making tin(II) chloride
Arsea
into 30% hydrochloric acid. I'm using a flask with a small outlet to
avoid air oxidation of the stannous ion. The solution I make is deep
purple color, however my references say SnCl2 solutions are clear. The
tin and HCl I'm using are industrial grade and so I'm wondering if the
color its due to an impurity that is strongly colored. The reaction
essentially goes to completion, with final HCl concentration < 0.3M.
Zigoteau
Hi, Arsea,
I'm not sure I can answer your question, but I've seen something that
is relevant.
Some time ago I had problems with adhesion to glass slides, and I found
a recipe which worked involving soaking the slides in a "solution" of
SnCl2 in water. So I dissolved SnCl2.2H2O out of the bottle in
ultrapure water. The result was milky white. The referee had handled
SnCl2 before and was unhappy with the description "solution". We split
the difference and called it a colloidal suspension. So clearly at
neutral pH in aqueous solution Sn2+ cross-links to form colloidal
particles of hydrated oxide.
It follows that if the result of dissolving Sn metal in HCl is clear,
the solution must still be quite acidic (low pH).
Books have been written about the optical properties of nanoparticles,
and if the size of the particle is right, and if there is perhaps some
absorbing impurity which couples resonantly to the optical field, you
can get all sorts of colors. Suspensions of silver nanoparticles can be
blue, and suspensions of gold nanoparticles can be various shades of
red. IIRC HCl often has transition metal impurities which might provide
the absorption. Ditto for Sn. SnO2 has quite a high refractive index,
so the diameter of a particle resonant at 500nm might be 100nm or less.
If this idea is correct, then lowering the pH by adding a bit more acid
should decolorize it. Does it?
Cheers,
Zigoteau.
rek...@gmail.com
hydroxide, and water complexes. According to a believable entry in
wikipedia, SnCl2(aq) can be oxidized on contact with air to form
aqueous SnCl4 and various Sn(OH)Cl solids. If your solution was the
slightest bit opaque, you have formation of solids.
Farooq W
> I was trying to make stannous chloride, by dissolving excess tin metal
> into 30% hydrochloric acid. I'm using a flask with a small outlet to
> avoid air oxidation of the stannous ion. The solution I make is deep
> purple color, however my references say SnCl2 solutions are clear.
Solid SnCl2 is colorless. Stannous chloride produces a milky white
suspension in water due to its hydrolysis, however strongly acidic
solution should be colorless.
>The
> tin and HCl I'm using are industrial grade and so I'm wondering if the
> color its due to an impurity that is strongly colored.
Though this might be relevant in you case (one can safely rule out gold
in your HCl), but traces of gold produce a purple coloration in the
presence of stannous and stannic chloride. I don't recall any metal
imparting a deep purple color if present in traces in HCl solution
(cobalt could be the culprit if large amount is present); Is the tin
metal pure? But finally Zigoteau's hypothesis on the color due to
colloidal particles sounds correct.
>The reaction
> essentially goes to completion, with final HCl concentration < 0.3M.
Does adding concentrated nitric acid remove the blue coloration? If you
are inclined to further experimenting, try adding some activated
charcoal to check if the purple color is removed/altered on shaking.
Wilco Oelen
Did the tin dissolve easily? I have some very pure tin (99.99+ %) which
does not dissolve at all in concentrated HCl at any appreciable rate.
If the tin contains some titanium, then you can get a very deep purple
color. I have dissolved titanium metal in 30% HCl and the color of
Ti(3+) ions is remarkably strong, even mg quantities of titanium give
strong purple colors. However, having titanium as an impurity in tin
metal seems a little odd to me.
I also have SnCl2 and indeed, this dissolves very easily in water, but
also in SnCl2, giving colorless solutions in strong acid and milky
white solutions (due to extensive hydrolysis) in slightly acidic to
neutral solutions. As far as I know, Sn(IV) species also are colorless.
Bob M
Are you sure you have tin granules Sn and not Titanium granules Ti.
Titanium dissolves in HCl to produce deep purple TiCl3.
Bob M
number6
Wilco Oelen wrote:
> Did the tin dissolve easily? I have some very pure tin (99.99+ %) which
> does not dissolve at all in concentrated HCl at any appreciable rate.
Very true ... Adding a little H2O2 is needed to dissolve it in
reasonable times ...
>
> If the tin contains some titanium, then you can get a very deep purple
> color. I have dissolved titanium metal in 30% HCl and the color of
> Ti(3+) ions is remarkably strong, even mg quantities of titanium give
> strong purple colors. However, having titanium as an impurity in tin
> metal seems a little odd to me.
>
You may have hit on it ... much easier to believe that Ti was mistaken
for Tin than there is some weird chemistry going on ... Like you I
doubt the impurity ... though TiCl4 and SnCl4 are used in conjunction
in some hot end coating for glass ...
Arsea
>
> Wilco Oelen wrote:
>
>
>
>>Did the tin dissolve easily? I have some very pure tin (99.99+ %) which
>>does not dissolve at all in concentrated HCl at any appreciable rate.
>
>
> Very true ... Adding a little H2O2 is needed to dissolve it in
> reasonable times ...
the reaction is very slow, and the coin sized chunks of tin metal I used
tool about 3 days at 75°C in concentrated HCl to dissolve in reasonable
amounts. Even after this time there is still about 3M of acid remaining.
Thanks for the H2O2 tip, for next time.
>
>>If the tin contains some titanium, then you can get a very deep purple
>>color. I have dissolved titanium metal in 30% HCl and the color of
>>Ti(3+) ions is remarkably strong, even mg quantities of titanium give
>>strong purple colors. However, having titanium as an impurity in tin
>>metal seems a little odd to me.
>>
>
>
> You may have hit on it ... much easier to believe that Ti was mistaken
> for Tin than there is some weird chemistry going on ... Like you I
> doubt the impurity ... though TiCl4 and SnCl4 are used in conjunction
> in some hot end coating for glass ...
>
If it is dissolved TiCl3 then is it possible to be a contaminate in the
HCl I'm using ? Say the HCl had been stored or processed in titanium
equipment. Maybe some TiCl4 (yellow) impurity is reduced to TiCl3
(purple) when the tin metal was added. My reason to believe the
contaminant to be in the acid was because the purple color appeared just
minutes after the addition of tin metal, and did not increase in color
concentration, even after the three days. If Ti were from the tin then
I'd expect the purple to increase as the reaction proceeded, but this
wasn't the case.
The metal appeared to be tin. I actually casted the metal chucks myself
and it melted at the expected temperature. The metal also made the
crystalline cracking sound as it was deformed and bent. Only tin posses
this property.
unknown
A far more interesting compound is stannous acetate readily made by
dissolving tin in glacial acetic acid.
Stannous acetate is used by the CIA as a poison that once ingested
mimics a heart attack; obviously used for untraceable assassinations.