Tin(ii) oxide preparation from pewter, failed attempt

Tin(ii) oxide, or stannous oxide, is a amphoteric oxide use in the maufacture of cranberry glass. It also finds some use as a catalyst for esterifications, however this is not common. It exists in three forms, a hydrated form which is a tan-cream coloured powder, a blue-black powder, and a metastable red powder. In the future, I intend to use tin(ii) oxide in a thermite reaction to produce tin metal.

The inspiration for this experiment came from this patent (link). First, a solution of tin(ii) chloride is prepared. Oxalic acid is then added which precipitates tin(ii) oxalate. The tin(ii) oxalate is then heated with ammonia to give the tin(ii) oxide as the blue-black powder form.

Tin(ii) chloride

To prepare the initial tin(ii) chloride solution, I used pewter, which is an alloy consisting of about 95% tin with the rest being copper, bismuth and antimony. These other components aren't a problem though as they are very unlikely to make it through the process.

To a 300ml beaker, I added 18.95g of pewter (powder and lumps). Using a graduated cylinder, I then added in 66ml of 33% hydrochloric acid and swirled the beaker. A fair amount of bubbling occurred which gradually diminished after a few minutes. I allowed The mixture to stand over night, then covered the breaker with cling wrap and heated the mixture at a low temperature no more bubbling occured, which took about 5 hours. Every now and then the mixture was stirred and a small amount of 33% hydrochloric acid was added to allow for liquid lost by evaporation. a small amount of hydrogen gas was given of at a steady rate for the first few hours, after which the hydrogen became indistigwishable from the bubbles of evaporating liquid. To remove undissolved material, I then filtered the mixture, collecting the clear filtrate in a 250ml beaker. The filtrate should be a roughly 40% solution of tin(ii) chloride. The next step is to convert this to tin(ii) oxalate.

Tin(ii) oxalate

19.12g of oxalic acid dihydrate and 45ml of water were added to a 250ml beaker and heated with stirring to around 60 C, whereupon the oxalic acid dissolved. The tin(ii) chloride solution prepared above was then added in small portions, with stirring in between additions while maintaining the temperature at around 60 C. Complete addition took around 40 minutes. With the first few additions, the mixture slowly became cloudy with fine white precipitate. The precipitate then redissolved towards the end of addition. The slightly yellow, clear solution was then taken off heat and allowed to cool for 1 hour. Beautiful needle-like crystals gradually precipitated as the solution cooled. These crystals should be the product, tin(ii) oxalate. After 1 hour of cooling, The crystals were filtered off and carefully washed on the filter with 100ml of cold water in portions and finally dried, yielding 6.05g of tin(ii) oxalate as white, needle-like crystals.

Tin(ii) oxide (failed)

The 6.05g of Tin(ii) oxalate prepared above was added to a 250ml beaker. 20ml of water was then added with stirring to form a suspension. 4ml of 25% ammonia solution was added and the mixture was heated to 60 C whereupon the tin(ii) oxalate dissolved. The temperature was maintained at 60 C for 40 minutes with occasional swirling of the beaker. After the first 20 minutes, an additional 4ml of 25% ammonia solution was added resulting in a white precipitate with a crystalline structure identical to the starting tin (ii) oxalate. After the 40 minutes of heating, no black tin(ii) oxide had precipitated as claimed by the patent and the experiment was abandoned.

Sn + 2 HCl ==> SnCl2 + H2   /   SnCl2 + H2C2O4 ==> SnC2O4 + 2 HCl

NH3 + H2O <==> NH4OH   /   SnC2O4 + 2 NH4OH ==> (NH4)2C2O4 + SnO + H2O

Ethyl acetate synthesis by Fischer esterification

Ethyl acetate, or ethyl ethanoate, is a simple ester widely used as a solvent and flavoring. It exists at room temperature as a colourless liquid with a sweet fruity smell. It also serves to some extent as a chemical precursor. There are a few reactions I plan to perform with ethyl acetate in the future.

The most well known method for producing ethyl acetate is Fischer esterification. I've included a brief description of the mechanism at the end of this post. The procedure consists of refluxing acetic acid with ethanol in the presence of a catalytic amount of sulphuric acid followed by some workup and purification steps.

The glacial acetic acid used in the procedure was first purified by drying over anhydrous copper(ii) sulphate, distilling, then drying again.

