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

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

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

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

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

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

Preparation of nitrourea

N-nitrourea, or N-nitrocarbamide, is an interesting compound with a variety of uses in organic synthesis. It is also a powerful secondary explosive and thus must be handled with caution. However, I'm not so interested in nitrourea as an explosive, but rather as a precursor to semicarbazide.

Nitrourea can be made by the dehydration of urea nitrate. The reaction is a bit finicky, and straying too far from the ideal temperature range can result in a significant loss of product.

To 150ml beaker, I added 17ml of 98% sulphuric acid. I placed the beaker in an ice bath (an arbitrary amount of sodium chloride was added to the ice bath to get it even colder). Once the temperature of the sulphuric acid had dropped to -2 C, I began adding 11g of urea nitrate in very small portions with constant stirring, at such a rate that the temperature never rose above 0 C. Complete addition of the urea nitrate took about 30 minutes. The viscosity of the mixture gradually increased and fine white particulate began to form. I continued stirring for a few minutes after all the urea nitrate was added.

Then I began adding pieces of ice (one at a time) to the mixture. This generates a fair amount of heat, so its best to start with small pieces as the nitrourea product is sensitive to warm-hot water. As more and more ice was added, lots of white nitrourea precipitated. Once the volume of the mixture had reached 125ml, I filtered off the nitrourea and washed it on the filter with 70ml of ice cold water to help remove sulphuric acid.

The product turned out to be very difficult to dry. I let it sit in a desiccator bag over anhydrous calcium chloride for over a week, after which it was still sticky from moisture. I decided this was as dry as I was going to get it, so I weighed the product and calculated the yield anyway. I got 4.33g of slightly wet nitrourea as a sticky white powder. Yield: 46%.

The nitrourea should be sufficiently pure and usable for most reactions. For maximum purity, the nitrourea can be recrystallized from ethanol. However I decided not to do this.

(NH2)2CO.HNO3 == -H2O ==> NH2CONHNO2

Urea nitrate

Urea nitrate, or more correctly, uronium nitrate, is an intriguing organic compound consisting of a protonated urea cation coordinated to a nitrate anion. It's perhaps most well known for its use as a high explosive, for which it has gathered an unpleasant reputation. I should probably point out that I'm not interested in using urea nitrate as an explosive, but rather to make nitrourea and eventually semicarbazide.

Urea nitrate is usually synthesized by the direct combination of urea and nitric acid. However I decided to use another method that avoids the use of expensive (for me anyway) nitric acid.

To a 200ml conical flask, I added 27g of potassium nitrate and 16ml of water. With a bit of warming, most of the potassium nitrate dissolved, giving a milky solution. I then added in a solution of 15.85g of recrystallized urea in 20ml of water. The resulting solution was cloudy, probably due to small amounts of undissolved urea and potassium nitrate in suspension. I stirred the mixture well, then added 30ml of 33% hydrochloric acid in portions with frequent stirring. This took a few minutes. With the first addition, a little bit of fizzing occurred, but after this, the mixture became clear and the later additions had no visible effect. Next I heated the mixture up until it began to steam but not bubble, then let it cool to room temperature.

I then chilled the mixture to 0 C and a large amount of crystalline urea nitrate precipitated. I filtered off the urea nitrate crystals, pressed them on the filter to expel as much liquid as possible, then dried them. I was left with 22.8g of urea nitrate as colourless-white crystals. This is a 70% yield.

(NH2)2CO + HCl ==> (NH2)2CO.HCl    /    (NH2)2CO.HCl + KNO3 <==> (NH2)2CO.HNO3 + KCl

Myristic acid synthesis

Myristic acid, or tetradecanoic acid, is a fatty acid with very few common uses. It is however, a useful organic building block. Under standard conditions, myristic acid is a low melting white crystalline solid. I'm not sure if I'll use my myristic acid for anything, although maybe I'll make some myristyl alcohol or a myristate ester.

Myristic acid can be made by the alkaline hydrolysis of trimyristin. This is a fairly easy and straightforward synthesis.

To a 500ml round-bottom flask, I added 1.43g of trimyristin and a solution of 11.43g sodium hydroxide in 48ml of water. I then added in 45ml of ethanol and attached a Liebig condenser to the flask. I refluxed the mixture for 1 hour and 30 minutes on a medium heat, adding in an additional 10ml of ethanol after the first 45 minutes. The trimyristin dissolved quickly and the liquid turned red, but after this, there wasn't much visible change. Once the reflux was complete, I let everything cool to room temperature. I then poured the mixture into a 200ml conical flask. At this point, the mixture consists of mostly sodium myristate dissolved in water. To get the myristic acid, the sodium myristate must be acidified with a strong acid. So I added 27ml of 33% hydrochloric acid. The red colour lightened to a yellow and the mixture became cloudy. I did add in a little bit more hydrochloric acid by accident, which was completely unnecessary.

