Sulphur monochloride, revisited

Sulphur monochloride is a reactive chemical used as a precursor to organosulphur compounds. Under standard conditions, it exists as the dimer disulphur dichloride (S2Cl2) which is a yellow-orange fuming liquid with an absolutely horrific smell. It is closely related to sulphur dichloride (SCl2) which is less stable and exists as a cherry red liquid.

Sulphur monochloride is made fairly simply by bubbling chlorine gas through molten sulphur. I did try making some sulphur monochloride in a previous post, however the results were far from satisfactory (link). As a warning, sulphur monochloride is a schedule 3 chemical weapons precursor due to its use in the manufacture of sulphur mustard (Cl-C2H4-S-C2H4-Cl).

Setup

To a two necked round-bottom flask, I added 48.32g of trichloroisocyanuric acid (TCCA). I attached 250ml separatory funnel to the center neck in which I (stopcock closed) poured 73ml of 31% hydrochloric acid. To the second neck of the round-bottom flask, I attached a vacuum adapter containing anhydrous calcium chloride with a stoppered end. To the vacuum inlet on the vacuum adapter, I attached one end of a PVC tube. The other end, I lead into a 1000ml conical flask containing 20g of sulphur. The final setup looked like this:

Reaction

I slowly heated the sulphur in the 1000ml flask, until it melted into a yellow-orange liquid. Then, very carefully I turned the stopcock on the separatory funnel a few degrees so the hydrochloric acid was introduced drop-wise to the flask. yellow-green chlorine gas was produced which quickly filled the flask. I allowed the chlorine to bubble through the molten sulphur in the conical flask at fairly rapid rate. As more and more chlorine was absorbed, the mixture became cherry red in colour due to the formation of sulphur chlorides.

When all the hydrochloric acid in the separatory funnel had been consumed, I allowed everything to cool to room temperature.

Workup and purification

I added the mixed sulphur chlorides to a two necked 500ml flask containing 6g of sulphur to which, I attached a liebig condenser circulated with cold water to the centre neck and a stopper to the side neck. I then refluxed this mixture for 20 minutes. The crude product is a mixture of sulphur monochloride and dichloride, refluxing with sulphur converts the dichloride to monochloride.

 After reflux, I allowed everything to cool to room temperature, then removed the condenser and attached a stillhead to the flask. I reattached the condenser to the stillhead and attached the rest of the equipment nescesery for simple distillation. The receiving flask was well sealed to the vacuum adapter to limit contact of the product with the air. I distilled off the sulphur monochloride till the distilling flask reached a state of near-dryness. I ended up with 10ml of fairly pure sulphur monochloride as a orange-red liquid.

Side note: all equipment that contacted the sulphur chlorides forms a dense coating of sulphur upon washing with water. This is very hard to clean, but hot toluene helps a lot.

C3N3O3Cl3 + 3 HCl ==> 3 Cl2 + C3N3O3H3

S8 + 8 Cl2 ==> 8 SCl2

S8 + 4 Cl2 ==> 4 S2Cl2

Ethyl benzyl ether via Williamson ether synthesis

Ethyl benzyl ether, or a-ethoxytoluene is an ether used in perfumes and as a flavoring. It also finds limited use in organic synthesis. At room temperature, ethyl benzyl ether exists as a high boiling colourless oily liquid with a pleasant fruity smell reminiscent of pineapple. I have no use for ethyl benzyl ether and only made it because I wanted to try making some kind of ether.

The easiest route to the compound is the famous williamson ether synthesis which involves the SN2 reaction between an alkoxide anion and an alkyl halide. For ethyl benzyl ether, ethoxide and a benzyl halide are used. Here is what I did.

To a two necked round bottom flask, I added 32ml of ethanol (reactant and solvent). I then attached a water circulated liebig condenser to the middle neck. Through the side neck, I added in 1.5g (0.216 moles) of lithium metal granules, then sealed the side neck with a stopper. The lithium began gently bubbling in the ethanol and slowly developed a white crust on the outside. To help the lithium react, I began refluxing the mixture. Lots of white precipitate began forming as the lithium was consumed. This should be lithium ethoxide, our nucleophile for the reaction. Once almost all the lithium had reacted and all that remained was one small piece, I added in 28ml (0.243 moles) benzyl chloride (our electrophile for the reaction) through the condenser and continued refluxing for 2 hours. With addition of the benzyl chloride, the lithium ethoxide dissolved and the mixture became yellow, and upon further reflux, orange. A fine white precipitate slowly built up during the reflux, even to a point where the mixture became noticeably viscous. This precipitate should be lithium chloride, the by-product of the reaction which is insoluble in ethanol. After, reflux, I allowed the mixture to cool to room temperature, then filtered it to remove lithium chloride. I then transferred the orange filtrate to a 250ml separatory funnel (stopcock closed), washing the round-bottom and filter flask out with 15ml of ethanol then transferring this also to the separatory funnel. I added in 50ml of water, then capped, shaked and vented the separatory funnel and allowed the layers to separate which took a while since an emulsion had formed. I drained off the aqueous phase, then washed the orange organic phase again with 50ml of water. After draining the lower aqueous phase off again, I transferred the still orange organic phase to a two necked 500ml round-bottom flask containing anhydrous calcium chloride. The calcium chloride unexpectedly dissolved into the water present and formed a small bottom layer. I sealed the side neck of the flask with a stopper and set the middle neck up for simple distillation.

