energetic tetraamminecopper(ii) salts

The complex ion of ammonia and Copper(II) forms a number of energetic salts with strongly oxidizing anions. 

Tetraamminecopper(ii) sulphate

TACS, [Cu(NH3)4(H2O)n]SO4

To a 250ml beaker, I added copper(ii) sulphate pentahydrate (25g, 0.1001 moles), water (25ml) and 25% ammonia solution (40ml, 9.03g NH3, 0.5302 moles NH3). With stirring, almost all the copper(ii) sulphate dissolved giving a rich dark blue coloured solution. I filtered the solution to remove any undissolved copper(ii) hydroxide and copper(ii) sulphate. To a the filtrate, I then added ethanol (38ml), whereupon a large amount of dark blue-purple precipitate formed. This should be the product, tetraamminecopper(ii) sulphate. I filtered off the precipitate, using the filtrate to wash the remaining solids in the beaker onto the filter. Finally I washed the product on the filter with 25% ammonia solution (5ml), ethanol (15ml), then acetone (15ml) and allowed the product to dry.

Yield: 18.64g (76% assuming monohydrate) of tetraamminecopper(ii) sulphate as a dark purple-blue crystalline powder

Tetraamminecopper(ii) bis(permanganate)

TACPM, [Cu(NH3)4(H2O)n](MnO4)2

In a 200ml conical flask, I dissolved tetraamminecopper(ii) sulphate (5g, 0.0203 moles) in water (30ml), giving a rich dark blue solution. I then chilled the solution down to 8 C and added cold saturated potassium permanganate solution (31ml). Without delay, I chilled the mixture down to 2 C, then filtered off the precipitated product immediately, pressing it on the filter to remove as much liquid as possible. The product was then allowed to dry.

Yield: 1.2g (15% assuming monohydrate) of tetraamminecopper(ii) bis(permanganate) as a brown, slightly violet powder

TACPM powder

Tetraamminecopper(ii) persulphate

TACPS, [Cu(NH3)4(H2O)n]S2O8

To a 200ml conical, I added tetraamminecopper(ii) sulphate (5g, 0.0203 moles) and water (25ml). With swirling, the tetraamminecopper(ii) sulphate dissolved, giving a beautiful dark blue solution. Using a graduated cylinder, I then poured in a solution of ammonium persulphate (5.13g, 0.0225 moles, 10% excess) in water (20ml). The flask was swirled, then chilled in a refrigerator for 1 hour, whereupon a mass of dark blue crystals separated. This should be the tetraamminecopper(ii) persulphate. The product was filtered off, pressed on the filter, then allowed to dry.

Yield: 3.37g (48% assuming monohydrate) of tetraamminecopper(ii) persulphate as fine deep purple-blue crystals

TACPS crystals

Tetraamminecopper(ii) dinitrate

TACN, [Cu(NH3)4(H2O)n](NO3)2

To a 500ml, two-neck round-bottom flask, I added copper metal turnings (5g). In the centre neck I placed a double jointed 250ml separatory funnel filled with 19ml of 68% nitric acid (stopcock closed). The 19ml of 68% nitric acid was prepared by by carefully diluting 13ml of anhydrous nitric acid with 6ml of water. To the other neck of the flask, I attached a vacuum adaptor with its other end sealed with a stopper. I then attached one end of a long PVC tube to the vacuum port on the adaptor with the other end of the tube leading to a bubbler trap containing sodium hydroxide solution. The trap serves to neutralize the highly toxic nitrogen dioxide gas produced in the reaction. Anyway, Once everything was ready, I slowly opened the stopcock on the addition funnel and allowed the nitric acid into the round-bottom flask containing the copper. lots of bubbling occured and large amounts of red nitrogen dioxide gas where given off. Once the reaction seemed to be largely complete, I added water (60ml) thourgh the separatory funnel. The green contents off the flask became more blue and most of the remaining nitrogen dioxide in the flask was driven off or dissolved. A small amount of bubbling copper remained so I disconnected the flask and allowed it to stand for a few days to insure the reaction was complete. I then poured the deep blue copper(ii) nitrate solution into a 150ml beaker and allowed approximately 2 thirds of the liquid to evaporate off.

