Sciencemadness Discussion Board
Not logged in [Login - Register]
Go To Bottom

Printable Version  
 Pages:  1  
Author: Subject: Test For Lead in Pewter
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 15-10-2010 at 13:10
Test For Lead in Pewter


Intro
Pewterware is cheaply available from antique and second hand stores; in some cases is essentially +90% pure tin. Pewterware is useful for casting as well as for use of a source of tin for reductions etc. For casting pieces used to serve food, pewter which contains significant amounts of lead is of course undesirable. Therefore, I have developed and optimized a test for lead in pewter. There are some kinks that could still be worked out (which would only serve to lower the LOD) but as it stands the test is more than sufficient for this application.

Pewter
"Modern pewter is tin hardened by adding a few percent of other metals to make it serviceable. The usual range of composition of English pewter or Britannia metal is tin 90-94%, antimony 5-7% and copper 1-3%. Old pewter generally consisted of an alloy of tin and lead. The Romans used tin 80, lead 20%, and tin lead alloys were in general use until fairly recent times; as however, much of the old pewter was made by melting down earlier pewter vessels the composition gradually became somewhat indefinite." --Thorpe's Dictionary of Applied Chemistry, 1943.

Test
The working idea is to dissolve the pewter in aqua regia and then add an iodide which should produce the insoluble PbI2 as a yellow precipitate.

Experimental

Materials


(everything)

Fuming Nitric acid (sulfate free!)
20% Hydrochloric acid
Decolorized Iodine Solution (sold at pharmacy in first aid department, essentially a solution of ammonium iodide in aqueous ammonia with some added EtOH)
Ca(NO3)2*3.7H2O

60 Sn, 40 Pb Solder (Positive control)
95 Sn, 5 Sb Solder (Negative Control)
Unknown Pewter Platter produced after 1964

Note: The nitric acid must be sulfate free as lead sulfate is insoluble in water, likewise SnSO4 easily hydrolizes to form the insoluble oxide. Since my nitric acid was prepared from H2SO4 and a nitrate, lower bp fuming acid should carry over very little H2SO4

First attempt
Little pieces of each sample were collected and digested with 2-3 mL of 1:4 nitric/hydrochloric acid. This was hastened by heating over open flame. Once the metal had fully dissolved, the samples were allowed to cool and the decolorized iodine was added dropwise; this rapidly reacted with the excess nitric acid forming molecular iodine and little precipitate.


(Approximate sample size)

Second Attempt
The digestion was just like the first, except after the metal dissolved, the excess HNO3 was removed by reduction with sodium metabisulfite. Unfortunately this produced a lot of precipitate (of PbSO3 presumably); upon addition of the iodide the PbI2 precipitate was hard to detect.

Third Attempt
Digestion as the first, but the solutions were concentrated to ca. .3 mL by boiling, more hydrochloric acid was added and the solutions were again concentrated. This was repeated until, in total, 3 concentrations were performed. Upon addition of the iodide, the positive controls showed very clear precipitate but the solution was yellow/red from liberated iodine. The Negative control showed no precipitate however was red/yellow in color. While this test is acceptable, it is not very decisive for very small sample sizes.

Fourth Attempt
Like third, except solutions were concentrated to dryness (without overheating) followed by re-suspension in HCl, concentration to dryness and finally resuspension in HCl.
The results for this were better than the 3rd attempt, but I was still unhappy with the decisiveness of the test. Some samples showed definite liberation of iodine (as could be noticed by boiling the solutions which caused the color to lighten to nearly colorless).

Fifth Attempt
The samples were dissolved in hydrochloric acid containing calcium nitrate, about a spatula tip of calcium nitrate was used and the solution volume was 2mL. The amount of calcium nitrate should be as little as possible to fully dissolve the metal. The digestion was again hastened by heating. After digestion, the sample was allowed to cool and iodide was added. For samples which show no precipitate but still had yellow coloration, an excess of iodide was added, and the solution was boiled for a minute or two. If after boiling the solution is appreciably colorless, the test is a negative. If the solution still has distinct yellow coloration that is not significantly changed by this process, the test is a positive for lead.


(Samples being digested. From left to right: Neg. Control in Aqua regia, Pos Control in Ca(NO3)2/HCl, Pos Control in Aqua Regia, Unkown in Aqua Regia)


(Same sample order as above after digestion, the nitric acid controls have been boiled with HCl. The unknown, is showing PbCl2 precipitate and is ready for another concentration)


(concentrating down an HCl/HNO3 Sample)


(Positives: Left to Right: Pos. Control Ca(NO3)2/HCl, Pos Control aqua regia concentrated to dryness (fourth attempt), unkown aqua regia concentrated (third attempt))


(Left: HCl/Ca(NO3)2 Pos Control (5th attempt). Right: 5th attempt Neg. Control after boiling off I2)

Discussion and Conclusion
The unknown was found to contain a significant proportion of lead.

