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blogfast25
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Yep, they're the real mccoy alright. Nice one!
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Lambda-Eyde
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AgCl dissolves in aqeous ammonia, giving the diaminesilver ion Ag(NH<sub>3</sub> <sub>2</sub><sup>+</sup>. Reduction with e.g. glucose will make a shiny silver mirror, which can be dissolved with
HNO<sub>3</sub>. 
[Edited on 28-3-2011 by Lambda-Eyde]
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Arthur Dent
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Question... Will it work if I use a vacuum buschner funnel with a #40 wheaton paper filter to dry off the precipitate?
Or will the precipitate clog-up the paper and the filter's frit? The idea of using ammonia to dissolve the AgCl is appealing. Perhaps I could use
ammonia to wash-off the funnel after the filtering?
Robert
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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blogfast25
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Robert, gentle vacuum assisted filtering should work well. You will always lose some product on the filter, of course. Just make sure the precipitate
is still wettish. If your filter is glass frit (but it doesn't appear to be) , you can carry out the acid digestion on it, it also resists heat well,
so no probs there…
Getting rid of most of the water should greatly help dissolving it into 68 % HNO3. Remember: you’re displacing a very strong acid! Heat will also
help: it will drive off the HCl, thereby pushing the equilibrium to the right (mass balance effect). Heat may also decompose the AgCl a bit, giving
the nitric something to get its teeth into (the silver). Multiply pronged attack!
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Arthur Dent
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Quote: Originally posted by blogfast25  | Robert, gentle vacuum assisted filtering should work well. You will always lose some product on the filter, of course. Just make sure the precipitate
is still wettish. If your filter is glass frit (but it doesn't appear to be) , you can carry out the acid digestion on it, it also resists heat well,
so no probs there…
Getting rid of most of the water should greatly help dissolving it into 68 % HNO3. Remember: you’re displacing a very strong acid! Heat will also
help: it will drive off the HCl, thereby pushing the equilibrium to the right (mass balance effect). Heat may also decompose the AgCl a bit, giving
the nitric something to get its teeth into (the silver). Multiply pronged attack!
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I'll use a Kimax 30ml buschner funnel with a glass frit and drop a small 1" circle of wheaton filter paper on top. I'll use Nitric to clean the
funnel.
Okay, i'll let the precipitate dry off a bit, as for the reaction, i'll do it outside because the mean orange clouds scare me  . When the reaction stabilizes, i'll bring it back inside and put it on my hotplate
for a bit. I have a sheet of pure Silver (JT Baker) so I might drop a tiny piece to see if it dissolves equally fast in the Nitric
Robert
[Edited on 28-3-2011 by Arthur Dent]
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blogfast25
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The ‘mean orange clouds’ should scare you:
Ag === > Ag+ + e
NO3(-) + 4 H+ + 3e ==== > NO + 2 H2O
Or: Ag + 1/3 HNO3(-) + H+ === > Ag+ + 1/3 NO + 2/3 H2O
(In reality the reaction products tend to depend on concentration, temperature and type of metal but NO is always part of the mix)
NO reacts immediately with air oxygen to NO2 (NO + ½ O2 === > NO2), which is the amber/reddish/brown gas (nitrogen dioxide) you’re seeing. It
smells somewhat like chlorine and is probably even more toxic. Stay AWAY!
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cyanureeves
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i always get that stubborn bluish grey stuff whenever i dissolve sterling and it turns out to be silver. it never re- dissolves for me.i've never
gotten a clear solution like nurd rage because sterling is an alloy. i bet you can turn it to metal silver with hcl acid and zinc.
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hkparker
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I have a tin(II) chloride question...
I've been trying to make some recently, and I was a little impatient with the reaction between my tin and HCl (I was using tin about the size of coins
and hadn't heated my HCl. I'm probably going to just reflux is next time. Easier then what I just tried to do...). But before I do that i'm curious
as to what happened with my last attempt...
