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Author: Subject: Small scale preparation of Stannic Chloride (SnCl4)
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[*] posted on 8-2-2014 at 11:55
Small scale preparation of Stannic Chloride (SnCl4)


Stannic chloride, SnCl4 or tin (IV) chloride can be prepared by direct union of chlorine with tin or tin (II) chloride but another potentially less hazardous route is also available: pyrolysis of K2SnCl6, a chemical that is quite easy to prepare with simple wet chemistry.

Pyrolysis: K2SnCl6(s) === > 2 KCl(s) + SnCl4(g)

To test this method some 7 g of K2SnCl6 I had prepared much earlier was first thoroughly dried on an electrical hot plate. Wet starting material would lead to hydrolysis of the SnCl4.

A rather ad hoc, erm… apparatus was cobbled together from a broken burette, bent in the middle, a one hole rubber stopper and a test tube:



Through the one hole rubber stopper was inserted a K-type thermocouple (- 50 to 1300 C) and the reagent then carefully loaded into the bent burette via the condenser side of the apparatus.

With everything in place, medium high Bunsen heat was then applied evenly over the ‘reactor’ area, avoiding of course heating the rubber stopper. From about 450 C I could start to see some change in the potassium hexachlorostannate and fumes were forming. As these seemed to have trouble condensing in the condenser area (the BP of SnCl4 is about 118 C), the test tube was immersed into water and ice in a 250 ml beaker.

Cranking up the heat to well over 500 C, droplets of a clear liquid could then be seen to come off the end of the condenser tube and into the test tube. The K2SnCl6 could also been seen to be bubbling, presumably due to boiling SnCl4. All in all about 0.5 ml of SnCl4 were thus collected before the burette had predictably started to sag and the rubber stopper was also starting to pass dense smoke (SnCl4).

The product, slightly turbid due to entrained salt, I think:



On stopping the heating, SnCl4 continued to evolve for some time, as can be seen here:



And this is the reactor area at the end of the experiment:



A better set up (the 'reactor' was definitely over-filled), possibly involving also mild vacuum, should allow to distil off good quality SnCl4 without too many problems.


[Edited on 8-2-2014 by blogfast25]




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[*] posted on 8-2-2014 at 14:59


Nice work. A couple questions:
How did you prepare your K2SnCl6?
Is there any reason that this experiment cant be conducted in a metal can? That would allow a much bigger batch size, and possibly the attachment of a proper condensor.




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[*] posted on 8-2-2014 at 15:57


@oscillator:

There are various ways of preparing K2SnCl6. My preferred one comes down to dissolving tin in a mixture of HCl and HNO3 (Aqua Regia), which results essentially in a H2SnCl6 solution. Solid KCl is then added and the slurry heated until all KCl dissolves. On cooling the salt crystallises out (it has a strong solubility-temperature dependency). I’ve posted the precise procedure somewhere here on this forum. One can also start from SnCl2 and Aqua Regia or SnCl2 dissolved in HCl, then cold oxidised with H2O2.

Re. metal containers, copper would be the best I think. I fear possible reduction of the Sn(IV) to (II) or even (0) with more reactive metals.

Having retrieved the thermocouple, I can see it was attacked quite badly. A fairly thick black-grey residue has coated it and is hard to remove. I wonder if some reduction of the SnCl4 by iron in the steel sheathing has occurred, so that the coating may contain elemental tin. The residue left inside the reactor is also very dark, unlike KCl.


[Edited on 9-2-2014 by blogfast25]




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[*] posted on 9-2-2014 at 09:23


With the reactor cold I took another look at it today and what struck me were some prominent purple patches, here’s one:



The photo doesn’t really do it justice, in reality it’s more like a lighter shade of chromium alum.

Most of it was recovered by cracking the glass. The material was then subjected to simmering 37 % HCl in which it dissolved quickly (bar a bit of insoluble residue) to a green solution.

