Saturday, 26 January 2019

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

Tuesday, 16 October 2018


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.


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

Wednesday, 1 March 2017

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).


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:


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

Friday, 10 February 2017

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

Saturday, 4 February 2017

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.


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

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

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

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


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

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

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

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

Friday, 27 January 2017

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

Thursday, 12 January 2017

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/cmwhich 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