What I thought was going to be a simple test was complicated by the ease at which iodide is oxidized by nitric acid to molecular iodine and nitrogen oxides. This made it necessary to have as little excess nitric acid as possible in the test solution. Removal of nitric acid, by way of HCl oxidation was an effective way to remove the bulk of the nitric acid, however enough nitric acid remained to greatly decrease the test sensitivity. Concentrating to dryness was fairly effective but upon redissolving in HCl the solution was sometimes faintly turbid, which made positives harder to distinguish from negatives. With both the third and fourth attempts, the results were not very consistent; some showed little iodine liberation while others showed significant free iodine. I attribute this to differences in boiling time for trial 4 and residual nitric acid adhering to the test tube walls for trial 5.

The method with Calcium nitrate suffered from the same problems as that with nitric acid, except calcium nitrate is much more convenient to use. Even without boiling the possible negative tests to remove molecular iodine, the negatives were more decisive than the nitric acid tests probably because I used much less total nitrate in these tests. It seems probable that if the nitric acid tests were performed with the minimum amount of nitric acid to dissolve the sample they would be more accurate.

In the fifth attempt, boiling the solution to drive off molecular iodine is only applicable when there is excess iodide with respect to nitrate in solution. If there is excess nitrate, all the iodide will simply be oxidized and driven off as iodine. I believe that the lingering yellow coloration in my negative controls is mostly due to lead impurity in my negative control.

I am sure someone will ask, why I didn't simply neutralize the solutions with base and then run the tests. I didn't do this because I didn't want even the slight precipitation of salts to interfere with the test. These insolubles make it hard to gauge how much of the precipitate is the lead iodide and how much is simply the basic nitrate or oxide of tin.

The final word is, if you want to see if your sample contains a few percent lead, follow the fifth attempt but don't do the second part to remove molecular iodine. If you want a more accurate test for lead do the step to remove molecular iodine.

I also want to point out that I am a inexperienced inorganic chemist, so be gentle with critique :P




"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
ldanielrosa
Hazard to Others
***




Posts: 115
Registered: 25-4-2007
Member Is Offline

Mood: transparent

[*] posted on 15-10-2010 at 23:19


That works too. I prefer to go with cheap ingredients when possible. Lead detection is touched on at http://www.youtube.com/thehomescientist . He uses acetic acid and three different precipitation tests.

It seems like a lot is going into this test. Are you trying to measure the lead content as well? The oxidation of iodide to iodine from the HNO3 makes it a bit touchy as well. Maybe I'm just too stingy with my nitric acid.

I don't have my kit fully assembled yet, but I'll have to see if tin, antimony or copper will give false positives. I'm recrystallizing some lead acetate right now to make a baseline test solution.

I'm also interested in the tin from pewter, but I'd like to get rid of the antimony. The lead and copper are less important for my purposes, as I'm thinking about making bronze so I'll be adding them anyway.

[Edited on 16-10-2010 by ldanielrosa]
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 17-10-2010 at 12:07


Very well reported, I must say...

A few points. Your mixture isn't really aqua regia, not in the strict sense of the word. Real AR requires also concentrated HCl. AR is often believed to be capable to dissolve anything metallic (because it dissolves gold and some other highly noble rare metals) but that's not always true...

In the case of a lead bearing tin alloy, the choice of a mixture of HNO3 and HCl seems a little odd. Dilute HNO3 vigorously attacks Sn but causes the insoluble SnO2 to form. On the other hand, PbCl2 is only sparingly soluble in water, but better in conc. HCl.

I think that therefore you're making your life a little unnecessarily difficult, if detecting Pb is your main purpose here. Instead I would digest the alloy in straight HNO3 (38 or 35 w% will do it, fuming nitric isn't needed), the tin forming insoluble SnO2, the lead going into solution as Pb(NO3)2. The two can thus be separated fairly easily. Concentrating the clear, lead nitrate bearing solution then allows testing for lead with Cl- (white PbCl2), iodide, S(2-) (black PbS) or chromate (red Pb chromate)... Or oxidise to PbO2 with ClO-...

I've got some lead containing solder, so I might test tube this myself...


[Edited on 17-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 17-10-2010 at 15:11


@Idanielrosa: Not a lot is going into the test; I just tried many different procedures to optimize it. In essence dissolve the metal in HCl with a nitrate salt and then add OTC decolorized iodine solution.