After filtering this was alkalised with strong KOH upon which a non-descript fluffy precipitate appeared. On adding a bit of 35 % H2O2 this turned brown-red, clearly Fe(OH)3.nH2O.

Figuring there might also be chromium in there (I think the thermocouple’s sheath is SS but I’m not sure) this slurry was filtered. pH of the filtrate was about 10. Lead (II) nitrate solution was added but no yellow PbCrO4 showed up, only a bit of PbCl2.

So it appears that some of the SnCl4 had reacted with the Fe in the thermocouple’s sheath but that only anhydrous Fe (II) chloride (and not FeCl3) had formed, presumably acc.

SnCl4 + 2 Fe === > Sn + 2 FeCl2

The Sn could explain the dark colour of most of the post reaction mix. Yesterday I fumed off quite a bit of mixed (K,Rb) hexachlorostannates in a nickel crucible and almost no discolouration of the mix was observed.

But none of this really explains the unusual purple colour of the material...

[Edited on 9-2-2014 by blogfast25]




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[*] posted on 9-2-2014 at 13:43


Well I have a bar of tin solder trying to dissolve in hydrochloric acid. So far all its done is make the solution go purple, sort of like a manganese(II) ion. Maybe thats what is causing the residue to be purple. What is your source of tin?



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[*] posted on 9-2-2014 at 14:33


Quote: Originally posted by Oscilllator  
Well I have a bar of tin solder trying to dissolve in hydrochloric acid. So far all its done is make the solution go purple, sort of like a manganese(II) ion. Maybe thats what is causing the residue to be purple. What is your source of tin?


I use different sources of tin, from 99 % metal to pewter to SnCl2 to SnO2. But I always recrystallize the K2SnCl2, using its poor solubility in 37 % HCl. Purity isn't an issue here, I think.

Contamination is unlikely to explain why so much iron was found in the post-reaction mix and why the probe seemed badly attacked.

Solvated Mn(II) is very light coloured (pink), it needs to be quite concentrated to show its colour.

Dissolving tin metal in HCl is a slow process, even at BP. Try adding some nitric acid. Tin dissolves in HCl/HNO3 mixtures very quickly.

[Edited on 9-2-2014 by blogfast25]




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[*] posted on 16-4-2014 at 14:10


I think about a reaction:

SnCl4 (anhydrous, liquid)(excess) + 2CuSO4(anhydrous) --> Sn(SO4)2 (solid) + CuCl2 (solid)

I need to know how fast this should react (I assume that will be a fast reaction) , because I need a good quantity of Sn(SO4)2 (anhydrous) in a short amount of time.
And... I need Sn(SO4)2 because I want to make SO3 by pyrolysing it at just 250 *C ... I hope that is a good way to obtain SO3 fast...
Sn(SO4)2 --(200-250*C)--> SnO2 +2SO3

Can somebody help me with this? (I mix those 2 together, and I can observe just that the mixture become more viscous...nothing else)
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[*] posted on 17-4-2014 at 04:33


Quote: Originally posted by DoctorZET  
Can somebody help me with this? (I mix those 2 together, and I can observe just that the mixture become more viscous...nothing else)


There simply is no reason why your proposed reaction should proceed, as far as I can see. And how would you suggest to separate the reaction products?




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[*] posted on 22-4-2014 at 05:27


Well... I think about this reaction, because I know that when 2 salts are melted together, a reaction should proceed (the most reactive metal ions in the mixture associate with the most electronegative ions)...and I have here one of the salts well melted (SnCl4) and the another salt (CuSO4) in a very fine powder form.

And about the mixture... I don't need to separate this 2 chemicals... I just need to remove the unreacted CCl4 and to heat the remaining salt up to 200-250*C to make SO3 (CuCl2 & SnO2 will remain in the crucible).