Being impatient with my HCl I decided to add some HNO3 to get things going, and I figured I would convert any tin nitrate to chloride later. I also
knew this would make tin(IV) but planned on reducing that. I was left with a clear solution of what I believe was SnCl6(2-). Jor told me that HNO3
oxidizes Sn to SnO2, but the presence of Cl- would likely form that complex. I neutralized with Na2CO3 and sure enough got a bright white precipitate
of (what I believe was) SnO2. So I dissolved that in HCl to yield SnCl4. Heres where things got tricky. As I wanted to make SnCl2, I though adding
tin metal would reduce it back. SnCl4 + 2Sn --> 2SnCl2. However immediately after adding the tin metal the solution went from clear, to yellow,
to green, to a brown. Any Idea on what this brown solution could be?
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blogfast25
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No but I'm guessing higher acid strength will remediate the problem...
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Waffles SS
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I have a another idea to produce Tin(II)chloride:
First Tin(II) sulfate by adding tin metal to copper sulfate solution:
Sn (s) + CuSO4 (aq) → Cu (s) + SnSO4 (aq)
Second adding calsium chloride to tin sulfate solution:
SnSO4 (aq) + CaCl2(aq) → SnCl2 (aq) + CaSO4(S)
Then filter calsium sulfate and evoparation solution under vaccum.
[Edited on 27-6-2011 by Waffles SS]
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blogfast25
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Calcium sulphate is a little soluble in water. Your tin sulphate is likely to be contaminated with CaSO4...
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LanthanumK
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For some reason, leaded solder dissolves in cold HCl in about 1 week, leaving behind a generous serving of lead powder, which burns nice and toxic.
Unleaded solder(?) hardly dissolves at all in HCl, so I just gave up attacking it that way.
hibernating...
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Neil
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| Quote: Originally posted by Arthur Dent |
Oh, and my acid is indeed 32% HCl and the solder according to the Kester data sheet specifies the correct amount of tin vs silver. Could it be that
the rosin core of the tin solder affects the reaction? I have not the ghost of an idea what is in that rosin, only that it says it's not recommended
for electronics.
Robert
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Rosin is in "Rosin" - Rosin comes from tree sap.
If it is not electrical solder but rather metal working solder then instead of Rosin it has an acid flux which can be anything from urea to metal
chlorides. If you want to find a specific one you often have to find the manufacture and search for an MSDS for "Acid Core" not "acid cored solder"
they tend to list the core as "acid core" on the solder msds sheets.
In dissolving thick tin plumbing solder I found that winding long tight coils and tossing those in a old HCl bottle with some 32% let the reaction
self heat and proceed easily, dissolving the whole roll in three days/three additions.
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Waffles SS
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Yes,
Barium sulfate is insoluble in water and we can use barium chloride instead of calcium chloride:
BaCl2(aq) + SnSO4(aq) = SnCl2(aq) + BaSO4(s)
Solubility of barium sulfate in water is 0.0002448 g/100 mL (20 °C)
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Waffles SS
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Difficult part is evoporation step tin(II)chloride decompose in hot water(hydrolyse occur even in more water)
Tin(II) chloride can dissolve in less than its own mass of water without apparent decomposition, but as the solution is diluted hydrolysis occurs to
form an insoluble basic salt:
SnCl2 (aq) + H2O (l) is in equilibrium with Sn(OH)Cl (s) + HCl (aq)
What we can do for evoporation step?evoporation under vaccum seems not work very well.How we can evoporate water?
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Lambda-Eyde
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Use an excess of HCl?
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blogfast25
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Carefully and by having an acid reserve at all times. It's a bit hit and miss for small quantities: boil in too far and you get hydrolysis, not far
enough and the product doesn't crystallise. But it worked for me.
The product of hydrolysis is still soluble in HCl though, so you can try over and over again untill you get it right.
Worth exploring would be what works with some chloride hydrates like AlCl3.6H2O or ZrOCl2.8H2O: saturate a cold solution of the salt with HCl (gassing
with HCl gas), the hydrate then precipitates.