First off, I set up an ice bath and placed in it, a 500ml round bottom flask. To the flask, I added 28ml (0.4797 moles) of ethanol and 27ml (0.4716 moles) of glacial acetic acid. Once the temperature of the mixture had dropped to around 12 C, I began adding 6ml of 98% sulphuric acid in very small portions while swirling the flask at such a rate that the temperature never rose above 20 C. Once complete addition of the acid was achieved, I attached a Liebig condenser, with cold water circulating, to the flask. I then gently refluxed the homogeneous mixture for 30 minutes using a water bath as the heat source. After reflux, there was very little if any visible change in the mixture, which remained a clear liquid. I allowed everything to cool to room temperature, then removed the condenser and set the flask up for simple distillation. I then distilled off approximately two thirds of the mixture into a 150ml beaker. Care was taken to cover the connection between the vacuum adapter and receiving flask with plenty of cling wrap to prevent evaporation of the distillate. I washed the distillate in the flask with 14ml of saturated sodium bicarbonate solution, then with 15ml of concentrated calcium chloride solution. Each of these washings resulted in a two-layered system with the upper organic layer containing the ethyl acetate. With each wash, the layers were thoroughly stirred together for 2 minutes. I then separated off the organic layer and dried it over 3A molecular sieves for 40 minutes. The dry liquid was then transferred to a 500ml flask and another simple distillation was performed.

The results of the distillation were unclear, so I added all the fractions back to the 500ml flask and performed the distillation again. The first fraction came over at about 66 C and was discarded. Only one other fraction came over which was presumably the ethyl acetate at 68-73 C. In the end, I was left with 10.5ml (0.1075 moles) of dry ethyl acetate as a colourless liquid with a sweet fruity ethereal aroma. If pure, this is a 22% yield. The density of the product was 0.79g/cm³ which unfortunately isn't overly close to the established value of 0.902g/cm³. Considering this, and the smell which is identical to online descriptions, I believe the product is mostly ethyl acetate, albeit not very pure.

I blame my extremely poor yield on the fact that I was working on a smaller scale then I'm used to and that I didn't dry the ethanol before the reaction.

Brief description of Fischer esterification: The acid catalyst first protonates the carbonyl oxygen on the carboxylic acid to form a charged oxonium ion. The oxonium ion causes the carbonyl carbon to have a partial positive charge. The alcohol then attacks the now partially positive carbonyl carbon forming a complex intermediate containing another oxonium ion, which when it encounters another alcohol molecule, protonates it, thus transferring the oxonium ion to the alcohol.

One of the resulting neutral molecule's hydroxyl groups is then protonated by the acid catalyst forming yet another oxonium ion. The O(+)H2 oxonium ion then breaks off the molecule as H2O, leaving behind the ester product with a protonated carbonyl oxygen. Finally another alcohol molecule comes along and grabs the hydrogen from the carbonyl oxygen, once again transferring the oxonium ion onto the alcohol and yielding the neutral ester product.

H2SO4 <==> 2 H [+] + SO4 [2-]

CH3-C(=O)-OH + H [+] <==> CH3-C(=O[+]H)-OH

CH3-C(=O[+]H)-OH + C2H5OH <==> CH3-C(-OH)(-OH)-O[+]HC2H5

CH3-C(-OH)(-OH)-O[+]HC2H5 + C2H5OH <==> CH3-C(-OH)(-OH)-OC2H5 + C2H5O(+)H2

CH3-C(-OH)(-OH)-OC2H5 + H [+] <==> CH3-C(-OH)(-O[+]H2)-OC2H5

CH3-C(-OH)(-O[+]H2)-OC2H5 <==> CH3-C(=O[+]H)-OC2H5

 CH3-C(=O[+]H)-OC2H5 + C2H5OH <==>  CH3-C(=O)-OC2H5 + C2H5O[+]H2

Ammonium permanganate

Ammonium permanganate is an intriguing salt which at room temperature exists as slightly bronze-metallic purple crystals. It is a moderately strong primary explosive but has very rarely found much use in this field due to its high sensitivity, short shelf life and low power in comparison to the more widely used explosives. I don't really have any use for the compound and mainly just made some because of its interesting composition and to test out its explosive properties.

To make the ammonium permanganate, I followed this procedure (link) which seemed to give very satisfactory results.