So, in order to make some SO3, I need to know if Sn(SO4)2 is really produced.
I think that I should make a small scale experiment (with max. 1g of CuSO4) to see if there will be produced any white fumes of SO3 .
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[*] posted on 22-4-2014 at 05:40


I think he is hoping when he fuses the salts he, the heat will drive the equilibrium to the right, as SO3 is liberated, SnCl4 + 2 CuSO4 ↔ 2 CuCl2 + SnO2 + 2 SO3. Fused salts tend to ionize, which will allow the reaction to proceed - I think.

This is simply speculation though, it seems like it could work...
It's hardly relevant to this topic, perhaps it should be split off?

[Edited on 22-4-2014 by Zyklonb]




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[*] posted on 22-4-2014 at 06:42


Soo..... this may seem a bit obvious but reacting this with Al will give anhydrous AlCl3 and Sn right? That seems to be a rather easy process. Maybe I should try it sometime.



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[*] posted on 22-4-2014 at 12:50


Quote: Originally posted by DoctorZET  
Well... I think about this reaction, because I know that when 2 salts are melted together, a reaction should proceed (the most reactive metal ions in the mixture associate with the most electronegative ions)...



Nope. You're quite confused, aren't you? Ions can't be electronegative: only atoms are. Chlorine atoms are electronegative because they have a tendency to absorb an electron and thus become an anion (the chloride ion).

SnCl<sub>4</sub> cannot be remotely described as a 'salt', it's an almost purely covalent compound (hence its low melting point). Trust me, it's not a 'molten salt'.

Tn IV sulphate can only exist in strong sulphuric acid solution, not as a solid.


[Edited on 22-4-2014 by blogfast25]




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[*] posted on 22-4-2014 at 12:53


Quote: Originally posted by bismuthate  
Soo..... this may seem a bit obvious but reacting this with Al will give anhydrous AlCl3 and Sn right? That seems to be a rather easy process. Maybe I should try it sometime.


Almost certainly and almost certainly quite violently. You need to calculate the reaction's change in enthalpy to be sure.

Much better is the recently tested reduction of ZnCl2 with Al powder:

http://www.sciencemadness.org/talk/viewthread.php?tid=30150#...




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[*] posted on 25-4-2014 at 06:26


Yes, Blogfast25 is right :

Tin tetrachloride is almost a covalent compound, meanwhile the copper sulfate is almost an ionic salt ... so, they can't react in a fast, visible way (even if one of them is liquid).

In order to run this reaction, some substances (able to create ions and act as solvents for both ionic and covalent compounds) are needed .
I found a similar situation in an old chemistry book (Authors: Mihai Zapan and Edith Beral ; The printing company "OLTENIA" - 1977 - S.R.Romania) that, if I simply fuse a covalent semimetal halide with an ionic salt of an unreactive metal, it's possible that nothing to happen ... But, if I do this with by adding a strong acid with traces of water (a few drops of acid - just a catalytic amount) in a solution of 2,2,2-trichloroethanol ... a reaction will take place as soon as the strong acid ionize the semimetal bonds.

I suppose the best acid that I can use in my reaction is conc. H2SO4 .
This is how the reaction should work (in trichloroethanol solution) :
2H2SO4 + SnCl4 --> Sn(SO4)2 + 4 HCl
4HCl + 2CuSO4 --> 2CuCl2 + 2H2SO4
The H2SO4 is always regenerated, so, hydrogen acts as a ion transporter from Cu to Sn and in reverse.
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[*] posted on 25-4-2014 at 08:14


Well, good luck with all that DZET, especially with:

'4HCl + 2CuSO4 --> 2CuCl2 + 2H2SO4'... which appears impossible to me as this would be the displacement of a strong acid (H2SO4) by a slightly weaker one (HCl). But in 2,2,2-trichloroethanol... who knows?

All this trouble for a bit of SO3? There are threads on SO3 generation that you might want to consult before resorting to... 2,2,2-trichloroethanol.

Also, if you believe 2H2SO4 + SnCl4 --> Sn(SO4)2 + 4 HCl will proceed, then what's the point to the CuSO4? Just try anhydrous SnCl4 and 98 % H2SO4.


