be so stable that even in almost neutral solution it does not precipitate and it is possible to isolate the tartrate compound. The potassium salt of
this complex is listed in the catalogues of many chemical suppliers.
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@ woelen
My abridged Mellors also mentions potassium antimonyl tartrate (Tartar Emetic). It claims this is made by boiling potassium hydrogen tartrate (cream
of tartar) with antimony trioxide. On cooling the concentrated solution, it separates in octahedral crystals. I was considering this simply prepared,
soluble antimony compound for doping my MnO2. I guess the organic component decomposes to CO and H2O, I'm not sure what would become of the K. However
I have found it melts at about 330 oC. so it is surprisingly stable, I don't know what the decomposition temperature is, but it is clearly above my
desired 320 oC.
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not_important
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| Quote: | | I guess the organic component decomposes to CO and H2O, I'm not sure what would become of the K. |
At that temperature range you'll end up with K2CO3
The ammonium equivalent might be useful, if it crystallises well enough; heating will result in just antimony oxides.
Those organic acids form complexes, the multiple carboxy and hydroxy groups are what does the trick. You could try direct dissolution of antimony
oxide in aqueuos solution of the organic acid, no promises.
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chloric1
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The tartrate intermediate shows promise especially since I can obtain tartrates cheaply.
Does the antimony have to be trivalent? If not you could dissolve pottery grade Sb2O3 in 30% HCl and then add 68% HNO3 and add this to your other
nitrates. This mixture should facilitate the complete removal of the chloride and leave only oxides on pyrolysis.
Fellow molecular manipulator
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Xenoid
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I'm not sure where I'm going with this but pottery grade Sb2O3 does indeed readily dissolve (forms complex antimony tartrate) in tartaric acid
solution. I dissolved 3 heaped teaspoons of tartaric acid in 100mls water, brought to boil and slowly stirred in about 2 level teaspoons of Sb2O3 most
of which dissolved. The slightly cloudy remaining solution was filtered and evaporated down to about 50 mls. I'm waiting to see what crystallises. I
guess it works because Sb2O3 is actually very slightly soluble, enough to get "complexed away". I put a bit of the solution on some metal and
pyrolised it. It seemed to swell up a lot which could disrupt any coating process, it disappeared in a final puff of smoke and a peculiar smell - ROFD
(rolls on floor dying). I need to investigate this a little more.
Rosco suggests using ammonium antimonate - I'm not sure how to make it!
Another possibility are the antimony alkoxides - eg methoxide, ethoxide, isopropoxide and butoxide - but I'm not sure how to make them either.
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chloric1
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| Quote: | Originally posted by Xenoid
I put a bit of the solution on some metal and pyrolised it. It seemed to swell up a lot which could disrupt any coating process, it disappeared in a
final puff of smoke and a peculiar smell - ROFD (rolls on floor dying). I need to investigate this a little more.
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I really don't like the sound of that Sounds too messy and possibly dangerously toxic. The tartrate possibly could have decomposed into aldehydes which
really scare me. Read the MSDS on acetaldehyde.
Fellow molecular manipulator
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Xenoid
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| Quote: | Originally posted by chloric1
I really don't like the sound of that Sounds too messy and possibly dangerously toxic. The tartrate possibly could have decomposed into aldehydes which
really scare me. Read the MSDS on acetaldehyde. |
Oh chloric! You had better give up chemistry right now if that worries you!
Nitrogen oxides are evolved during the pyrolysis of the nitrates. Your coating procedures should be done in an open, well ventilated area. We are only
dealing with milligram quantities for each coat after all...
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Rosco Bodine
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@Xenoid ,
You could just dissolve the Sb2O3 in HCl , oxidize with H2O2 to the higher state ,
then neutralize with ammonia to get the hydrated oxide ,
(antimonic acid ) , filter , then redissolve in more ammonia
to get the ammonium antimonate .
I think the Sb2O3 will also dissolve directly in H2O2
to give a pentavalent Sb hydrosol .
I'll try to dig up the refs again , I'm sure I've posted them
before but it was months ago .
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Xenoid
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| Quote: | Originally posted by Rosco Bodine
I'll try to dig up the refs again , I'm sure I've posted them
before but it was months ago . |
Yeah! Thanks Rosco!
All the info. is on about page 4 of the "perchlorate (not) with graphite" thread!
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chloric1
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| Quote: | Originally posted by Xenoid
Oh chloric! You had better give up chemistry right now if that worries you!
Nitrogen oxides are evolved during the pyrolysis of the nitrates. Your coating procedures should be done in an open, well ventilated area. We are only
dealing with milligram quantities for each coat after all... |
Xenoid yes I am aware that baking of the nitrates releases dangerous nitrogen oxides. Like chlorine, chlorine dioxide etc they are rather acrid and
they will let you know if your ventilation is adequate or not. Acetaldehyde is rather different. It seems that acetaldehyde has, I don't know, less
of a warning affect. I have has some respiratory irritations from very minute exposures.