To the concentrated solution, I then added 32ml of 25% ammonia solution. I large amount of heat was given off and a precipitate formed. I stirred the mixture, pressing the precipitate against the sides of the beaker, whereupon it appeared to redissolve. Upon cooling, a large amount of dark blue crystals precipitated. This should be the tetraamminecopper(ii) dinitrate. Once the mixture had cooled completely to room temperature, I filtered off the product, pressing it well on the filter then scrapped it off into a 150ml beaker containing 20ml of acetone. The mixture was stirred well for a few minutes, then the product was filtered off and pressed again. Lastly the product was dried.

Yield: 15.9g (74% assuming monohydrate) tetraamminecopper(ii) dinitrate as dark purple-blue crystals 

TACN crystals

4-Aminophenol via hydrolysis of paracetamol

4-Aminophenol is essentially a benzene ring substituted with a hydroxyl and amino group on opposite sides. Pretty simple huh. It forms colourless crystals at room temperature, but is often impure, appearing orange-red. It finds it's main use as a photographic developer, under the marketing name Rodinal. I decided to try making 4-Aminophenol from paracetamol I'd extracted from pills. The idea and basic guidelines for this synthesis came from Chemplayer, (a fantastic original source of chemistry videos). Whereas Chemplayer performed an alkaline hydrolysis on the paracetamol, I opted to do an acid hydrolysis. This was my only modification.


1) Paracetamol extraction

40 500mg Paracetamol tablets were ground up and added to a beaker. An arbitrary amount of methanol was then added (100-150ml), and the flask was swirled for a while. The mixture was then filtered, and the solids on the filter washed with a bit more methanol. The filtrate was then left to evaporate, leaving the product. An unexpected thing happened after a couple of hours, the methanol solution turned red. The red colour is apparently due to oxidation of the product to brightly coloured benzoquinone imines. This could  be avoided by evaporation in vacuum, leading to a pure colourless product. I didn't do this however, and was left with 15.72g of white-pink brittle paracetamol crystals. probably still fairly pure, lets hope. If so, this is a 79% recovery yield.


2) 4-Aminophenol


To a two-necked 500ml round bottom flask was added 15.72g of paracetamol and 69ml of 5% hydrochloric acid (10% excess of a 1 to 1 molar ratio). The side neck of the flask was stoppered, and the other neck was set up for reflux with a water circulated liebig condenser. The mixture was then refluxed for 1 hour on medium heat. The paracetamol dissolved into solution as the mixture heated up, giving a clear yellow/brown solution which did not change appearance as the reflux progressed. After reflux, the apparatus was dismantled, and the solution transferred to a 250ml beaker (the flask being washed with did of water to ensure all product was transferred). The beaker was then heated till approximately one 1/4 of the liquid had boiled off. The solution (now diminished in volume) was then placed in a fridge, and allowed to cool to around room temperature. A mass of transulest needle like crystals had formed in the flask once cooled (presumebly 4-aminophenol hydrochloride). The solution was stirred, then the pH tested with litmus and found to be acidic (unsurprisingly). I proceeded to make up about 30ml of 20%ish  w/w sodium hydroxide solution in a separate beaker, and added this to the reaction mixture in small portions, checking pH in between additions until the pH was close to neutral. This was done because in acidic conditions, the amino group in 4-aminophenol is protonated, forming a somewhat soluble, undesirable salt form of the product. In basic conditions, 4-aminophenol is deprotonated to give a water soluble 4-aminophenolate. So at neutral pH, just the non-salt form of 4-aminophenol should be present which can be more easily extracted then salt forms, and is the form I wanted. As the sodium hydroxide solution was added, the "4-aminophenol hydrochloride" crystals dissolved and were replaced with a fine off-white precipitate which made the solution thick at a neutral pH. This should be the non-salt, freebase 4-aminophenol. The mixture was chilled in a freezer for ghg hours to allow more product to crystallize out. The product was then filtered off, washed with a bit of ice cold water, and finally dried, giving 3.9g of 4-aminophenol. The product looked pure enough as a tan powder, but never the less, I decided to recrystallize it from water. this turned out not to be such a great idea, as the product is sensitive to atmospheric oxidation when heated. As a result, the final recrystallized product appeared as fine dark drown crystals. if pure, this represents a yield of 34%. Pretty shitty yield and product, but it was fun.