[Edited on 25-4-2014 by blogfast25]




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[*] posted on 25-4-2014 at 13:05


Quote: Originally posted by blogfast25  

'4HCl + 2CuSO4 --> 2CuCl2 + 2H2SO4'... which appears impossible to me as this would be the displacement of a strong acid (H2SO4) by a slightly weaker one (HCl).

HCl(aq) is a stronger acid than H2SO4.

Quote:
Originally posted on Wikipedia
This is a list of strong acids with pKa < -1.74, which is stronger than hydronium ion, from strongest to weakest.

Hydroiodic acid HI (pKa = −9.3)[1]
Hydrobromic acid HBr (pKa = −8.7)[1]
Perchloric acid HClO4 (pKa ≈ −8)[3]
Hydrochloric acid HCl (pKa = −6.3)[1]
Sulfuric acid H2SO4 (first dissociation only, pKa1 ≈ −3)[3]
p-Toluenesulfonic acid (pKa = −2.8) Organic soluble strong acid
Methanesulfonic acid (pKa = −1.92) Liquid organic strong acid[4]


[EDIT] Just in case:A more reliable source then Wiki


[Edited on 25-4-2014 by Zyklonb]




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[*] posted on 25-4-2014 at 13:41


Quote: Originally posted by Zyklonb  
Quote: Originally posted by blogfast25  

'4HCl + 2CuSO4 --> 2CuCl2 + 2H2SO4'... which appears impossible to me as this would be the displacement of a strong acid (H2SO4) by a slightly weaker one (HCl).

HCl(aq) is a stronger acid than H2SO4.


Hardly relevant in aqueous solution- they're both strong, and neither the chloride ion nor the hydrogen sulphate ion are going to accept a proton when there is water around. In a different solvent, aqueous pKa values don't work.

The above reaction isn't going to happen. The reverse reaction could be made to occur at higher temperatures, as HCl is much more volatile than sulphuric acid.




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[*] posted on 26-4-2014 at 03:52


I also think '2H2SO4 + SnCl4 --> Sn(SO4)2 + 4 HCl' isn't going to work. See for instance the lab preparation of PbCl4 by pouring cold conc. H2SO4 on (NH<sub>4</sub>;)<sub>2</sub>PbCl<sub>6</sub>. The resulting PbCl4 doesn't decompose to Pb(SO4)2.



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[*] posted on 12-6-2015 at 17:40


I just finished an attempted preparation of tin(IV) chloride by passing dry chlorine gas over some strips of tin metal. It was working fairly well for a while, and the tin was readily dissolving, forming a clear liquid. After a while though, the liquid began to turn greensish yellow, and little brown flecks were floating around in it. I have no idea why it would do this. My first thought was that the green color was from chlorine being dissolved in it, but I don't know what the brown stuff is. I suppose some impurities could have come from the PVC tubing that I used, but that doesn't seem very likely to me. The first picture shows the setup, minus the suckback trap that I added prior to starting. The second picture shows the suckback trap. The third picture shows the product after it began to turn yellow. Any ideas about why this happened?

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[*] posted on 12-6-2015 at 18:01


I can only put this down to impurities in the tin.

Most of the tin salts I've prepared came from a single, large old tankard of pewter. Tin (and pewter) dissolves well in aqua regia. The obtained solutions were always greenish. I later managed to prove w/o reasobable doubt that the green colour was due to chromium (III)! Unexpected but true.

I'm not claiming this is the case in your experiment (I don't know that) but maybe it's food for thought? Do you have any idea about the purity of the tin you used?

Yellow could point to FeCl3 of course because chlorine will oxidise Fe to Fe(III).

Nice experiment, BTW!


[Edited on 13-6-2015 by blogfast25]




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[*] posted on 12-6-2015 at 18:41


Quote: Originally posted by blogfast25  
I can only put this down to impurities in the tin.

