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

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

Nitric acid

Nitric acid, or hydrogen nitrate, is a very common laboratory acid with a wide array of uses. One of the main uses is for nitrating various compounds. There are many explosives based around nitro groups, which makes nitric acid incredibly useful as an explosives precursor. I plan to use nitric acid to make 4-nitrotoluene which in turn I will eventually use to synthesize a local anesthetic called benzocaine.

Nitric acid is usually supplied at an azeotropic concentration of 68%. However the nitric acid produced in this procedure is almost pure.

To a 1000ml round-bottom flask, I added 101g of powdered potassium nitrate and 54ml of 98% sulphuric acid. I then quickly set the flask up for simple distillation with the receiving flask in an ice bath. I then began distilling the mixture at a medium heat, gradually increasing the temperature whenever the yellow nitric acid stopped coming over. At the start of distillation, the flask was full of red nitrogen dioxide gas. This quickly faded to a yellow once the nitric acid started coming over. Gradually increasing the temperature throughout distillation is needed to free the nitric acid from all the sulphate salts formed. Pure nitric acid initially comes over at about 70-80 C then later on, azeotropic nitric comes over at over 100 C. However the vapour temperature was never allowed to rise above 100 C, so when nothing came over below this, the distillation was ended.

I collected the yellow nitric acid in the receiver and transferred it to a glass bottle for storage. I got 38ml of practically pure nitric acid which corresponds to a 91% yield.

Such concentrated acid is very dangerous, and despite what one might expect, its best to wear no gloves at all when handling it since anything above 87% nitric acid can ignite latex and nitrile.
The yellow colour of the nitric acid is due to a small amount of decomposition to nitrogen dioxide. This really isn't a problem though, as the contamination is very minimal.

KNO3 + H2SO4 ==> KHSO4 + HNO3

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

Boron nitride

Boron nitride is an interesting compound used in ceramics and cosmetics. It's moderately heat resistant and can withstand temperatures up to 2,800 C in the absence of oxygen. Boron nitride occurs in several different forms, which are structurally very similar to the allotropes of carbon. Graphite, diamond and lonsdaleite each have boron nitride analogs.

There are several good ways to make boron nitride. I tried a couple of them out. The reaction of urea with boron trioxide I found to be by far the most reliable, so this is the one I decided to present in this post.

The first step is to convert boric acid to the required boron trioxide.

To a metal can, I added 16g of boric acid. I then strongly heated the can over a camping stove. The boric acid began to melt and give off steam. The mixture became a glassy bubbling syrup as the boric acid was converted to boron trioxide. After about 15 minutes, the bubbling had almost completely stopped and I allowed the mixture to cool. Upon cooling, the boron trioxide solidified into a brittle glass-like substance. I scraped the product out of the can and was left with 8.6g of boron trioxide.

With boron trioxide prepared, I moved on to making the boron nitride.

I ground up the 8.6g of boron trioxide (made above) with 20g of urea. I then poured this mixture into a metal can, which I heated over a camping stove for 30 minutes with gradually increasing heat and fairly constant stirring. The mixture formed a bubbling melt, which diminished after a few minutes to a solid mixture. If the heat is increased too rapidly, a significant amount of boron trioxide can fail to react. I made this mistake, so after the 30 minutes of heating, I added an additional 10g of urea and continued heating for 10 minutes. This probably wouldn't have been necessary had I not started heating so strongly at the beginning.

Anyway after this, I allowed the mixture to cool to room temperature. I then crushed the mixture up and added it to a 150ml beaker containing 75ml of dilute hydrochloric acid. I left this to soak overnight. The next morning most of the white boron nitride had settled to the bottom of the beaker. Above it there was a thin layer of brown material. I stirred the mixture up, waited until the most of the boron nitride had settled, then carefully decanted and discarded as much of the supernatant liquid-brown material suspension as I could. I then washed the residual boron nitride with 100ml of cold water. I filtered off the boron nitride, washed it on the filter with 20ml of ethanol, then dried it. I was left with 2.15g of white boron nitride which is a 35% yield.

The procedure I was following (here) called for boiling the product in water to purify it. However I've found in previous runs that this tends to completely destroy the product. It's well known that hot water reacts with boron nitride, so I'm not sure why they suggest this method of purification.

2 B(OH)3 ==> B2O3 + 3 H2O    /   B2O3 + (NH2)2CO ==> 2 BN + 2 H2O + CO2

B2O3 + 3 C2H5OH ==> (C2H5)3BO3 + B(OH)3

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

Trimyristin from nutmeg

Trimyristin, or glyceryl trimyristate, is an interesting triglyceride that occurs naturally in nutmeg. It doesn't have many uses except in scientific research. Trimyristin can be hydrolysed to glycerol and myristic acid, and in the future I plan to try this out. Ground dry nutmeg typically contains about 25% trimyristin (by weight) so this is a great source of the compound.

The extraction of trimyristin from nutmeg is a classic experiment. Usually diethyl ether is used as the extraction solvent, but I tried using a different method.