Left = extracted paracetamol  /  Right = 4-aminophenol

For the hydrolysis:  (OH)C6H4NHCOCH3 + H2O  ==> (OH)C6H4NH2 + CH3COOH

with the excess HCl:  (OH)C6H4NH2 + HCl ==>  (OH)C6H4NH3(+)Cl(-)

neutralizing with NaOH:  (OH)C6H4NH3(+) Cl(-) + NaOH ==> (OH)C6H4NH2 + NaCl + H2O

5-Bromovanillin

3-Bromo-4-hydroxy-5-methoxybenzaldehyde, or 5-bromovanillin, is a relatively obscure trisubstituted benzaldehyde used an intermediate in the synthesis of syringaldehyde and other related compounds. At room temperature, it appears as a light yellow crystalline solid with an intoxicating smell rather like vanillin. It can be made with relative ease via low temperature bromination of vanillin.


Method:

Into a  500ml two necked round bottom flask, I poured a solution of 30.42g of vanillin in 200ml of methanol. A separatory funnel (stopcock closed!) containing 35.2g of dry Bromine was attached to one neck of the flask, and the other was stoppered off. The flask was placed in an ice bath (with salt added to decrease temperature in bath further). After leaving contents of flask to cool for a while, the addition of the bromine was begun, a couple of drops at a time , swirling the flask to mix contents each time. Care was taken to make sure the temperature did not rise much over this period. Addition of all the bromine took about 1 hour. After this, I removed the mixture from the ice bath, and stirred it periodically for 1 hour. Mixture was then cooled in another ice bath to around 0 C, whereupon 200ml of water were gradually added (In same way as the bromine, via separatory funnel). The mixture was then filtered to collect the precipitated product, which was washed on the filter with 400ml water, then 200ml 70% methanol solution (0 C). After drying, I obtained 45g of 5-bromovanillin as a fine light yellow crystalline powder, a 97% yield if pure.

C6H3(OCH3)(OH)COH + Br2 ==> C6H2Br(OCH3)(OH)COH + HBr

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

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

Isopropylamine extraction from glyphosate herbicide

Isopropylamine, or 2-aminopropane, is an organic amine used mainly in herbicides and as a chemical intermediate. It is a colourless, highly volatile liquid with the strong rotten-fish smell typical of many amines. There are several reactions I plan to try out with isopropylamine in the future.

Isopropylamine can be fairly easily extracted from glyphosate herbicides. The herbicides contain the amine in the form of the glyphosate salt. The glyphosate is precipitated by adding hydrochloric acid, which leaves the isopropylamine in solution as the hydrochloride salt. Sodium hydroxide is then added to the solution to freebase the amine, which is distilled out of the mixture in fairly pure form.

To a 1000ml beaker, I added 500ml of Number 8 Ready To Go Glyphosate 360 Weedkiller (0.789 moles of isopropylamine glyphosate). With stirring, I poured in 88ml of 33% hydrochloric acid (0.789 moles of hydrogen chloride) and left the mixture to stand for 10 minutes. With addition of the hydrochloric acid, the mixture rapidly became cloudy as a dense white precipitate began to form. After 1 minute, a dense mass of crystals had precipitated. The crystals should in theory consist of N-(phosphonomethyl)glycine (glyphosate), however it's possible that the hydrochloric acid may cause the surfactants in the herbicide to precipitate as well. Anyway, I filtered the heavy white precipitate off, collecting the filtrate in a 1000ml conical flask. 40g (1.00 moles) of sodium hydroxide were then added to the mixture in small portions, with swirling of the flask between additions. A fairly violent reaction was observed with each addition and immediately an intense aroma of rotten fish was noticed. Once the addition was complete, I stoppered the flask and allowed it to cool to room temperature. I then poured the contents of the flask into a 1000ml flat-bottom boiling flask. Thankfully isopropylamine does not form an azeotrope with water, so distillation can be used to extract the amine. A stillhead, condenser, vacuum adapter and receiving flask were attached and a distillation was performed. Boiling water was used as the heat source. The isopropylamine gradually came over at 40-50 C as a colourless liquid.