Was I being silly about my worries? Yes, but many organics can be insidious and extra measures of caution need to be implemented. Before you say it,
nitrogen dioxide is insidious too, but its brown colored and for our uses it will be mixed with free HNO3 vapor giving it more of an unpleasant acrid
smell. A little squirel cage can blow on the heating apparatus to a designated area or over a bisulfite solution. Bisufites absorb aldehydes and
ketones AFAIK. It also drops gold out of solution. Sodium Metabisulfite is the
lab technician's best friend
Fellow molecular manipulator
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Rosco Bodine
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stannous nitrate Sn(II) nitrate Sn(NO3)2
US3243385 Stannous Nitrate see example 2 (attached)
Here's the only reference which I have found to production
of the +2 valency tin nitrate . This could be useful in the
formation of a tin *nitrate* variant of the mixed valency inorganic polymer , of the Pytlewski patent US3890429 .
868 grams of granulated tin metal is dissolved in 6200 grams
of 30% HNO3 , temperature being kept below 20 C , resulting
in a clear solution of tin(II) nitrate .
The use of higher concentrations of HNO3 and higher temperatures produces a different product , a precipitate
of Sn(IV) oxide .
Evidently for obtaining the Sn(IV) nitrate , only an indirect route is practical , where hydrated stannic oxide
( stannic hydroxide Sn(OH)4 ?? or "stannic acid" ?? )
is precipitated from a solution of an Sn(IV) compound
upon careful *partial* neutralization by a base slowly with vigorous stirring , whereupon the hydrated stannic oxide is obtained as a precipitate ,
rinsed by decantation , and
drained , dried without heating . This hydrated stannic oxide
can then be used to neutralize HNO3 completely , keeping
temperature well below 50 C , to form a solution of Sn(IV) nitrate .
Evidently both valency forms of Sn nitrate are *unstable*
with regards to both pH and temperature , but this not necessarily a bad thing for our purposes , simply something which must be accommodated in
handling of the materials
during preparation and use . The materials may not be
storage stable either and may have to be freshly prepared
for use .
The mixed valency polymer may get around that issue .
Also in the presence of a third metal nitrate , stability
may be greatly enhanced so that storage stable compositions are possible . A third metal could be
perhaps Mn for example , *if* we are lucky  , or
perhaps Bi or Sb or Co .
Anyway , it is known that Fe , Cr , and Al do most definitely
stabilize the Sn(IV) nitrate , having a radical effect in that regard . So these alone would have interest , particularly
also with regards to a coating scheme for gouging rods .
The plural metal oxide gels which are thermal decomposition products via similar processes are reportedly strongly adherent to substrates , and that
would seem likely for
these materials as well when applied at an optimum dilution .
I wish I could be more definitive , but the literature is not extensive concerning this somewhat obscure information
which of course leaves experiments to be the best source or perhaps the only source of further information .
[Edited on 29-12-2007 by Rosco Bodine]
Attachment: US3243385 Stannous Nitrate see ex 2.pdf (154kB)
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chloric1
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myu lt5kjloutwwe33
Fellow molecular manipulator
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Rosco Bodine
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| Quote: | Originally posted by chloric1
myu lt5kjloutwwe33 |
translation ???
keyboard malfunction ????
[Edited on 29-12-2007 by Rosco Bodine]
Attachment: Tin nitrate related Gmelin.pdf (187kB)
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Rosco Bodine
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There is an interaction between iron nitrate , and metastannic acid which renders the metastannic acid
soluble . A similar effect occurs with aluminum nitrate and
chromium nitrate , and possibly for other nitrates derived
from metals which form oxides of the form M2O3 .
This was described by Frederik Hendrik van Leent , phd. ,
( Royal Dutch Chemical Society ) in 1898 .
Recueil des Travaux Chimiques des Pays-Bas et de la Belgique , 1898 , 17 , 86-93
Attached is an English language article related to this anomalous solubility for metastannic acid in the presence
of iron , aluminum , or chromium nitrates .
Attachment: Tin nitrate M2O3 nitrate soluble complexes.pdf (430kB)
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Rosco Bodine
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A more modern article tending to confirm van Leents observations regarding iron nitrate and SnO2 is attached .
[Edited on 6-1-2008 by Rosco Bodine]
Attachment: MSE06.pdf (373kB)
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Rosco Bodine
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The SnO2 precursor for the Iron doped SnO2 described in the above article is gotten from ordinary SnCl2 in aqueous solution subjected to atmospheric
oxidation by stirring in an open vessel with the solution exposed to air for a couple of days . The SnCl2 slowly hydrolyzes and oxidizes to a
metastannic acid hydrosol which is evidently reactive with
iron nitrate to form the same hydrosol , as that which would be gotten from dissolving a mixture of tin and iron in nitric acid .
This "soft chemistry" approach is an alternative route having more modern description of the product hydrosol , which is believed by me to be the same
material as Dr. van Leent described a hundred years ago .
Attachment: Chimie douce preparation, characterization and photocatalytic activity of nanocrystalline SnO2.pdf (176kB)
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