@Blogfast: It is very close to aqua regia, as I am using what is essentially 95% HNO3 (for the reason stated in the materials section); I tried with some HNO3 that came over at the tail end of a nitric acid distillation, this produced a lot of the insoluble sulfate.

HNO3 and HCl were used because of the very thing you stated with tin, aqua regia is known to dissolve tin without passivation/oxide formation. For small sample sizes the lead choride precipitation does not occur. I originally was going to use dilute nitric acid, but all references i saw stated that tin nitrate could be formed by dilute nitric acid, but they never stated how dilute, so I went with aqua regia (which I have no problem working with). Also the HCl serves as a reducing agent to destroy excess nitrate.

If anything, I think a better way would be to dissolve the metals using hydrogen peroxide and acetic acid, boil the solution to destroy the H2O2 and add iodide. I didn't go this route because all I had was 3% hydrogen peroxide available and I wasn't sure how quickly/efficiently this would dissolve the solder.

A tip though; If you have rosin core solder, melt it first, to get out the bulk of the rosin and then dissolve it.

Thanks for your input though. Nice to see your perspective as inorganic is really not my thing.

[Edited on 10-17-2010 by smuv]




"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 23-10-2010 at 08:36


I carried out my own tests:

0.87 g of tin/lead solder (15 y old 60 Sn / 40 Pb - admittedly this is more than just a lead 'contamination') alloy and 0.91 g of lead (roadkill: from a car wheel balancing weight) where treated with about 10 ml of 38 w% HNO3. With the solder reaction started immediately but the lead was more sluggish. After briefly heating the solder it dissolved in about 15mins, forming a precipitate of SnO2. The lead by contrast had to be steam bathed for about an hour.

At the start:



The solder solution/precipitate was then filtered and slightly diluted and divided into three portions. The third was repeatedly (2 x) boiled to almost nothing to get rid of excess HNO3. The filtrate is slightly yellow because the solder contains a fluxing resin core.

Solder filtrate:



NaCl (left), K2Cr2O7 (middle) and KI (right) were then respectively added to tube 1, 2 and 3:



It's always a joy to see the typically delayed precipitation of PbCl2: in this case it was like snowing inside the tube, with crystals as long as 4 mm needles formed. With KI, no I2 was formed.

The lead nitrate solution had developed a white precipitate but clearly crystalline in nature. The clear liquid was decanted off and some 10 ml cold water added to the precipitate which dissolved quickly and completely upon shaking. That liquid was then added to the rest and divided over three test tubes.

NaCl (left), K2Cr2O7 (middle) and KI (right) were then respectively added to tube 1, 2 and 3:



Here the higher concentration of Pb caused the PbCl2 to drop out more quickly in a 'cheese-like' precipitate. Notice also how the iodide tube did develop some iodine: here the tube hadn't been boiled to expel HNO3.

Some interesting additional observations were also made. Turns out that the SnO2.n H2O precipitate was soluble in hot 22 w% HCl, from which with 5 M NaOH fresh Sn(OH)4 can be precipitated which redissolves effortlessly to Na2Sn(OH)6 with more 5 M NaOH.

That would explain why smuv's experiments the tin did dissolve into 'aqua regia' and a mixture of HNO3 and HCl may indeed be the best solvent for tin. Sn (IV), at least according to the potential series, can be reduced to Sn (II) by H2 (but H2 cannot further reduce Sn (II) to Sn (0), at least not in watery solution). So adding Zn and HCl should make the hydrogen in statu nascendi reduce the Sn (IV) to Sn 2+. Precipitating with ammonia should cause Sn(OH)2.n H2O to form, while the Zn ammonia complexes (Sn(OH) is also amphoteric - see stannites - but I doubt if NH4OH is alkaline enough). This may be a faster route to Sn (II) salts than direct dissolution in HCl which is reported to be very slow...


Some more results with HNO3/HCl mixtures with surprising results to be posted shortly...


[Edited on 24-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 23-10-2010 at 09:03


Using the same solder I then carried out some simple test tube experiments on about 1 g solder each: #1: a mixture of 50/50 38 % HNO3/22% HCl, #2: straight 38 % HNO3 and a #3: mixture of 50/50 38 %HNO3/water.

The results were a little surprising.

#1 started off quite well and no SnO2 or NO was being formed. But the reaction died down quickly and was revived with some heat, then died again. Finally I kept it on steam bath for quite a while but reaction was very, very sluggish. No SnO2 formed and no NO2 could be smelled or seen and unreacted solder was left over. On cooling, a small amount of PbCl2 crystallised out. Adding alkali and a small amount of precipitate, very floccular, formed which redissolved in an excess alkali.

I wonder whether in the case of # 1 precipitate was perhaps Sn (II), not Sn (IV), as no by-product of oxidation (NO/NO2) was observed...