[Edited on 5-7-2011 by blogfast25]
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Waffles SS
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"Have an acid reserve at all time" is good idea but how we can get rid of HCl in final crystal?(by this method there is some HCl in final product)I
think we should wash final product with suitable solvent ,Tin(II) chloride is soluble in ethanol and acetone(i think we should find another solvent
for this purpose)also it think result of tin chloride hydrolysis will react with HCl again and make to Tin(II) Chloride again
Sn(OH)Cl + HCl =SnCl2 +H2O
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LanthanumK
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When evaporated, much of the HCl fumes away, decreasing the amount of HCl in the final crystals.
hibernating...
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Arthur Dent
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It's pretty hard to crystallize. And it is indeed necessary to have an excess of HCl in the solution. Half of my yield was simply "evaporated" at room
temperature on a glass plate, giving the familiar needle-like structures, that can be scraped-off the glass and be put in a vial (still a bit moist
and acidic). That is a long process but it did yield fairly pure, snow white powder that dissolves readily in very little water.
The other half, I tried to evaporate on a hotplate at low temp with the solution in a large petri dish. Most of it hydrolized quite rapidly, turning
into a beige, opaque, insoluble paste, but around the rim, a cluster of perfectly square crystals formed (about 1 to 1.5mm) which I salvaged one by
one with tweezers and washed with acetone. They look like perfectly clear pickling salt crystals. This second method is hardly recommended because a
lot of the Tin Chloride was lost when it hydrolized into tin hydroxide and oxide (about 75%).
Robert
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Waffles SS
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@Arthur dent(robert),
Tin chloride solution even is sensitive to air :
| Quote: |
Solutions of SnCl2 are also unstable towards oxidation by the air:
6 SnCl2 (aq) + O2 (g) + 2 H2O (l) → 2 SnCl4 (aq) + 4 Sn(OH)Cl (s)
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"Evaporating at room temperature on a glass plate" means long time contact with air oxygen!.(are you sue you didnt get Tn(IV)chloride?)
also you "tried to evaporate on a hotplate at low temp with the solution in a large petri dish" and you saw insoluble paste?i think this is impossible
to see insoluble paste if your solution contain HCl
Sn(OH)Cl (s) + HCl(aq) = SnCl2(aq) +H2O(l)
I am interested to try evaporation under vaccum but i am not sure the result change
[Edited on 6-7-2011 by Waffles SS]
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LanthanumK
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Last time I tried evaporating a SnCl2 solution, I experienced the sentence of doom for heatless evaporation: deliquescence.
hibernating...
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Cloner
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| Quote: Originally posted by Waffles SS | @Arthur dent(robert),
Tin chloride solution even is sensitive to air :
"Evaporating at room temperature on a glass plate" means long time contact with air oxygen!.(are you sue you didnt get Tn(IV)chloride?)
[Edited on 6-7-2011 by Waffles SS] |
The boiling point of tin(IV)chloride is much lower than the BP of tin(II)chloride so this is likely. For the purpose of making conducting glass this
is fine though.
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Waffles SS
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| Quote: Originally posted by Cloner | | Quote: Originally posted by Waffles SS | @Arthur dent(robert),
Tin chloride solution even is sensitive to air :
"Evaporating at room temperature on a glass plate" means long time contact with air oxygen!.(are you sue you didnt get Tn(IV)chloride?)
[Edited on 6-7-2011 by Waffles SS] |
The boiling point of tin(IV)chloride is much lower than the BP of tin(II)chloride so this is likely. For the purpose of making conducting glass this
is fine though. |
@cloner thanks but:
Boiling point of tin(IV) chloride is 114c
We are talking about evaporation" at room temperature" not at boiling point of water!
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Arthur Dent
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Yes indeed it is sensitive to air, but much more in the absence of HCl, so as I mentioned, my solution is evaporated at room temperature until the
formation of needles (still moist a bit) and harvested at that moment. The vial containing the Tin Chloride has the characteristic smell of HCl.
As for the insoluble paste, when the HCl is completely driven off by heat, the Tin Chloride undergoes hydrolysis and becomes a cruddy, insoluble goop
that leaves a precipitate in water. I've tried it twice and got the exact same results, under heat, a small quantity of the solution reaches the
ideal level of water/HCl and turns into lovely, big crystals, the rest turns into tin hydroxide and oxide.
Robert
[Edited on 6-7-2011 by Arthur Dent]
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