Most of the tin salts I've prepared came from a single, large old tankard of pewter. Tin (and pewter) dissolves well in aqua regia. The obtained solutions were always greenish. I later managed to prove w/o reasobable doubt that the green colour was due to chromium (III)! Unexpected but true.

I'm not claiming this is the case in your experiment (I don't know that) but maybe it's food for thought? Do you have any idea about the purity of the tin you used?

Yellow could point to FeCl3 of course because chlorine will oxidise Fe to Fe(III).

Nice experiment, BTW!
Thank you! That's interesting about the chromium, but the tin that I used should be pure. I got it from my school chemistry lab, and I am pretty sure that it is reagent grade. I will try to test it with ferrocyanide to see if there's any iron in it though. It's just so annoying to work with because of the fuming.



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[*] posted on 13-6-2015 at 06:21


Dissolve a bit in Fe-free 33 % HCl (that's quite slow), then test the SnCl2 solution for Fe?



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[*] posted on 15-7-2015 at 08:45


I got around to distilling my cloudy SnCl4 today, with great success. I set up a tiny simple distillation apparatus and sucked the air out through the drying tube and let it back in slowly to try and keep the inside of the apparatus dry. Then I brought it to a boil and let it distill until there was only residue left in the boiling flask. The thermometer read a consistent 115°C once it began distilling. It is still slightly yellowish, and I think this is likely from dissolved chlorine, but it is crystal clear, and I'm satisfied with it.

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[*] posted on 16-7-2015 at 07:40


Some commentary from Atomistry.com on SnCl4 (link: http://tin.atomistry.com/stannic_chloride.html ):

"Stannic Chloride, SnCl4, was first obtained by Libavius in 1597, by distilling tin or its amalgam with corrosive sublimate. The liquid thus obtained was termed by its discoverer Spiritus argenti vivi sublimati, but was known later as Spiritus fumans Libavii. It was known in 1630 that a solution of this substance is obtained by dissolving tin in aqua regia; and a Dutchman named Drebbel discovered its value as a mordant by which the red colour of cochineal could be fixed on cloth.
Stannic chloride is prepared by passing chlorine over molten tin, tin-foil, or tin-plate. This latter method is employed in the recovery of tin from tin-plate waste. The reaction is accompanied by the evolution of heat and sometimes light. Stannic chloride may also be obtained by the action of chlorine on stannous chloride, and of chloroform vapour on the dioxide. It is reduced to stannous chloride by hydrogen at 1000° C. Stannic chloride is a thin, colourless, fuming liquid, which has a density of 2.2788 at 0° C.; it forms a solid at low temperature, which melts at -33° C., and boils at 113.89° C. under 750 mm., 114° C. under 757 mm., and 114.1° C. under 760 mm. pressure; its critical temperature is 318.7° C., and critical pressure 36.95 atmospheres. The heat of formation of this compound from its elements is 127,250 calories, and its heat of solution 29,920 calories. Stannic chloride dissolves very readily in alcohol, and also in ether, benzene, nitrobenzene, carbon disulphide, and other non-hydroxylic solvents."

And also of possible interest:

"Stannic chloride dissolves rhombic sulphur, yellow phosphorus, iodine, and bromine; with turpentine it reacts so vigorously as to set fire to the hydrocarbon. It dissolves in excess of water, forming a solution from which hydrated stannic oxide separates on boiling or standing; nevertheless it unites with smaller quantities of water to form various solid hydrates. Demarchy first observed, in 1770, that with one-third of its weight of water stannic chloride forms a crystalline mass, which was termed butter of tin. This hydrate, which is SnCl4.5H2O, is also called oxymuriate of tin. Hydrates containing 3, 4, and 8 molecules of water are also known, and separate from the saturated solution with which they are in equilibrium at different temperatures. Thus SnCl4.8H2O is stable up to 19° C., SnCl4.5H2O is stable between 19° C. and 56° C., SnCl4.4H2O between 56° C. and 63° C., SnCl4.3H2O between 63° C. and 83° C.
The hydrolysis of stannic chloride proceeds according to the scheme:
SnCl4 + 4H2O ⇔ Sn(OH)4 + 4HCl;
the extent of the hydrolysis as measured by conductivity when equilibrium is attained depends, not only on the amount of water present, but also on the temperature, the raising of the temperature tending to retard the hydrolysis which is an exothermic reaction. The combination of stannic chloride with hydrochloric acid to form H2SnCl6 also influences the equilibrium, ..."