After a while, I decided that just hot water was not a satisfactory heating source, so I placed a hotplate below the water bath to maintain a temperature of around 90 C. The heat was occasionally adjusted so the vapour temperature stayed at 40-52 C. When no more isopropylamine came over at this temperature, the distillation was ended. I dried the isopropylamine over sodium hydroxide, then redistilled it, collecting the fraction boiling at 30-41 C. Finally, I transferred the product to an amber vial for storage. I got 18ml (0.209 moles) of isopropylamine as the dry freebase. This represents a 26% yield. The density of the product was 0.601g/cm³ which is pretty close to the 0.988/cm³ of pure isopropylamine.

The distasteful smell of isopropylamine wasn't quite as bad as I was expecting. However, the foul odour was still quite unpleasant whenever I was working directly with the compound. Fishy ammoniacal, with a slight rancid note is how I would personally describe it.

The addition of hydrochloric acid at the beginning of the procedure is probably not necessary. I only added this step because I wanted to recover the glyphosate in solid form for easy disposal.


C3H9N.C3H8NO5P + HCl ==> C3H8NO5P + C3H9N.HCl

C3H9N.HCl + NaOH ==> NaCl + H2O + C3H9N

Silicon production from quartz

Silicon (Si) is a chemical element with a wide array of applications. It's used in high strength alloys such as ferrosilicon, but by far the biggest and most well known is its use as a semiconductor in circuit boards. With the rise of computer technology, silicon is playing an ever larger part in the modern era.

I investigated two methods for preparing elemental silicon, the first led to an extremely impure product so I chose not to present it. The second method is a thermite reaction between silicon dioxide and aluminium. Since this reaction isn't thermally self sustaining, a side reaction of sulphur and aluminium is introduced to provide the necessary heat. Many sources of silicon dioxide can be used for the reaction such as silica gel and sand, but I chose to use quartz.

To begin, I added 11.87g (0.1976 moles) of finely powdered quartz, 15.83 (0.4936 moles) of sulphur and 13.19g (0.4888 moles) of 325 mesh aluminium powder to a plastic container. I stirred this mixture intimately until a uniform gray powder was obtained, then poured the powder onto a brick. I placed a piece of magnesium ribbon in the mixture and ignited the ribbon with a torch. A very intense, but fairly slow burning thermite reaction began, giving off large amounts of heat and smoke. Once the thermite was complete, I scraped up the residue and added it to a 500ml beaker containing 200ml of water. After a few seconds, the mixture began giving off a huge amounts of hydrogen sulphide gas, this lasted about 30 minutes. I allowed the mixture to stand with occasional stirring for about 24 hours to insure the evolution of hydrogen sulphide gas was complete. The mixture became a clear liquid with a thick gray precipitate of aluminium hydroxide. Anyway, I stirred the mixture to form a suspension, waited a few seconds, then decanted off the aluminium hydroxide suspension.

With an additional 200ml water added each time, the decantation step was repeated 4-5 times until almost all the aluminium hydroxide had been removed. Left at the bottom of the beaker were a lots of silicon beads, which appeared metallic with a slight bluish luster. I added 50ml of 33% hydrochloric acid and left the silicon to soak in this for 15 minutes with occasional swirling of the beaker. I then decanted off the acid and added 200ml of water to the silicon. I filtered the mixture through a regular kitchen sieve to remove the fine grains of unwanted material and washed the silicon beads caught by the sieve with two 200ml portions of water. After drying, I was left with 3.19g of elemental silicon representing a 57% yield.