#2 and #3 behaved quite similarly, #2 a little faster but both finished completely w/o external heat and both generated a white precipitate, presumably SnO2. But the precipitate of #3 dissolved better and quicker in hot HCl than that of #2. #2 generated much NO/NO2 but #3 didn't seem to produce any!

Smuv's acid mixture of 1:4 nitric/hydrochloric acid based on pure HNO3 is of course quite different from my HNO3/HCl mixture. I've got 65 % HNO3 and 36 % HCl in the pipeline and will experiment more to find a solvent for pewter or tin/lead solder that doesn't cause the tin to precipitate as SnO2.
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 23-10-2010 at 13:52


@blogfast: With dilute solutions of "HNO3" produced from Ca(NO3)2 and 20% HCl, full digestion of the metal could be achieved in less than 15 min, with intermittent heating over open flame. Maybe #1 was slow for you because you had a lower HCl concentration which caused passivation by surface oxide?

When I digested only tin containing solders, the HCl/HNO3 solution stayed pretty much colorless (even though the nitric acid was orange to begin with). The lead bearing solders developed orange coloration but, not extremely intense. The pewter samples seemed to develop more orange coloration than the lead ones, I am not sure why.

As I stated in one of my replies, according to literature (which is congruent with my experiences) aqua regia will not precipitate tin as the oxide.





"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 24-10-2010 at 04:42


Quote: Originally posted by smuv  
@blogfast: With dilute solutions of "HNO3" produced from Ca(NO3)2 and 20% HCl, full digestion of the metal could be achieved in less than 15 min, with intermittent heating over open flame. Maybe #1 was slow for you because you had a lower HCl concentration which caused passivation by surface oxide?

When I digested only tin containing solders, the HCl/HNO3 solution stayed pretty much colorless (even though the nitric acid was orange to begin with). The lead bearing solders developed orange coloration but, not extremely intense. The pewter samples seemed to develop more orange coloration than the lead ones, I am not sure why.

As I stated in one of my replies, according to literature (which is congruent with my experiences) aqua regia will not precipitate tin as the oxide.



Yes, agreed on all points.

In my case the light yellow colouration comes from the flux I think. But it doesn't bother me: in my case the flux comes up floating on the solution as an amber blob, easy to remove. And it's dissolving pewter that I'm after anyway...

Where you write "(even though the nitric acid was orange to begin with)", did you mean the nitric or the actual aqua regia? The latter is supposed to be orange due to the presence of nitrosyl chloride (NOCl), a strong oxidiser.

I've some Ca(NO3)2 but in solution. I remember it being extremely hygroscopic, almost impossible to keep as a solid and I've no other nitrates (KNO3 or NaNO3) right now but could easily make some... So I might try that... Otherwise I'll juts wait for the 65 % HNO3.
View user's profile Visit user's homepage View All Posts By User
chemrox
International Hazard
*****




Posts: 2245
Registered: 18-1-2007
Location: Milky Way
Member Is Offline

Mood: intensely curious

[*] posted on 24-10-2010 at 13:06


Can't you simply dissolve in HCl and treat with iodide with a yellow solution as positive?



Aviator-talk to me about flying
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 25-10-2010 at 07:16


As luck would have it, I just managed to buy about 0.5 kg of pewter (an old but ugly 'tankard') cheaply, so I've now got plenty tin to carry on testing... Will try adding some nitrate to HCl shortly.

[Edited on 25-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 25-10-2010 at 09:23


Et voila, 2 quick tests with the pewter, #1: 0.4 g with a few ml of 22 % HCl; #2: 0.9 g with 1 ml HNO3 38 w% and a few ml of HCl 22 %.

First up #1 which at RT (but it's cold in my lab) reacted slowly but very noticeably. On heating in bunsen flame the reaction rate becomes entirely acceptable, vigorous I'd say and finished in about 10 min, reheating regularly when needed. On steam bath, dissolving this pewter with conc. HCl should be doddle.

#2 was very similar but faster I'd say (I didn't compare in real time). No Sn oxide was formed. A slight yellow colour developed (NOCl?). The question is what's the oxidation state of the Sn in this run: +II or +IV? Do nitrate ions oxidise Sn +II to +IV? Since as air oxygen does it I'm guessing NO3- can do it too but I saw or smelled no NO/NO2... I'll have to look that up. The oxidation state +IV is what I need for my purpose.

In both cases there was a small amount of unknown black insoluble residue, the solutions were filtered.

Here's the two test tubes after very careful neutralisation with 5 M NaOH. That doesn't tell much at all about the oxidation state of either, although I assume that #1 is pure Sn +II.