To a 150ml beaker, I added 6.18g (0.0391 moles) of potassium permanganate and 17g (0.3178 moles) of ammonium chloride. I then added in 116ml of water using a graduated cylinder and stirred the mixture to dissolve everything. A dark purple solution was obtained and the beaker became intensely cold to the touch. Using a retort stand, I secured the beaker over a water bath heated by a hotplate. Medium heat was then applied and a thermometer was inserted into the mixture to monitor the temperature. Most of the time the temperature of the mixture stayed at around 80 C and care was taken not to let it go much higher then this. The liquid was occasionally stirred with the thermometer. Once the volume of the liquid had been reduced to 74ml, I took the mixture off heat and rapidly filtered it while hot into a 250ml conical flask to remove brown manganese dioxide. This is formed by some of the heat sensitive permanganate ions decomposing.

Anyway, I then chilled the flask containing the filtrate in an ice bath to 5 C whereupon a fair amount of ammonium permanganate precipitated as dark purple felted, short needle-like crystals. The mixture was then immediately filtered to collect the product, pressing it on the filter to remove as much liquid as possible. After drying, I obtained 1.68g (0.0123 moles) of ammonium permanganate which if pure, corresponds to a yield of 31%.

Left = dry ammonium permanganate  /  Right = mildly energetic decomposition of  ammonium permanganate (results in large cloud of manganese dioxide)

The reaction is a relatively simple double displacement with ammonium permanganate being forced out of solution by exploiting its radically reduced solubility in ammonium chloride solution at low temperatures. Hence the large excess of ammonium chloride.


KMnO4 + NH4Cl ==> NH4MnO4 + KCl

tert-Butyl chloride

Tertiary butyl chloride, also known as 2-chloro-2-methylpropane (IUPAC) is a tertiary alkyl halide used industrially as a precursor to pharmaceuticals, agrochemicals and fragrances. It also finds some use in organic chemistry as a reagent for introducing the tert-butyl group into molecules. Under standard conditions, tert-butyl chloride is a clear colourless mobile liquid with a boiling point of 51 C. I don't really have any particular use for the compound, although I might try synthesizing a tert-butyl ether sometime in the future.

tert-Butyl chloride can be easily prepared in the lab from tert-butanol and hydrochloric acid. I've included a description of the reaction mechanism at the end of this post. A large excess of hydrochloric acid is required to shift the equilibrium in favor of the product. Here is what I did.

To a 500ml round-bottom flask, I added 50ml (0.5228 moles) of tert-butanol and 150ml (1.3549 moles) of 33% hydrochloric acid (159% excess). I loosely sealed the top of the flask with some cling wrap and a few rubber bands, then gently shook the flask for around 1 minute, which resulted in a cloudy and completely homogeneous mixture. After shaking, I allowed the mixture to stand and separate into two layers for 20 minutes. A clear upper layer separated almost immediately and gradually became larger. After the 20 minutes, about two thirds of the mixture were carefully poured into a 150ml beaker and a syringe was used to remove the upper tert-butyl chloride layer, which was added to a 250ml flask. Once no more tert-butyl chloride could be syringed off, the remaining liquid was discarded and the remaining one third of liquid in the round-bottom flask was added to the 150ml beaker and the process was repeated (using a syringe to transfer the upper tert-butyl chloride to the 250ml flask. The tert-butyl chloride was then thoroughly washed with 90ml of saturated sodium bicarbonate solution to remove all traces of hydrochloric acid. After removing and discarding the lower aqueous layer, I washed the tert-butyl chloride with two 50ml portions of water in succession. The crude product was then added to a 500ml round-bottom flask containing an excess of anhydrous calcium chloride.

After a bit of swirling, the tert-butyl chloride lost its cloudiness and became clear as the calcium chloride removed water. I set the flask up for simple distillation, using a 250ml flask as the receiver. I then began the distillation. Only one fraction passed over, which was the tert-butyl chloride at 48-49 C. After distillation, the product still looked a bit cloudy, so I added a bit of anhydrous calcium chloride, whereupon it became crystal clear. Finally, I transferred the product to a bottle for storage. I got 34ml (0.3126 moles) of tert-butyl chloride as a colourless liquid which works out to a respectable yield of 60%. The product had a density of 0.782g/cm³ which is very close to the established 0.851g/cm³ of pure tert-butyl chloride. This shows my product is relatively pure.