Left = silicon beads  /  Right = quartz and aluminium thermite with sulphur

3 SiO2 + 4 Al ==> 2 Al2O3 + 3 Si    /    2 Al + 3 S ==> Al2S3

Al2S3 + 6 H2O ==> 2 Al(OH)3 + 3 H2S   /   Al2O3 + 6 HCl ==> 2 AlCl3 + 3 H2O

Al(OH)3 + 3 HCl ==> AlCl3 + 3 H2O

Benzyl chloride via free-radical halogenation

Benzyl chloride, or a-chlorotoluene, is an extremely useful reagent used in organic synthesis. Commercially, it's mainly used to make benzyl esters, which find use in the perfume industry. Unfortunately, benzyl chloride is an extremely dangerous chemical and has been used historically in chemical warfare due being a strong lachrymator (severely burns eyes, like a tear gas) and very toxic. I plan to use benzyl chloride to make benzylamine.

I made some benzyl chloride using a procedure that I modified slightly (link). Toluene is halogenated directly with chlorine in the presence of a halogen lamp to encourage the equillibrium shift of diatomic chlorine molecules to free radical chlorine atoms. Benzal chloride and benzotrichloride are also produced, but the quanitity of these side products can be kept reletivly low by using a large excess of toluene.

To a 500ml round-bottom flask, I added 100ml (0.9442 moles) of technical grade toluene and 2.5ml of 33% hydrochloric acid. I attached a Liebig condenser to the flask and refluxed the mixture on a medium heat for 2 hours with a halogen lamp placed a couple of centimetres from the flask. During reflux, I slowly added in 36.56g (0.1573 moles) of trichloroisocyanuric acid (TCCA) in small portions through the condensor. Complete addition took 1 hour. Sometimes the small amounts of the TCCA become stuck in the condenser so a long rod was used to push it down into the flask. With the first few additions of TCCA, a considerable amount of heat was given off and the formation of yellow-green chlorine gas was observed. For the second hour of reflux, the flask was ocasionally swirled. Anyway, after the reflux, I turned off the heat and left the mixture to stand overnight. The next morning, I filtered off the precipitated white cyanuric acid, transferring the orange filtrate to a 500ml round-bottom flask. I attached a stillhead and condenser to the flask and began distilling, using a thermometer to track the fractions.

The first fraction to come over was excess toluene at 100-124 C. I collected around 35ml of this. The vapour temperature then rose and I began to collect benzyl chloride at 160-186 C. After all the benzyl chloride had come over, there was nothing left in the distilling flask and I ended the distillation. I then redistilled the benzyl chloride, this time collecting the fraction boiling at 160-182. The residue in the boiling flask should consist of mainly benzal chloride and benzotrichloride. Anyway, I dried the benzyl chloride over anhydrous calcium chloride, then transferred it to an amber vial for storage. I got 16ml (0.139 moles) of benzyl chloride which is a 29% yield. The product had a density of 0.921g/cm³ which isn't too far away from the theoretical 1.1 g/cm³. This shows the benzyl chloride I made is probably moderately pure.

The lachrymatory properties of benzyl chloride were actually nowhere near as bad as I was expecting. My eyes were attacked viscously when I was washing the glassware and briefly during the initial reflux but that was about it.

The pour yield could likely be dramatically improved by using a larger excess of toluene. Adding the TCCA more slowly probably would have also helped to increase the yield. I did notice there were a lot of higher chlorination products left over in the distilling flask after the second distillation.

C3N3O3Cl3 + 3 HCl ==> 3 Cl2 + C3N3O3H3

C6H5CH3 + Cl2 ==> C6H5CH2Cl + HCl

Iron(iii) oxide

Iron(iii) oxide, otherwise known as ferric oxide, is a brick red compound used in the product of iron and steel. It also finds use as a pigment. iron(iii) oxide happens to be the principle constituent of rust and this is probably how most people encounter the compound. Just for fun, I decided to try making some.

18.3g (0.0658 moles) of iron(ii) sulphate heptahydrate were dissolved in 75ml of water in a 250ml beaker. In a separate flask, I dissolved 11.06g (0.1316 moles) of sodium bicarbonate in 110ml of water. Since the dissolution of sodium bicarbonate is endothermic, slight heating was needed to dissolve everything. I then added the iron(ii) sulphate sulphate to the sodium bicarbonate solution in small portions. A white precipitate of iron(ii) carbonate formed and the mixture foamed as carbon dioxide was given off. After stirring for a few minutes, I filtered the mixture. By this time, the iron(ii) carbonate precipitate had turned green and in some places red.