 The first fraction to come over was a small amount of water carrying some ethyl benzyl ether at 92-102 C. The fraction boiling between 185-190 C was collected in a storage vial. This is the ethyl benzyl ether. After distillation, I discovered that quite a substantial amount of the ethyl benzyl ether had been steam distilled over with the water in the lower fraction. So, I separated off the upper ethyl benzyl ether layer and performed another distillation, again collecting the fraction boiling at 185-190 C. This was then combined with the earlier collected product in the vial. Finally I added some anhydrous calcium chloride to the vial to dry the product. In the end, I got 16ml of dry ethyl benzyl ether as a colourless oily liquid with a strong citrus-pineapple smell. This works out to a reasonable 52% yield.

The density of the product was 0.914/cm³ which is quite close to the established value of 0.938g/cm³ suggesting relative purity. However, there is probably some unreacted benzyl chloride present, as I made computation error and accedently used an excess of benzyl chloride.

The lithium metal first reacts with the ethanol, irreversibly generating the ethoxide anion which is a powerful nucleophile. The ethoxide then undergoes a SN2 reaction with the benzyl chloride which is electrophilic producing ethyl benzyl ether.

2 Li + 2 C2H5OH ==> 2 LiC2H5O + H2

LiC2H5O <==> Li (+) + C2H5O (-)

C6H5CH2Cl + C2H5O (-) ==> C6H5C(+)H2(C2H5O)Cl

C6H5C(+)H2(C2H5O)Cl ==> C6H5CH2OC2H5 + Cl (-)

Li (+) + Cl(-) ==> LiCl

Benzyl chloride via nucleophilic substitution

Some time ago, I made some benzyl chloride by the halogenation of toluene (link). The process was on the whole quite unpleasant, very long and gave a miserable yield. It is the primary industrial method for producing benzyl chloride. I deciding to try the more common laboratory method, which is much better suited to small scale operations. Benzyl chloride is the only chlorinated organic compound produced, so the product is generally more pure, in addition the procedure is relatively short.

As a warning, benzyl chloride is highly toxic, lacrymatory, corrosive and potentially carcinogenic. This procedure should only be attempted by an experienced chemist with a good respirator in a well ventilated area. All contact should be avoided.

To a 1000ml flat-bottom boiling flask, I added 575ml of 33% hydrochloric acid. I then slowly added 78ml of benzyl alcohol while swirling the flask slightly. Once addition of the alcohol was complete, I attached a cold water circulated condenser to the flask, with one end of a tube sealed to the end of the condenser. The other end of the tube was suspended just above some sodium hydroxide solution in a separate vessel. This serves as a gas scrubber to neutralize the hydrogen chloride gas that escapes the system. Anyway, I then refluxed the mixture for 10 minutes. Even before the reflux began, an upper layer of clear liquid had separated from the rest of the liquid in the flask.  Throughout reflux, the reaction flask was frequently swirled until the layers merged. After reflux, I allowed everything to cool to near room temperature with the help of an icebath. During this time the layers had completely separated out into a cloudy lower aqueous phase and a clear organic upper phase. The organic phase should be mostly benzyl chloride with some dissolved benzyl alcohol. Anyway, I then poured as much of the mixture into a 250ml separatory funnel as could fit and drained off the lower aqueous layer. The rest of the mixture was added to the separatory funnel in portions, draining off the lower aqueous waste teach time until nothing but the entirety of the organic benzyl chloride layer was left in the funnel. The benzyl chloride was then washed with two 50ml portions of saturated sodium bicarbonate solution in succession, stoppering, shaking, and venting the separatory funnel each time. Note that for these washes, benzyl chloride forms the bottom layer. I then washed the benzyl chloride one last time with 100ml of saturated sodium chloride solution for which, the benzyl chloride formed the upper layer. After draining off the aqueous layer and discarding it, I drained off the benzyl chloride into a 200ml conical flask and dried it over anhydrous calcium chloride.

I then poured the benzyl chloride into a 500ml round-bottom flask and setup for simple distillation. The fraction boiling at 170-182 C was collected. This should be relatively pure benzyl chloride. I then dried the product again over anhydrous calcium chloride and transferred it to an amber glass bottle containing 3A molecular sieves for storage. I got 50ml of benzyl chloride which works out to a respectable yield of 58%. The product was also quite pure with a density of 1.06/cm3!

The reaction is a nucleophilic substitution, which unusually can proceed via SN1 and SN2 mechanisms. Substitution with primary alcohols (of which benzyl alcohol is one) almost exclusively proceed via SN2. This is because the carbocation intermediate involved in SN1 is much less stable for primary alkyls than it is for secondary or tertiary. However there is another factor that can come into play called resonance. Resonance allows the positive charge of carbocations to be slightly spread out across the whole molecule, and thus increase the stability. The aromatic ring in benzyl alcohol creates resonance and so the carbocation that would form in an SN1 reaction is more stable. This is why substitution with benzyl alcohol progresses via both SN1 and SN2.

SN1

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

C6H5CH2OH + H (+) ==> C6H5CH2O(+)H2

C6H5CH2O(+)H2 ==> C6H5C(+)H2 + H2O

C6H5C(+)H2 + Cl (-) ==> C6H5CH2Cl

SN2

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

C6H5CH2OH + Cl(-) ==> C6H5C(+)H2(OH)Cl

C6H5C(+)H2(OH)Cl ==> C6H5CH2Cl + OH (-)

OH(-) + H (+) ==> H2O

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