Anyway, after a few minutes, a white precipitate of myristic acid began to collect on the bottom of the flask. I chilled the mixture down to get as much myristic acid to crystallize out as possible. A bit of unreacted trimyristin floated on top of the liquid and this was carefully decanted off. I then filtered off the myristic acid, washed it with 50ml of water, then dried it. I got 0.36g of myristic acid, which is a 27% yield.

I'm honestly not sure why the yield was so low.

(C13H27COO)3C3H8 + 3 NaOH ==> 3 NaC13H27COO + C3H8(OH)3

NaC13H27COO + HCl ==> C13H27COOH + NaCl

Oxidation of toluene to benzoic acid

Benzoic acid is an aromatic carboxylic acid used as an organic building block. It occurs naturally in a wide range of plants, and is used as a food preservative. It's also an ingredient in many cosmetics. I plan to use benzoic acid to synthesize Benzamide and eventually aniline.

Benzoic acid can be made very easily from toluene. I did a test run to see if I could get this to work.

To a 250ml beaker I added 15g of potassium permanganate and 157ml of water. With a bit of stirring, most of the potassium permanganate dissolved giving a dark purple solution. I then added this solution (and the small amount of undissolved potassium permanganate) to a 500ml round-bottom flask. Next I added 35ml of technical grade toluene and attached a liebig condenser to the flask. I then strongly refluxed the mixture for 2.8 hours. The mixture gradually darkened as the potassium permanganate was converted to brown manganese dioxide. After 2.8 hours of reflux, I allowed the mixture to cool to room temperature. The purple colour of permanganate had completely gone, indicating all potassium permanganate had been consumed. I filtered the mixture to remove the manganese dioxide by-product. Manganese dioxide is a useful chemical to have, so I kept mine rather then discarding it.

 Anyway, I transferred the clear filtrate to a 250ml beaker, a small layer of excess toluene floated on top of the surrounding liquid. I removed the toluene with a syringe and saved it for future runs. After this, I was left with about 125ml of liquid in the beaker. This is a solution of potassium benzoate in water. I slowly added 33% hydrochloric acid to the solution and fluffy white crystals of benzoic acid precipitated. I kept adding hydrochloric acid until no more precipitate formed.

I then filtered off the benzoic acid and dried it. I got 2.45g of benzoic acid as fine fluffy white crystals.

The reason for the small amount of product obtained was my scales which were broken and gave a false reading.
Due to this I actually used a lot less than 15g of permanganate. I have no idea how much permanaganate I used so I can't calculate the real yield.

2 KMnO4 + C6H5CH3 ==> KC6H5COO + 2 MnO2 + KOH + H2O

KC6H5COO + HCl ==> C6H5COOH + KCl

Ammonoylsis of diethyl oxalate to oxalamide

Oxalamide or ethanediamide, is an organic compound used as a fertilizer. It's also used as an additive in some rocket fuel mixtures to help slow the burning rate. At room tempurature, oxalamide exists as a white solid. When heated past 350 C, it decomposes releasing deadly cyanogen gas. I did plan to use oxalamide for an experimental synthesis of cyanoformamide, but I decided that this was too dangerous.

Oxalamide can be made by the ammonoylsis of diethyl oxalate. This is a quick and very easy synthesis.

First I added 24ml of 25% aqueous ammonia solution to a 100ml beaker. Then using a syringe, I added 10ml of diethyl oxalate. At first the diethyl oxalate just formed a bottom layer, but after a few seconds, I noticed a white precipitate (presumably oxalamide) beginning to form. I began stirring the mixture and rapidly more oxalamide precipitated. The mixture got thicker as the precipitate built up. After about 1 minute of stirring, the mixture had become a thick white paste. Once this point was reached, I transferred the pasty mixture on to a sheet of paper to dry. After drying, I was left with 6.42g of a white powder, which is almost definitely oxalamide. If pure, this represents a 99% yield!

 A nice feature of the reaction is that oxalamide is the only non-volatile solid involved. This means a fairy pure product can (in theory) be obtained just by allowing the mixture to dry. The only thing that could contaminate the product is unreacted diethyl oxalate. For this reason, I used a 10% excess of ammonia solution to ensure all the diethyl oxalate reacted.


(C2H5)2C2O4 + 2 NH3 ==> (CONH2)2 + 2 C2H5OH

Diethyl oxalate

Diethyl oxalate, or diethyl ethanedioate, is a interesting ester used as an organic building block. It has an ethereal smell similar to grape juice. Diethyl oxalate is also used as a specialty solvent for various resins, and as a pigment carrier. In the pharmaceutical industry, diethyl oxalate is used to produce barbiturate steroids. I plan to use it to prepare ethyl acetopyruvate and oxalamide. Diethyl oxalate can be made from dry ethanol and anhydrous oxalic acid. I tried this out.