Anyway, I dried the filtered off iron(ii) carbonate, which became completely red. Whats happening is the iron(ii) carbonate, which is a white solid, slowly reacts with oxygen in the air to form iron(iii) oxide which is red. Once the material was completely dry, I transferred it to a crucible and gassed it with a butane torch for 10 minutes to make sure the conversion to iron(iii) oxide was complete. During heating, the colour became a darker, richer red as everything was oxidized to iron(iii) oxide. Once the iron(iii) oxide had become uniformly red, I transferred it to filter, washed it on the filter with 50ml of boiling water then finally dried it. I got 3.23g (0.0202 moles) of iron(iii) oxide. The product is probably not pure and likely contains some Iron(ii,iii) oxide.

2 NaHCO3 + FeSO4 ==> Fe(HCO3)2 + Na2SO4   /   Fe(HCO3)2 ==> FeCO3 + H2O + CO2   /

2 FeCO3 + O2 ==> Fe2O3 + CO2   /   FeCO3 ==> FeO + CO2   /   2 FeO + O2 ==> Fe2O3

 4 FeO ==> Fe3O4 + Fe   /   4 Fe + 3 O2 ==> 2 Fe2O3   /  4 Fe3O4 + 2 O2 ==> 6 Fe2O3

Nitrostarch

Nitrostarch is not a specific compound, but rather a mixture of different esters of nitric acid and starch. It's a secondary explosive and was used as a filler in grenades in world war 1. The energetic properties of nitrostarch are very similar to nitrocellulose, both consisting of chains of glucose molecules with varying degrees of nitration.

Nitrostarch is easily prepared by the mixed acid nitration of starch. Normally a mixture of nitric and sulphuric acids is used, but I chose to use potassium nitrate and sulphuric acid which forms nitric acid in situ.

To begin, I poured 36ml of 98% sulphuric acid into a 250ml beaker in an ice bath. Once the temperature of the acid had dropped below 10 C, I began adding 11.2g of potassium nitrate in small portions with good stirring. During the additions, the temperature was never allowed to rise above 20 C. Once all the potassium nitrate had been added, I added 10g of corn starch to the mixture in one portion with good stirring. The mixture became quite viscous and turned slightly yellow. I stirred the mixture continuously for 15 minutes, then left it to stand for 2 hours with occasional stirring, maintaining the temperature at below 10 C. After this, I began cautiously adding 400ml of water to the reaction mixture. Once a small amount of water had been added, the mixture became fluid enough to be easily poured, so I transferred the mixture to a 1000ml flask before adding the rest of the water. I left the mixture to stand until all the nitrostarch had settled to the bottom of the flask, then decanted and discarded as much of the cloudy supernatant liquid as I could without losing any nitrostarch. I then added a further 400ml of water to the nitrostarch and repeat the above process (letting then nitrostarch settle then decanting as much of the supernatant liquid as possible).

Again, I added 400ml of water to the nitrostarch that was left but this time also added in 3ml of 100% undiluted nitric acid. I boiled the mixture for 10 minutes with good stirring, covering the top of the flask to help prevent loss of the volatile nitric acid. Still heating, I then neutralized the mixture with sodium bicarbonate until the liquid became neutral-slightly alkaline. By mistake I sometimes added a bit to much bicarbonate and the mixture bubbled over a couple off times. I think this may have led to fairly substantial loss of product. Anyway, after the neutralization, I filtered off the nitrostarch, washing it on the filter with 1000ml of water. After drying, I was left with 2.62g of nitrostarch as a white powder with a slight yellow colour.

Left = nitrostarch   /   Right = deflagration of nitrostarch

KNO3 + H2SO4 ==> KHSO4 + HNO3

HNO3 + 2 H2SO4 <==> 2 HSO4 (-) + NO2 (+) + H3O (+)

5-Sulphosalicylic acid

5-Sulphosalicylic acid is an organic compound possessing both sulphonic and carboxylic acid functional groups. It's used in integral colour anodizing, however its main use is as a reagent to measure protein levels in urine. Under standard conditions, 5-sulphosalicylic acid exists as a white crystalline solid.

