Sciencemadness Discussion Board

Antimony and Tin Nitrates

dann2 - 19-4-2007 at 16:12

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

not_important - 19-4-2007 at 21:49

This is because they are only stable in strong nitric acid. Diluting the acid causes hydrolysis and results in the precipitation of hydrated oxides.

Rosco Bodine - 20-4-2007 at 04:25

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]

not_important - 20-4-2007 at 08:31

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).

Magpie - 20-4-2007 at 09:19

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]

Antimony Nitrate

Xenoid - 17-12-2007 at 21:59

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.


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.

The_Davster - 17-12-2007 at 22:18

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.

not_important - 18-12-2007 at 00:02

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.

Xenoid - 18-12-2007 at 00:20

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!

woelen - 18-12-2007 at 00:37

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.

Xenoid - 18-12-2007 at 01:16

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.


@ 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.

not_important - 18-12-2007 at 02:24

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.

chloric1 - 18-12-2007 at 17:10

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.

Xenoid - 18-12-2007 at 17:36

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.

chloric1 - 18-12-2007 at 18:51

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.



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.:o:o

Xenoid - 18-12-2007 at 19:03

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.:o:o


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...

Rosco Bodine - 18-12-2007 at 21:52

@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 .

Xenoid - 18-12-2007 at 22:37

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!

chloric1 - 29-12-2007 at 07:33

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:D;):D

stannous nitrate Sn(II) nitrate Sn(NO3)2

Rosco Bodine - 29-12-2007 at 07:41

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 :P , 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 :D
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)
This file has been downloaded 504 times


chloric1 - 29-12-2007 at 10:27

myu lt5kjloutwwe33

Rosco Bodine - 29-12-2007 at 11:07

Quote:
Originally posted by chloric1
myu lt5kjloutwwe33

translation ??? :P
keyboard malfunction ????

[Edited on 29-12-2007 by Rosco Bodine]

Attachment: Tin nitrate related Gmelin.pdf (187kB)
This file has been downloaded 856 times


Rosco Bodine - 6-1-2008 at 08:27

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 - 6-1-2008 at 08:46

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)
This file has been downloaded 962 times


Rosco Bodine - 12-1-2008 at 07:36

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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abstract concerning van Leent article

Rosco Bodine - 12-1-2008 at 07:49

Here is an abstract summarizing the more extensive
article by van Leent concerning the anomalous "solubility"
of what is evidently a doped tin oxide dispersion , in
the form of a hydrated metastannic oxide sol .

I suspect that the van Leent material could accurately be called a clathrate , and similarly there are many other of the
doped tin oxide materials which are probably also
coordination compounds .

Edit: errata - ( coordination compounds ) should read
*inclusion compounds*

[Edited on 10-2-2008 by Rosco Bodine]

Attachment: Abstract for van Leent article.pdf (103kB)
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tentacles - 14-1-2008 at 14:23

So, Rosco, would metastannic acid (H2SnO3) be appropriate for a DTO precursor? The metastannic acid ppt that I had previously formed (and foolishly disposed of) was certainly water soluble, as long as there was not too much HNO3 in solution. It would seem like the decomposition would leave SnO2 (once the water was driven off)

[Edited on 14-1-2008 by tentacles]

Rosco Bodine - 15-1-2008 at 09:16

There seems to be no reason why the anomalously soluble
binary of metastannic acid with iron nitrate or perhaps chromium nitrate would not make an excellent
DTO precursor , since the metastannic acid is a higher intermediate in the pyrolytic decomposition sequence of other materials making their way to a baked DTO coating .

Iron is an already identified dopant of value for the DTO coatings useful as anodes . The proportions might have to be adjusted and the pH adjusted in mixture with other components so that everything stays soluble , but yeah I think it certainly has promise as a DTO precursor particularly since it forms a highly loaded "solution" ,
and the SnO2 is "already there" as the oxidized Sn not
requiring aerobic baking to develop , but only being baked to dehydration and sintering .

This contrasts with a chlorides precursor which has to hydrolyze and lose its chlorine , absorb oxygen and dehydrate to end up being SnO2 , which requires longer baking and higher temperature and would never really be so complete for the material buried deep in the layer .

So .....the binary composition with metastannic acid is a higher developed precursor , for the DTO .

dyer's tin nitrate described (more or less)

Rosco Bodine - 23-2-2008 at 10:29

This bit of obscure information was gotten from an english trade journal circa 1877 , The Textile Colourist , and is the
single description which I have been able to find concerning
a preparation of a reported *Stannic* nitrate , directly from
tin metal using nitric acid .

By this method the product appears to actually be a mixture
70% stannic nitrate with 30% stannous nitrate by analysis
methods of that era . The nitric acid concentration used is
reportedly d 1.17 , 28.5% HNO3 , with the dissolving of the tin 1 part in dilute acid 8 parts by weight , the process done
slowly so the exotherm is kept small , reaction temperature
kept below 100F according to one source and 60F according to another . The resulting soluble nitrates of tin is a solution of density 1.3 having 14% tin content as the metal .