Brief description of the SN1 reaction: a hydrogen ion from the hydrochloric acid first attacks the hydroxyl group of a tert-butanol molecule, forming an intermediate compound containing a positively charged OH2 group. The OH2 group is not very stable and acts as a leaving group, breaking off the molecule as H2O and transferring its positive charge to the alpha carbon. This leaves a carbocation which is then attacked by the chloride anion left over by the hydrochloric acid, forming the final product, tert-butyl chloride.


HCl <==> H (+) + Cl (-)

(CH3)3COH + H (+) ==> (CH3)3CO(+)H2

(CH3)3CO(+)H2 ==> (CH3)3C (+) + H2O

(CH3)3C (+) + Cl (-) ==> (CH3)3CCl

Boron, revisited

In a previous post (link), I described how I prepared a very crude sample of the element boron. I decided to have another go at boron using a different method to see if I could get a nicer sample. In my post on making silicon, I used a side reaction of aluminium and sulphur to thermally sustain the reaction. This gave nice large beads of silicon and I hoped that it might do the same for boron. Although the melting point of boron is significantly higher (over 2000 C) this didn't seem to be a problem.

To begin, I added 8g of finely powdered boron trioxide and 10g of powdered sulphur to a steel can. I then added in 11.2g of 325 mesh aluminium powder and stirred the mixture thoroughly to intimately mix the chemicals. Next, I poured the mixture onto a brick and placed a small amount of potassium nitrate/magnesium flash powder on top, then inserted a strip of magnesium metal. The flash powder and magnesium ribbon serve to kick start the thermite reaction. Using a torch, I carefully ignited the flash powder which immediately set off the thermite. Large amounts of heat and flame were given off and near the end, I noticed the fire was tinted green, presumably due to the presence of boron. Once the thermite was finished and the residue had cooled sufficiently, I used a hammer to remove as much of the residue as I could from the brick. I then added the residue to a 500ml beaker containing 200ml of water. Rapidly after addition, lots of bubbling was observed, which continued for over an hour. This is aluminium sulphide formed in the helping reaction reacting with the water to form aluminium hydroxide and hydrogen sulphide gas. After about 2 hours, the rate of gas evolution had decreased significantly and I decided to move on to the next step. I swirled the beaker, waited a few seconds, then quickly decanted as much of the gray aluminium hydroxide suspension as I could without losing any of the other material on the bottom of the beaker.

I repeated this process until almost all the aluminium hydroxide had been removed. To the material remaining in the beaker, I slowly added 50ml of 33% hydrochloric acid. A very vigorous reaction was observed with some foaming and lots of bubbling. I allowed the mixture to stand for about 2 hours to make sure the hydrochloric acid had dissolved all that it could, then decanted off the acid. At this point, I got my first proper look at what seemed to be the boron. Most of it seemed to be present as a black powder mixed in with lots of coloured impurities and hopelessly unrecoverable. However, there were a few beads of boron which looked fairly easy to separate. I washed the boron with 20ml of ethanol and after decanting the ethanol, then carefully picked out as many pieces of the black metallic boron as I could.

After drying, I was left with 0.7g of what I presumed to be boron. It certainly resembles many of the pictures I've seen on the net. In conclusion, this method doesn't seem very viable to produce large amounts of the element, but it's certainly a good way to get a nice sample.

2 Al + B2O3 ==> 2 B + Al2O3

2 Al + 3 S ==> Al2S3

Al2S3 + 6 H2O ==> 2 Al(OH)3 + 3 H2S

Bromine, preparation and isolation

Bromine is a fascinating element that exists as a strongly fuming dark red, almost black liquid at room temperature. It finds extensive use as a reagent, particularly in organic chemistry and is used commercially in flame retardants and as a disinfectant for pools.

There are many ways to make bromine. After considering a few different methods, I decided the oxidation of bromide with chlorine (generated in situ from hydrochloric acid and TCCA) was the best. So I went ahead and tried it out.

Unfortunately, the handing of pure bromine poses some extreme safety hazards, being incredibly toxic, corrosive and volatile. It can easily kill you if you're not prepared, and that's not an exaggeration. This procedure should only be performed by a chemist with lots of experience, in a fume hood or outside with a respirator.