View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 25-10-2010 at 13:11


About the flux. I mentioned in my first post, you can remove that by melting the solder before you digest it. I never noticed anything I could attribute to flux after doing that.

About the orange color of the NA. It was red fuming acid, upon addition to the soldering alloys it lightened in color. I think the tin reduces some of the species to less colored products.

So if I understand this correctly blogfast, you used HCl alone to digest mug that had lead free pewter? Or did it contain lead?

It was a good idea blogfast to buy something meant to be eaten out of, even if it is old there is a much lower chance of it containing lead.

@Chemrox: does lead dissolve in conc. HCl with an appreciable rate?





"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 26-10-2010 at 05:14


Quote: Originally posted by smuv  

So if I understand this correctly blogfast, you used HCl alone to digest mug that had lead free pewter? Or did it contain lead?

It was a good idea blogfast to buy something meant to be eaten out of, even if it is old there is a much lower chance of it containing lead.



Yes, 0.4 and 0.9 g of it. I didn't detect lead but didn't really test for it either. I will when I scale up (this afternoon).

The choice of pewter was rather by accident than by design: a charity junk shop just happened to have a pewter tankard and I jumped on it. A lot of pewter still around is in the form of eating utensils of course, that was one of the main uses.

Redox potential series shows that the reduction of HNO3 to NO is capable of oxidising Sn (II) to Sn (IV), easy peasy. This aft I'll see if I can see any difference between the 2 oxide samples...
View user's profile Visit user's homepage View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 26-10-2010 at 12:41


Some more tests...

The oxides in both test tubes were dissolved in HCl, both dissolved well, #2 (very subjectively) a wee bit slower.

The black, insoluble residues of both test were decanted and combined and subjected to 38 % nitric. The residues reacted vigorously, SnO2 and NOx were formed, it would appear the residue is mainly tin. Filtering gave a lime green filtrate:



This did not respond to K2Cr2O7 indicating this pewter is relatively lead-free. Also it didn't respond to NH3 1.5 M: no precipitate and no blue complex: this rules out Cu2+...

Two other, larger scale test were then carried out with the aim of synthesising some more (NH4)2SnCl6.

First a test with 3.85 g pewter and 20 ml HCl 22 %. This is the stoichio amount (+ small excess) for obtaining SnCl4 (even though in the first step only SnCl2 is formed). Dissolution on steam bath took about 20 mins but left some black insoluble residue behind. The solution was filtered and 3 g of HNO3 38 % was added: this is stoichoi + plus small excess for 3 Sn (II) + 2 HNO3 --> 3 Sn (IV) + 2 NO (not fully balanced here). That was simmered for a few minutes. No NOx was observed. Then about 3 g of solid NH4Cl was added and the solution simmered with reflux until boiling behaviour became odd. On cooling to RT (then refrigerator), typical crystalline sheets of (NH4)2SnCl6 appeared. Nice crop too!

The second test (with 5.1 g pewter) was very similar but here the HNO3 was added right from the start to the HCl. Events unfolded quite differently though.

Although this too started out fine at BP, about half way the reaction completely died, due I believe to passivation of the tin which manifests itself in a stubborn black sludge coating the metal. It was removed, the solution set aside and the metal scrubbed with sand paper and dissolved in virgin HCl (w/o HNO3). Some black residue remained. I combined the first solution with the latter and added some more HNO3 (I patently forgot to filter first!). On simmering something strange happened: all of a sudden the solution 'whooshed', threw some outside the beaker (not very much) and dissolved also the residue! On top of that, the solution was now greenish:



This was brought to the boil and solid NH4Cl was added and something strange happened again: a whitish precipitate formed. SnO2? (NH4)2SnCl6? Something else? What ever it was on adding quite a bit of HCl 22 % and heating, the precipitate dissolved completely: by then the volume was about 125 ml. That was boiled down to half way and I got my supper. On returning, in the cool solution quite few white crystals had formed. Very well formed. Except... they didn't look like any of the (NH4)2SnCl6 previously obtained. I collected the crystals on a plastic tea strainer, then washed them with cold DIW: they are essentially white. They dissolve in water and tested negative for Pb2+ (I suspected maybe Pb(NO3)2 because they look like it) but positive for behaviour of (NH4)2SnCl6: the solution sheds SnO2 on heating and also when strong alkali is added.

Finally the rest of the solution was reduced and cooled: more crystals formed...


[Edited on 26-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 26-10-2010 at 13:58


Very cool blogfast. I love how experimental and photo-rich this thread has become.

Keep in mind that there are non-negligable amounts of copper and antimony in lead free pewter. Cupric oxide could be responsible for some of the insoluble crap.