To get a good yield, the hydrated oxalic acid used must be made anhydrous.

To a large crystallizing dish, I added 100g of oxalic acid dihydrate. I then placed the crystallizing dish
on a hotplate and lowered in an overhead thermometer. I heated the oxalic acid at around 140 C until no more water boiled off. During this process, the oxalic acid melted and a large amount of vapor
(mostly steam) was boiled off. This is not a very effective way to produce anhydrous oxalic acid, but I got there in the end.

To a 500ml round-bottom flask, I added the anhydrous oxalic acid prepared above. I then added 200ml of ethanol and a drop of 98% sulphuric acid, then set up for reflux. I allowed the mixture to reflux for 2 hours. The anhydrous oxalic acid dissolved but apart from this, there was not much visible change in the mixture. After 2 hours of reflux, I removed the condenser and set up for simple distillation. The first fraction came over at 71-81 C,  consisting of ethanol/water azeotrope. About 175ml of this was collected. The temperature of the distillate then rose to 100 C, and I changed the receiving vessel. I kept collecting distillate until nothing was left in the distilling flask. The temperature had climbed to 220 C at the end of the distillation. I had collected 50ml of clear liquid, predominantly diethyl oxalate.

To this liquid, I added 60ml of saturated sodium chloride solution. I shook the mixture up then let it settle. The diethyl oxalate separated, forming a layer above the surrounding liquid. I collected the diethyl oxalate product with a syringe, dried it over anhydrous calcium chloride and finally transferred it to an amber glass bottle for storage. I got 42ml of diethyl oxalate which is a 34% yield.


Adding a drop of sulphuric acid isn't absolutely required as oxalic acid is a strong enough acid to catalyse the reaction.

H2C2O4 + 2 C2H5OH ==(- H2O)==> (C2H5)2C2O4 + 2 H2O

2-Aminophenol via hydrolysis

2-Aminophenol, or 2-hydroxyaniline, is an organic compound used in photography. It's a valuable reagent used to synthesize many heterocyclic compounds such as benzoxazoles (many of which are important in the pharmaceutical industry). I plan to use 2-aminophenol to synthesize catechol. Although I'm not sure if I'll ever get around to it.

2-Aminophenol is formed via the acid hydrolysis of 2-benzoxazolinone. I tried this out.

To a 500ml round bottom flask, I added 27ml of 33% hydrochloric acid and 70ml of water. I then added 1.8g of 2-benzoxazolinone and set up for reflux. Using a boiling water bath as the heating source, I gently refluxed the mixture for 2 hours. The 2-benzoxazolinone gradually dissolved leaving a clear brown-yellow solution. After the reflux, I allowed the mixture to cool down to room temperature. As the mixture cooled, fluffy white crystals of 2-aminophenol began to precipitate.

Once the mixture had fully cooled, I filtered off the 2-aminophenol product. After drying I was left with 0.62g of slightly brown 2-aminophenol which is a 42.6% yield.

A note on safety: although the deadly gas phosgene is produced in this reaction, it immediately reacts with the water present as soon as it is formed. Due to this, the risk of phosgene poisoning is extremely low.


C6H4(O)NHCO + 2 HCl ==> C6H4(OH)NH2 + COCl2  /  COCl2 + H2O ==> CO2 + 2 HCl

1,3,5-Trimethylbenzene

1,3,5-Trimethylbenzene commonly called mesitylene, is an aromatic hydrocarbon occasionally used as a ligand or solvent. Its also used as an additive to plane fuel and as a precursor to 2,4,6-trimethylaniline. Upon mild oxidation with Manganese dioxide, mesitylene forms 3,5-Dimethylbenzaldehyde, I may have a post on this in the future. Mesitylene can be made by the dehydration of acetone with sulphuric acid. I tried this method out.

To a 1000ml round-bottom flask in an ice bath I added 127ml of acetone. Once the acetone had cooled to 5 C, I (over the course of 5 hours) began adding 90ml of 98% sulphuric acid in small portions. I made sure the temperature never rouse above 22 C during the additions. Towards the end of this, the mixture started becoming red in colour. I left the mixture to stand for 24 hours then set up for simple distillation. as the distillation progressed a green liquid started coming over. Soon after this, huge amounts of sulphur dioxide started being generated. The distillation setup must be equipped with a gas outlet to dispose of the toxic sulphur dioxide. Anyway, after a few minutes, a yellow oil began to come over. it formed a layer on top of the previous distillate. I presumed this was the crude mesitylene.

I stopped the distillation when the foam in the distilling flask reached the bottom of the still-head.
Using a syringe I collected the top mesitylene layer in the receiving flask. I washed the mesitylene with 20ml of dilute sodium hydroxide solution, then dried it over anhydrous calcium chloride.
I then distilled the crude product collecting the fraction boiling at 100-200 C. As I was collecting the mesitylene distillate I accidentally spilled some. Due to this I only got 3ml of mesitylene product and my yield was 2.1%. The procedure I followed claimed a 25% yield.