To a 500ml round-bottom flask, I added 36g (0.1549 moles) of trichloroisocyanuric acid (TCCA) and a solution of 95g (0.9233 moles) of sodium bromide in 200ml of water. With addition of the bromide, a  yellow-red colour was observed, likely due to a small amount of chlorine being given off by the TCCA and liberating some bromine. Anyway, I then carefully poured in 34ml (0.1549 moles) of 16% hydrochloric acid, swirled the flask, then quickly attached a stillhead greased with 98% sulphuric acid. The mixture rapidly became very red with addition of the acid. I then attached the rest of the distillation equipment (condenser, vacuum adapter and receiving flask), greasing all the joints with 98% sulphuric acid. A tube was attached to the vacuum adapter and led into a bubbler trap to contain escaping bromine fumes and the receiving flask was placed in an ice bath. I then commenced distillation on medium heat. After a minute or two, I noticed red-orange bromine fumes had filled the whole apparatus and the mixture in the distilling flask became intensely dark red. A few minutes later, bromine began distilling over at a rate of about 1 drop per second. As the distillation continued, more and more water came over with the bromine. I kept distilling, until no more red bromine fumes could be seen in the distilling flask, gradually increasing the heat to keep the distillate coming over.

Once the apparatus had cooled to room temperature, I removed the receiving flask which contained a two layer system of bromine water floating on bromine. I carefully removed the top water layer with a pipette and discarded it. I then washed the bromine thoroughly with 15ml of 98% sulphuric acid to remove any water that might still present . After leaving the layers to separate for a few seconds, I removed the lower bromine layer using a pipette and transferred it to an amber glass vial for storage. A small amount of 98% sulphuric acid was also added to the vial to form a layer above the bromine and help prevent evaporation. The lid was sealed thoroughly with a generous amount of PTFE tape.

I got about 42g (0.2628 moles) of anhydrous bromine, representing a 57% yield. To test the purity, I placed one drop of the product on top of an upside-down beaker and allowed to evaporate. No residue remained, indicating my bromine to be fairly pure.

I should probably mention that my method of storage is not suitable for keeping bromine indefinitely. For permanent storage, ampoules are definitely the best option.

Also, this procedure was by know means designed by myself, I was simply following a brilliant video by the youtube chemist, NileRed (link). It's worth noting that only 1/3 of the total amount of hydrochloric acid needed stoichiometrically is actually required. NileRed seemed to think that this might possibly be because the TCCA is able to function as a source of chlorine directly to some extent.

C3N3O3Cl3 + 3 HCl ==> 3 Cl2 + C3N3O3H3

Cl2 + 2 NaBr ==> 2 NaCl + Br2

Sodium silicate and the chemical garden

Sodium silicate, often referred to as waterglass or liquid glass, is not one compound, but rather a series of compounds with the general formula (Na2SiO2)nO. It is used in high temperature glues for furnaces, in water treatment and in the past, perserving eggs as well as many other things. Sodium silicate also finds use in famous chemical demonstrations known as chemical gardens. In these demonstrations, lumps of transition metal salts are added to sodium silicate solution. The corresponding transition metal silicates gradually begin to form in such a way that an incredible plant-like structure results.

I decided to make some sodium silicate and try the demonstration.

To begin, I added 3.5g (0.0875 moles) of sodium hydroxide to a 125ml beaker. I poured in 12ml of water and swirled the flask. The sodium hydroxide rapidly dissolved, giving a clear solution. I then began adding 5.25g (0.0874 moles) of powdered silica gel in small portions with good stirring, heating the beaker between additions to dissolve everything. After all the silica gel had been added, a small amount of white solid remained which did not dissolve with heating. The solution was filtered to remove this. The resulting slightly yellow liquid should be a fairly concentrated solution of sodium silicate.

The sodium silicate solution prepared above was transferred to a beaker and diluted with 90ml of water. I added a few medium sized pieces of chromium(iii) chloride hexahydrate to the beaker and left the mixture to stand for 12 hours. After the first few hours, I noticed the start of the garden. The chromium(iii) chloride crystals expanded into a plant-like structure of green-black chromium(iii) silicates and a few columns of  the  the material rose up through the liquid, then later collapsed. This was the result after approximately 12 hours:

This is a fairly basic chemical garden. Often a whole range of transition metal salts are added to give colour variation.

2 n NaOH + n SiO2 ==> (Na2SiO2)nO + n H2O   

2 CrCl3 + 3 Na2SiO3 ==> 6 NaCl + Cr2(SiO3)3

8 CrCl3 + 6 Na4SiO4 ==> 24 NaCl + Cr8(SiO4)6

2 CrCl3 + Na6Si2O7 ==> 6 NaCl + Cr2Si2O7