About the (NH4)2SnCl6, why don't you take a sample of the product you are unsure about and heat it in a test tube with aqueous NaOH. If the smell of ammonia is noticed as well as precipitation of Stannic oxide (that part you already recorded) I think it is pretty fair to say you have ammonium hexachlorostannate,

I also would like to warn you, that most 'pyrex' kitchenware is no borosilicate glass. So watch out with heating over naked flame if it isn't. You can check by looking at the edges of the glass, if they are strongly blue/green they are normal soda-lime glass, if they are only faintly green you have borosilicate.




"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 27-10-2010 at 04:44


Quote: Originally posted by smuv  
Very cool blogfast. I love how experimental and photo-rich this thread has become.

Keep in mind that there are non-negligable amounts of copper and antimony in lead free pewter. Cupric oxide could be responsible for some of the insoluble crap.

About the (NH4)2SnCl6, why don't you take a sample of the product you are unsure about and heat it in a test tube with aqueous NaOH. If the smell of ammonia is noticed as well as precipitation of Stannic oxide (that part you already recorded) I think it is pretty fair to say you have ammonium hexachlorostannate,

I also would like to warn you, that most 'pyrex' kitchenware is no borosilicate glass. So watch out with heating over naked flame if it isn't. You can check by looking at the edges of the glass, if they are strongly blue/green they are normal soda-lime glass, if they are only faintly green you have borosilicate.


Well, the green is definitely not due to copper (cupric chloride is greenish but no reaction with NH3 was obtained).

Re. the (NH4)SnCl6, it will be analysed shortly by another method, all being well. But firstly I'll have an attempt at determining the Sn content of the pewter. This should be possible by titrating as follows: Sn2+ + 2 Fe3+ --> Sn4+ + 2 Fe2+ (this reaction should proceed according redox potential series), against KSCN as indicator.

If it works it should be possible also to titrate Sn4+ (and thus (NH4)2SnCl6) by firstly reducing it back to Sn2+ first (with hydrogen in statu nascendi: Al ribbon + acid). This would be similar to redox titrating titanium (reduced first to Ti3+).

Note that I've synthesised the hexachlorostannate also from crude SnO2, so I know what it looks like... (see 'SnO2 fusion' thread, if you're interested...)

Re. 'pyrex-not pyrex', my experience with Pyrex (TM) kitchenware in the lab, including open flame (nearly all the time), has been very positive over the many years. Some 'beakers' have been going for years now. I've occasionally lost one due to thermal shock but IMHO that's more because they're very thick walled which on sudden cooling caused a massive temp. differential between the cold side and the hot side, with resulting high stresses...

Thanks for your feedback...


[Edited on 27-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 27-10-2010 at 08:26


Quote:
Well, the green is definitely not due to copper (cupric chloride is greenish but no reaction with NH3 was obtained).


While your observations may not support it, there is a very good chance of copper being in your pewter. All the metals in modern pewter serve distinct purposes, therefore the absence of copper would lead to a very inferior pewter.

Tin -- Base metal
Antimony -- to stop tin rot
Copper -- to harden the alloy (and probably make it flow better while casting)

About everything else, I like where you are going with this, it will be pretty useful to have a relatively easy method of determining the tin content in pewterware. The problem is, you are essentially assuming that no other ions in solution will be reduced by the aluminum ribbon. This is all well and good for antimony, but if there is copper present, which there almost always is, in theory this should be reduced to copper(I) which will readily reduce Fe(III). While I don't have much experience reducing metals with aluminum ribbon, I think that it may be unreliable for quantitatively reducing all Sn ions in solution, because most of the aluminum will be consumed simply reducing water. I would pick another reducing agent. Can SO2 reduce Sn(IV) to Sn(II)?






"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 27-10-2010 at 13:07


Firstly, it's not the aluminium that does the reducing here but the hydrogen, in particular monoatomic hydrogen (hydrogen radicals, if you prefer), generated by Al + acid ---> Al3+ + H* (2 H* --> H2). This method (also with zinc or iron) is often used as a reduction agent. It reduces for instance Ti (IV)aq to Ti (III)aq. SO2? Worth looking into, IMHO...

Re copper, I'd have to look that up, I think it might be reduced to elemental copper, as antimony. But I have to look both up.

Here's what I did.

1.92 g of pewter was first attacked with 25 ml simmering HCl 22% and when that reaction showed signs of slowing down, 5 ml of HNO3 38 % were added. The reaction that followed was almost instantaneous: large eruption of NOx, remaining residue completely disappeared and a clear, lime green solution formed. This is the first time things were so straight forward with my reagents and pewter.

The solution was then quantitatively transferred to a Nalcene round calibrated 250 ml measuring flask and diluted to the mark, resulting in a stock solution of approx. 0.1 N (=0.05 M) Sn, probably mainly Sn (IV).