3 (CH3)2CO  ==H2SO4==>  C6H3(CH3)3 + 3 H2O

Nitrocellulose

Cellulose hexanitrate, commonly called nitrocellulose, is a very interesting explosive. When ignited unconfined, it bursts into a large ball of flame. This happens so fast that it's actually safe to ignite it on your hand. When the nitrocellulose is tightly confined it detonates, resulting in a violent explosion. Nitrocellulose was first prepared by Henri Braconnot in 1832. The synthesis of nitrocellulose is a chemistry classic.

Nitrocellulose is usually prepared from sulphuric and nitric acids, however I substituted the nitric acid with potassium nitrate.

To a beaker in an ice bath, I added 50ml of 98% sulphuric acid. I let the acid cool for 5 minutes, then over the course of 30 minutes added 25g of potassium nitrate in portions. This forms the required nitric acid in situ. After the nitrate addition, using a glass stir rod, I pushed pieces of cotton wool in until no more could be covered by the mixture. I left this overnight. The cellulose in the cotton reacts with the acids to form nitrocellulose. Anyway the next morning, I removed the nitrocellulose and added it to a bucket of water. I left it in the bucket for about 15 minutes to soak, then swirled it around in the water until it uncaked and started looking like cotton wool again.

I washed the nitrocellulose with 500ml of saturated sodium bicarbonate solution to remove traces of sulphuric/nitric acid. Then I washed it with water again to remove the sodium salts from the product.
I then dried it. I couldn't weigh the nitrocellulose as the pieces were too big to fit on my scales.

Left = dry nitrocellulose  /  Right = nitrocellulose deflagration unconfined
H2SO4 + KNO3 ==> HNO3 + KHSO4  /  HNO3 + 2 H2SO4 ==> 2 HSO4 (-) + NO2 (+) + H3O (+)

Acetone semicarbazone

Acetone semicarbazone is an organic compound derived from semicarbazide. It's got quite an interesting structure. I plan to use acetone semicarbazone to make semicarbazide.

I made acetone semicarbazide using a synthesis from www.prepchem.com which I modified slightly.

To a 250ml conical flask, I added 2.29g of hydrazine sulphate, 2g of sodium carbonate and 2ml of water.

After stirring, I poured in a solution of 1.14g of sodium cyanate in 25ml water along with 2ml of acetone. I stirred the mixture well for 10 minutes then left it to stand for 24 hours. After this, I filtered the mixture, collecting the filtrate. I evaporated the filtrate down to dryness, then to it I added 50ml of ethanol. The original procedure called to use a soxhelt extractor but I don't have this. Anyway I stirred the mixture so the ethanol could dissolve as much product as possible. Then I filtered the mixture again collecting the ethanolic filtrate.

I dried the filtrate in a crystallizing dish and was left with a pathetic amount of acetone semicarbazone. The amount was so small that I didn't bother weighing it. So I'm not sure what the yield was.

I think a larger quantity could have been obtained if I had used a soxhelt extractor or a least used hot ethanol. Warming the mixture before the cyanate addition would probably have helped as well.

see here for how this reaction works

Ethyl bromide

Ethyl bromide, also called bromoethane, is a volatile organic used as a ethylating agent. It's a very useful reagent to have. I plan to use ethyl bromide to make tert-pentanol and ethylamine. Ethyl bromide must be handled with care as it's a suspected carcinogen.

Ethyl bromide is easily prepared from ethanol, sulphuric acid and sodium bromide. I chose to use this method.

To a 1000ml round-bottom flask I added 74ml of ethanol and 50ml of water. I placed the flask of ethanol and water in an ice bath and slowly over the course of 30 minutes added 80ml of 98% sulphuric acid. I took care to ensure the mixture didn't heat up too much during the addition. After this, I added 70g of sodium bromide and set up for simple distillation. In the receiving flask I placed 25ml of water. I began distilling the mixture at a medium heat. Ethyl bromide began distilling over. When no more ethyl bromide passed over, the distillation was stopped.

The ethyl bromide had formed a layer underneath the water in the receiving flask. The ethyl bromide layer was collected with a syringe and washed with sodium bicarbonate solution. Then it was loaded into a watertight bottle and 25ml of sodium bicarbonate solution was added on top of the product to prevent it from evaporating. The lid was sealed with duct tape.
I got about 25ml of ethyl bromide which is roughly a 33.5% yield.

C2H5OH + H2SO4 ==> C2H5HSO4 + H2O  /  2 NaBr + H2SO4 ==> Na2SO4 + 2 HBr

C2H5HSO4 + HBr ==> C2H5Br + H2SO4