When I amass a sufficient quantity of pure tin, I believe my 1st effort, on a small scale preparation, would be by the action of Cl2 (pretty sure dried chlorine gas is required else Cl2 + H2O could form HCl, for example, creating some SnCl2 also, see http://tin.atomistry.com/stannous_chloride.html , and also HOCl, possibly (?) creating an unstable Sn(ClO)4, ... , and please see http://www.sciencemadness.org/talk/viewthread.php?tid=9713 as a possible prep guide for dry chlorine gas) on molten tin (which melts 100 C lower than Pb, at just over twice the boiling point of water at 232 C). The action apparently puts on a light display, and being exothermic, would reduce the amount of applied heat. In fact, here is Wikipedia's comment on SnCl4 preparation (see https://en.m.wikipedia.org/wiki/SnCl4 ) to quote:

"It is prepared from reaction of chlorine gas with elemental tin at 115 °C (239 °F).

Sn + 2 Cl2 → SnCl4 "

I would then cool the vapors, collect and especially protect the SnCl4 from moisture.

For example, a very simple embodiment would be to place 11.9 grams of Sn (around 2 cc) in the corner of a large vessel containing at least 4.5 liters of dry chlorine gas. Place ice bags around the top of the vessel and mildly heat the corner with the Sn powder until melted and then as needed. Expected yield 0.1 mole or 2.6 grams (only about 1 cc ) of SnCl4. Repeat procedure until desired amount of SnCl4 has been formed. Note, incidental SnCl2 from unremoved water could further react with Cl2 to form SnCl4.

Also, interestingly, the method via aqua regia is nearly 400 years old!

Blogfast is right on the hazards of large scale chlorine generation, which the above (to be successfully demonstrated) preparation hopefully mitigates by iterative generation of controlled portions of Cl2 instead.

[Edited on 16-7-2015 by AJKOER]
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[*] posted on 13-4-2016 at 18:41


Preparation of Anhydrous SnCl4

Introduction & Background
Wanting some anhydrous SnCl4 I searched my usual sources but no one had it. Looking in Brauer (ref 1) I found a facile synthesis using tin shot and Cl2. In perusing the posts above in this thread I noted that zts16 had successfully made SnCl4 using a similar procedure. Having both of the needed reagents at hand I decided to make it myself.

Sn + 2Cl2 --> SnCl4 + heat

SnCl4 Characteristics
SnCl4 is a covalently bonded Lewis acid, soluble in many organic solvents. It is a liquid with specific gravity of 2.23 and has a boiling point of 114°C (Wiki). In its anhydrous form it is aggressively hygroscopic, fuming in humid air to form a pentahydrate. This, then, IMO, is a very interesting reagent.

Procedure
For this preparation I followed Brauer’s method A. With this method elemental tin shot is contacted with a fast stream of dry Cl2.

My reaction vessel was a 20mm x 150mm test tube with tublature as shown in the photo below:

chlorination of Sn (starting).jpg - 91kB
chlorination of Sn (starting)

Cl2 was generated in a 500mL RBF by dripping 20%HCl on TCCA (ref 2) as shown in the photo below:

Cl2 generator.jpg - 85kB
Cl2 generator

A. Formation of Anhydrous SnCl4
The 20% HCl was fed using a 12vdc peristaltic pump with a 5/32” OD polyethylene hose (ref 3). See photo below. This worked exceedingly well compared to trying to keep a slow steady rate using a p-e funnel.

peristaltic pump.jpg - 92kB
Peristaltic pump

The Cl2 was dried using con H2SO4 as shown in the photo below:

Cl2 dryer (con H2SO4).jpg - 100kB
Cl2 dryer

The HCl feed rate in this picture was 40d/min, producing a steady, fast stream of Cl2. The Cl2 is fed to the bottom of the reactor. Unreacted Cl2 is vented to my hood suction.