The titrant solution was ferric alum 0.1 N (0.1 M) but years old. Kept in the dark and well stoppered it looked in good condition with no precipitation or observable hydrolysis. Of course it would at least have to be re-standardised but here the purpose was to show a principle.

My titration station:



20.0 ml of stock solution were pipetted into a clean conical flask and 25 ml HCl 22 % and 0.6 g of Al added. The Al used is extremely low in C, BTW from similar experience (titration of Ti (IV)aq. ) After a bit of heat the reaction starts. And here's the snag: fairly small amounts of an amorphous looking black powder started to form. After all Al was digested, 3 ml of KSCN (Fe3+ indicator) were added. I decided to titrate with the precipitate present. It's clear the reaction Sn (II) + 2 Fe (III) ---> Sn (IV) + 2 Fe (II) [1] proceeds: each time you add some titrant solution, temporarily the red colour of FeSCN (2+) appears then fades quickly as the Fe (III) is reduced away.

A the start of a reduction: no black residue formed yet:



But there's a snag: endpoint determination was difficult with a slow and 'spongy' endpoint. End point was around the 20 ml area.

The second test was identical, except that I filtered the reduced solution quantitatively (it was then clear as water) and then titrated it. That improved things but endpoint was till difficult to determine, I believe because at near endpoint of [1] the reaction speed had dropped to almost nothing. Near the end, when you add another drop the solution turns, say amber red, but in minutes gets back to completely clear! A remedy would be to increase the concentration of both solutions to about 0.5 N, instead of 0.1 N.

But it's been shown that titrating Sn with Fe3+ after reducing to Sn (II) should be possible in the right conditions.

As regards the nature of the black residue, I'll get some more tomorrow to investigate...

[Edited on 28-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 27-10-2010 at 20:15


Quote:
Firstly, it's not the aluminium that does the reducing here but the hydrogen, in particular monoatomic hydrogen (hydrogen radicals, if you prefer), generated by Al + acid ---> Al3+ + H* (2 H* --> H2). This method (also with zinc or iron) is often used as a reduction agent.


I know where you are coming from with this, but the whole nascent hydrogen thing has been disproven decades ago. Of course, it is mentioned in older texts, before proven false. Instead these reactions generally proceed via electron transfer directly from the metal. Therefore, the metal is the reducing agent, not the hydrogen.

About the black powder, you should dissolve some of that aluminum in HCl to make sure that the aluminum itself is not the source of this precipitate (it may be very pure, but its good to cover all bases).

About the endpoint, I wonder if there are any readily available redox indicators that could be used instead of the SCN-.





"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 28-10-2010 at 04:54


Two pix were added to the post above.

If it really is direct electron transfer, the copper (if any) should plate out: in a neutral solution of Cu (II) aluminium goes into solution while being plated with copper (a classroom redox demo!) I was thinking about that last night and will test this with CuCl2, lots of HCl and the same Al used above... In a neutral solution of Sn (IV) (precarious because of hydrolysis!) reduction would almost certainly also take place with aluminium but to what new oxidation state?

In any case, the redox potentials show that Sn (IV) + H2 ---> Sn (II) + 2 H+ must proceed.

Most of the remainder of the pewter stock solution will now be reduced in the same conditions as above, thereby obtaining more of said residue for investigation.

I'm very much inclined to exclude Al as the material: I've used this grade too often for redox titrations of Ti (III): the solution always cleared up completely before titrating with ferric alum...

I think, reading up on antimony, that Sb is a more promising candidate. Acc. Holleman, (again 'statu nascendi') hydrogen bubbled through a soluble antimony compound solution yields small amounts of stibine - SbH3. But this, according to the same source, is very vulnerable to oxygen, which causes elemental Sb to form. That kind of fits, I think. I'll also carry out a flame test on the pewter stock solution, copper might show up.

Re. indicators, yes, there exist of course redox indicators but I haven't got any of them. Potentiometric endpoint determination may also be possible here but I need to look onto that... If I'm right that reaction speed really is the problem, then not even a perfect system of endpoint determination would solve the problem: only a catalyst would...
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 28-10-2010 at 18:18


How about methylene blue.



"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 29-10-2010 at 11:46


Yes, methylene blue would be one of them. Need to find out what's redox potential at the turn, though. And I haven't got any right now.

The test with CuCl2 + HCl + Al surprised me a little. Here's a pic in mid-reaction: Cu is dropping out like the clappers, much faster than would happen in neutral conditions (all Cu dropped out):



Regards the residue, I'm now more or less convinced that it has to do with the Al and with the acid reserves in the sample, which may have been borderline. These are teething problems that should be easy to iron out...