5-sulphosalicylic acid is easily prepared by the sulphonation of salicylic acid with concentrated sulphuric acid. I decided to try this synthesis out.

To a 500ml round bottom flask, I added 3.71g of salicylic acid. I carefully added 18ml of 98% sulphuric acid and stirred to thoroughly mix the two chemicals. I then set the flask up for reflux with a Liebig condenser. I refluxed the mixture for 1 hour and 30 minutes using a boiling water bath as the heat source. After a few minutes of reflux, the suspended salicylic acid dissolved, giving a blood-red liquid, which gradually turned brown. Once the reflux was complete, I allowed the mixture to cool to room temperature. A small amount of solid material separated. I then, very slowly and with rapid stirring, poured the mixture (including the solid material) into a 250ml beaker containing 140ml of saturated sodium chloride solution. Initially not much seemed to happen, but after a few seconds, a fine white precipitate formed. This is the product, 5-sulphosalicylic acid dihydrate.

I filtered the 5-sulphosalicylic acid off, washing it on the filter with two 15ml portions of cold water. I then dried the 5-sulposalicylic acid. This product however did not seem to be very pure at all, so I recrystallized the 5-sulphosalicylic acid from saturated sodium chloride solution. After recrystallization, I was left with 3.68g of fairly pure 5-sulphosalicylic acid dihydrate as a fine white crystalline powder. This represents a 54% yield.

C6H4(OH)COOH + H2SO4 ==> C6H3HSO3(OH)COOH + H2O

Copper metal, two ways

Copper is possibly the oldest metal known to mankind. It was the first metal cast into a shape and even the first to be smelted from ore. I thought it might be interesting to try some different approaches to making copper metal from its compounds. There are many ways it can be done, and in this post I present two methods that I investigated.

The first involves the displacement reaction between magnesium and copper(ii) sulphate.

To a 150ml beaker, I added 11.55g of copper(ii) sulphate pentahydrate and 50ml of water. With a bit of swirling, the copper sulphate dissolved, giving a nice blue-coloured solution. I then added in 1.07g of powdered magnesium metal in portions, over the course of about 2 minutes. A vigorous and exothermic reaction immediately began. The mixture started bubbling, giving off steam and releasing a large amount of heat. Gradually, the mixture turned a mud-green colour, which became a brown and finally an orange-brown. At this point it became evident that the colour was caused by small particles of copper metal in suspension. Once all the magnesium had been added, I left the mixture to stand until all the orange-brown copper particles had settled to the bottom of the beaker. I then decanted as much of the liquid off as possible without losing any copper. To the residual copper and liquid that was left, I added 50ml of 5% acetic acid.

I gave the mixture a stir, then left it to stand for about 10 minutes. Then I filtered the mixture, washing the copper on the filter with two 50ml portions of water. And finally I dried the copper powder. 1.95g of copper powder was collected, corresponding to a 70% yield.

The other method I investigated was a thermite reaction between copper(ii) oxide and magnesium.

3.77g of copper(ii) oxide and 1.15g of powdered magnesium metal were added to a mortar and pestle and ground into a fine powder. The powder was then spread on a brick and ignited with a butane torch. A violent thermite reaction began, giving off large amounts of fire and smoke. Once the reaction was complete, I collected the residue and added it to a 250ml beaker. I then added 30ml of 16.5% hydrochloric acid to the beaker and swirled this for a few minutes. I filtered the mixture and transferred the copper on the filter back to the 250ml beaker. I added an additional 30ml of 16.5% hydrochloric acid, stirred the mixture for a few minutes, then filtered off the copper, and washed it on the filter with 50ml of water. After drying, this gave 0.09g of very dirty copper metal, which represents a miserable 3% yield.

Left = copper made by the first method    Right = Copper(ii) oxide thermite from the second method

CuSO4 + Mg ==> MgSO4 + Cu  /  CuO + Mg ==> MgO + Cu

2 CH3COOH + Mg ==> Mg(CH3COO)2 + H2  /  2 HCl + CuO ==> CuCl2 + H2O

2 HCl + MgO ==> MgCl2 + H2O  /  2 HCl + Mg ==> MgCl2 + H2