Per Brauer the reactor was charged with enough tin shot to fill the tube to about 2cm below the overflow into the 100 ml RBF. This amounted to 73g. Based on this charge, my theoretical yield of SnCl4 was 160.2g, or 71.8ml.

The stoichiometric Cl2 needed was calculated to be 1.23 moles. I was prepared to deliver 2 moles in 2ea 1 mole batches, assuming that the heterogeneous reaction would not be very efficient.

As the tin is converted to the liquid SnCl4 it fills then overflows the reactor tublature. A 100mL RBF is used as a reservoir to catch the liquid as it overflows.

Shown below is the status after consumption of the 1st mole of Cl2.

chlorination after 1 mole of Cl2.jpg - 83kB
Product status after delivery of 1 mole of Cl2

The reaction produces considerable heat and it was necessary to continually add ice to the water bath enclosing the reactor. The bath temperature held about 16°C.

The system was shut down after running out of 20% HCl after the generation of the 1st mole of Cl2. The generator was then emptied of the waste product cyanuric acid and refilled with a 2nd charge of 80g of TCCA. The 20%HCl for each batch was made up from 133ml of 31.45% muriatic acid + 93mL of water. A 2nd mole of Cl2 was then sparged into the reactor as before. After the 2nd mole had been delivered there was very little tin shot left in the reactor, ie, 1.2g.

chlorination of Sn near the end of 2nd mole of Cl2.jpg - 76kB
Production after delivery of the 2nd mole of Cl2

Some flakes of Sn were added to the SnCl4 in the RBF to react with dissolved Cl2 and it was left overnight. The yellow color of the crude product was evidence of significant dissolved Cl2.

B. Distillation
Today I set up for simple distillation of the crude SnCl4 using the 100mL RBF overflow vessel as pot. I left the tin flakes in the pot and added a couple pottery chards. A CaCl2 guard tube was attached to the vacuum adapter tublature. The system was backflushed with argon before attaching the pot.

distilling SnCl4.jpg - 88kB
Distillation of the product

The distillation was rapid (several drops/sec) and without incident. Initially the condensate was a clear yellow but quickly became crystal clear. The boiling point locked on to 112.5°C and held this temperature throughout the distillation. I have no explanation for why it didn’t rise to the literature value of 114°C.

Results
The yield was 147.9g, or 66.3mL. The %yield was 92.3%.

Discussion
Generating my own Cl2 and using it successfully requires, above all, good plumbing. Systems using Cl2 usually have internal gas pressure above atmospheric. This is especially true with this synthesis. About 3.6”H2O back pressure is generated by the hydrostatic head of the con H2SO4 in the dryer. Then, added to this is about a 10”H2O hydrostatic head due to the 4.5” of SnCl4 (sp. gr. = 2.23) in the reactor. All total the Cl2 pressure in most of the system is 13.6” H2O, or ½ psi. Since I have inevitable small leaks and sometimes startup fumbles resulting in major excursions, my good fume hood serves as an important backup.

References
1.Handbook of Preparative Inorganic Chemistry, 2nd ed, by Georg Brauer (1963), forum library.
2. “Chlorine – Illustrated Practical Guide,” by Len1, Prepublication http://www.sciencemadness.org/talk/viewthread.php?tid=9713
3. “Building a Peristaltic Pump for Small Volumes/Rates,” post by Magpie, http://www.sciencemadness.org/talk/viewthread.php?tid=59870#...

Questions, suggestions, and comments are welcomed.



[Edited on 14-4-2016 by Magpie]

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The single most important condition for a successful synthesis is good mixing - Nicodem
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