Residue:



It glistens because it's still wet.

Yesterday I made 250 ml stock solution that was approx. 0.5 N in Sn (IV) and 250 ml of ferric alum ( 0.5 N, as yet not standardised) (in 0.1 M HCl) as titrant solution. This works much, much better than 0.1 N: endpoint is still a bit shaky but much easier to establish. Tomorrow the 0.5 N solution will be standardised.

250 ml of approx. 0.5 N ferric alum in 0.1 M HCl:



Finally, another 10 g pewter was dissolved into HCl/HNO3 and an excess of NH4Cl added. The solution was then gently simmered down until the first crystals appear.

After a bit of standing (solution still hot - about 50 C) and a nice crop of (still presumed (NH4)2SnCl6) crystals has already appeared (yield determination tomorrow):



I'll also try and make the potassium salt, which is likely to be even less soluble at RT and which has been used for the separation of Rb and Cs.

Finally regarding the antimony: it should be detectable as (black) Sb2S3 in not too acidic conditions by saturating with H2S.

O/T: silly chemists' joke:

Two amphoterics at the bar.

Says the first one: hey, you're putting them down fast, watch out or you'll get acid burn!

Number two: you're right, I'll slow down because I've got a date later on and I'm hoping to touch first base!



[Edited on 29-10-2010 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
smuv
International Hazard
*****




Posts: 842
Registered: 2-5-2007
Member Is Offline

Mood: Jingoistic

[*] posted on 29-10-2010 at 18:24


Methylene blue: E0 = 0.53v @ pH = 0.

Really odd your mug doesn't have copper in it. Does the metal seem rather soft?




"Titanium tetrachloride…You sly temptress." --Walter Bishop
View user's profile View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 30-10-2010 at 04:13


Quote: Originally posted by smuv  
Methylene blue: E0 = 0.53v @ pH = 0.

Really odd your mug doesn't have copper in it. Does the metal seem rather soft?


No, not really. I've known softer pewters. It does hacksaw with difficulty like all (relatively) soft metals. It's also a rather young artefact, probably less than 20 years old: the bottom is flat glass (completely see-through) presumably for those who want to combine their drinking sessions with spot of submarine watching...
View user's profile Visit user's homepage View All Posts By User
blogfast25
International Hazard
*****




Posts: 5670
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline


[*] posted on 31-10-2010 at 09:48


First off, standarsisation problems. The idea was to standardise the ferric titrant using Fe3+ + I- --> Fe2+ + 1/2 I2, titrating the I2 with sodium thiosulphate. The thiosulphate can then be standardised with Cu2+ + 2 I- --> CuI + 1/2 I2.

The thiosulphate solution was years old but seemed in good condition and give a titre of 0.99 against a purposely prepared primary standard of 0.2 M Cu(NO3)2. But the standardisation of the ferric titrant was erroneous. I believe I wasn't using enough KI and now I'm clean out of it anyway, so all this will now have to wait till I get some more KI or NaI, or make some (I've got some sublimed, prilled iodine).

I'm not so interested in determining the Sn content of the pewter anymore (it appears satisfactorily high for my purpose) but it would be great to determine the Sn content of the ammonium hexachlorostannate (theor. about 32 w%) as it would confirm the composition.

The last batch of the product from about 10 g of pewter gave a yield of about 55 % (based on pewter = 100 % Sn), really well formed white crystals. It does indicate the product is quite soluble at RT. The crystallisation from a hot saturated solution appears to start at about 50 C. From the supernatant liquid plus some extra HCl all the other odds and sods of (NH4)2SnCl6 have now also been recrystallised.

From 5 g of pewter some (presumably) K2SnCl6 was synthesised uneventfully by replacing NH4Cl by KCl in the recipe. A nice crop of remarkably similar white crystals were obtained, yield of which will be determined tomorrow.

What really remains a mystery is the green colour. Believing it might be ferrous stuff, I tested it with KSCN after adding peroxide, also hypochlorite, but no FeSCN 2+ showed up.

What is most striking is how it shows up: as explained I use HCl to dissolve most of the pewter, then add enough HNO3 to oxidise all Sn (II) to Sn (IV) and dissolve the last of the black residue that appears insoluble in HCl alone. At that point the solution is still ever so slightly turbid and looks a bit greyish. Sometimes the nitrate then leaves the party altogether with a puff of NOx and the turbidity is then instantaneously replaced with the green colour. Sometimes the turbidity just disappears over a few seconds, being replaced by the green colour.

This strongly points to something in the pewter being dissolved ate the end of the dissolution process and forming a green compound. But what?
View user's profile Visit user's homepage View All Posts By User
 Pages:  1  

  Go To Top