From this source and from others there is indication that the presence of some HCl enhances the storage stability , also
the presence of NH4NO3 as a minor proportion . From other sources as a more general consideration with regards to the delaying of precipitation by hydrolysis of soluble tin compounds , a small percentage of iron as an impurity is
stabilizing . Less explicitly reported are indications that
also tartaric and acetic acids may be stabilizing against precipitation , all of these materials seeming to act as dispersants which hinder the nucleation and growth of colloidal particles , thereby hindering precipitation from hydrolysis .

Obviously this mixture of stannic and stannous 70/30 is
a mixed valency composition , and what would be the properties similarity with a single valency stannic nitrate
derived from a nitric acid neutralization of precipitated alpha stannic oyhydroxide is unknown . But one thing that is shown by this old reference is that the nitrates of tin in
at least the simplest case of reacting dilute nitric acid with tin metal , can be made and has sufficient stability if kept cool in storage , to be useful at least for some weeks time of keeping . And this may perhaps also be true for the 100%
stannic nitrate composition , particularly if some stabilizing benefit is accomplished by means of those things identified to hinder precipitation in storage being added as stabilizers .

Attachment: Dyers Stannic Nitrate.pdf (449kB)
This file has been downloaded 479 times


woelen - 23-2-2008 at 11:21

It is funny to read such an old article from 1800 or so. Especially all the different types of nitric acid are nice to read. Apparently, at that time, one could not make pure HNO3 and acid of different purities is regarded another type of acid.

Rosco Bodine - 23-2-2008 at 11:44

Yeah I noticed that the nitric acid which was regarded
as being the "good stuff" was Dutch manufacture ,
while the English made acid was simply adequate :D
and I'm sure you got a good laugh from that :D

A proper chemist should only use politically correct HNO3 ;)

[Edited on 23-2-2008 by Rosco Bodine]

microcosmicus - 23-2-2008 at 12:48

As for different concentrations being regarded as different chemicals, remember
that Lavoisier thought that acetic acid and glacial acid were different chemicals.
This and similar confusions in elementary chemistry such as HO rather than H2O
only got straightened out for good by the middle of the nineteenth century.

However, I do not think that this is exactly the reason for the different names in
the article. Remember that this is written for people who dye cloth, not for chemists.
Therefore the names given there would be the names which suppliers used for
the acids they sold at the time. Even though chemists certainly understood what
nitric acid was by 1877, suppliers probably still sold varying dilutions and grades of
purity under different names, so naturally a recipe for dyeing cloth would use these
names to make it easier for the intended audience to use the recipe. While the
chemists understood what was going on and adjusted their terminology accordingly,
it took a lot longer for people selling industrial chemicals to catch up and start
using names like "10% nitric acid" instead of "double aqua fortis".

To some extent, I encounter a similar situation when I go to the pottery shop
and have to translate chemical names into their terminology in order to be
understood. They even still use some old alchemists' terms like "crocus
martis" and "pearl ash" as well as trade names instead of chemical names.

Rosco Bodine - 24-2-2008 at 17:35

On the first page of this thread a patent US3243385 was posted which described production of stannous nitrate .
It appears that the clear solution described in example 2 of that patent is more likely a solution predominately the +IV
stannic nitrate , and the patents identification is in error .

The figures for quantities and strengths of HNO3 have been compared with the old Dyers Nitrate of Tin , and the preparations are very close to the same . :D:D:D:D

Yes that's four smileys for the +IV valency Tin Nitrate .

more ancient art involving tin nitrate

Rosco Bodine - 29-2-2008 at 14:44

Here is some more interesting observations
recorded in the older literature ,
concerning some anomalously soluble ,
reasonably stable mixtures of tin nitrate with
tin chloride in stable complex mixtures .

It is my belief that Ordways colored mixed valency
compounds of tin , in varying proportions of both
nitrate and chloride is very likely due to the formation
of Pytlewski polymers later described in US3890429
and US3676186 , which are also reported colored materials
yet derived from precursors which are colorless . Also
is probably present some double salt complex with
ammonium chloride and possibly also ammonium nitrate .
These products formed in the reaction system likely account
for the anomalous solubility and stability for the mixture
under lower conditions of acidity than would be allowable
for the individual components to remain in solution .
The anomalously soluble complex mixtures are very possibly sols ,
and not simple solutions . This would also agree with
the hypothesized Pytlewski polymer formation , given that the materials
can be dried (without heating) to a residue and the residue
can be redispersed in plain water to reconstitute the original sol .

At any rate this seems to be a valid method for getting
a highly concentrated , high tin content sol or solution , without
an excessively acidic condition for maintaining the solution or dispersion ,
which may be useful as a precursor for baked SnO2 coatings
that should result from the thermal decomposition and
dehydration and sintering of these compositions .

[Edited on 1-3-2008 by Rosco Bodine]

Attachment: Tin Nitrate stable complex mixtures (1857 Chemical Gazette article).pdf (908kB)
This file has been downloaded 1828 times