To a 250ml conical flask, I added 40g of dry ground nutmeg. To this I added 100ml of hot (nearly boiling) ethanol and stirred vigorously for 15 seconds. I then immediately filtered the mixture through a cloth, collecting the orange filtrate in a 200ml conical flask. I washed the nutmeg on the filter with 25ml more hot ethanol. It's important that the whole filtering process is performed quickly while the ethanol is still hot, otherwise the trimyristin will start crystallizing out. Anyway, the 200ml conical flask containing the filtrate (already starting to precipitate some product) was chilled to about 2 C. A lot more trimyristin crystallized out and the mixture was filtered to collect it.

After filtration, I was left with fine crystals of almost white trimyristin. I decided to perform a recrystallization from ethanol to purify my product. After recrystallization, I was left with 0.84g of trimyristin as very fine white crystals. The product was easily melted by hot water, which is a good sign as trimyristin is said to melt at about 56 C. My guess is the trimyristin isn't extremely pure but still definitely usable for most things.

Left = recrystallized trimyristin   Right = ground nutmeg

This extraction was poorly planned and performed. I probably could have gotten more product by boiling the nutmeg in ethanol at the start. My procedure was loosely based on this.

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

Sodium pyroantimonate

Sodium pyroantimonate is an interesting salt used as a glass clarifier. It's also used in monochrome picture tubes and glass fibers. At room temperature, sodium pyroantimonate exists as white crystals.
I chose to make it because the synthesis seemed interesting. I probably won't use the product for anything though.

To a 125ml beaker I added 20ml of 35% (by weight) sodium hydroxide solution. Next I added 4g of crude antimony trisulphide. The antimony trisulphide dissolved giving a brown-yellow solution. Slight heating was needed to get all of it dissolved. Since my antimony trisulphide was impure, some elemental antimony remained. This was filtered off before continuing. The solution was transferred to a 1000ml conical flask and diluted with 130ml of water. I heated the mixture up to 70 C and slowly added 50ml of 6% hydrogen peroxide in portions with stirring while maintaining the temperature at 70-85 C. The colour of the mixture changed to a light yellow and some crystals of sodium pyroantimonate precipitated. I let the mixture cool down to room temperature then filtered off the sodium pyroantimonate.

 I washed the product with 10ml of equivolume water/ethanol mixture and then 10ml of anhydrous ethanol, then dried it. I got 1.5g of dry sodium pyroantimonate.

Sb2S3 + 6 NaOH ==> Na3SbS3 + Na3SbO3 + H2O

3 H2O + Na3SbO3 + H2O2 ==> NaSb(OH)6 + 2 NaOH

 NaOH + H2O2 + Na3SbS3 + 3 H2O ==> NaSb(OH)6 + 3 NaSH

Preparation of mixed polysulphanes

Polysulphanes are interesting compounds. At room temperature they exist as toxic yellow liquids. The three stable polysulphanes are disulphane, trisulphane and pentasulphane. They are all very sensitive to alkalies so the glassware used to make these compounds must be washed with acid to remove trace amounts of alkali. They have few applications and uses.

I decided to try making disulphane. The reaction produces all three polysulphanes. In a future post I will perform a distillation to isolate disulphane.

The first step is to prepare a solution of sodium polysulphides.

To a 500ml beaker, I poured a solution containing 17g of sodium hydroxide and 150ml of water. I added in 20g of sulphur and began boiling the mixture. The sulphur gradually dissolved and the mixture turned a beautiful dark red colour. After all the sulphur had dissolved, 120ml of dark red liquid were left. This is the polysulphide solution.

Next the sodium polysulphides must be acidified to yield the polysulphanes. All equipment used in this procedure was washed with 5% acetic acid to prevent decomposition of the polysulphanes.

In a 500ml beaker, I chilled down 20ml of the solution prepared in step one to 0 C. The reaction is best performed at -15 C but 0 C was as cold as I could get. After this temperature had been reached, I added the solution to 60ml of 33% hydrochloric acid (also chilled to 0 C). If the starting solution and the acid aren't chilled, only hydrogen sulphide and sulphur will be formed. Anyway, after the addition, a bit of hydrogen sulphide was produced along with some solid sulphur. I stirred the mixture then let it settle. After this, the polysulphanes could be seen as a vivid yellow liquid at the bottom of the container separate from the surrounding liquid. 

The product was collected with a syringe and placed in an acid-washed vial. I got 0.7ml of mixed polysulphanes.

S8 + 12 NaOH = 4 Na2S + 2 Na2S2O3 + 6 H2O  /  S8 + 4 Na2S = 4 Na2S3

S8 + 8 Na2S = 8 Na2S2  /  S8 + 2 Na2S = 2 Na2S5  /  Na2Sx + 2 HCl = H2Sx + 2 NaCl

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 (+)

Synthesis of boron

Boron is an interesting element used in high strength fibers. It burns with a nice green flame. I plan to use it to make boron tribromide. I tried making some boron from boric acid.

The first step is to convert the boric acid to boron trioxide.

To a metal can, I added 15g of boric acid. Then I heated the can on a medium heat. After 5 minutes the boric acid had began to melt and decompose. Ten minutes after this, the mixture had turned to a bubbling sticky mess. I kept the heat on for another 10 minutes. After this, a sticky glassy mass remained in the can. Upon cooling, it solidified into rock-hard crystals of boron trioxide. The boron trioxide was scraped off the can and collected.

The final step is to reduce the boron trioxide to elemental boron.

I ground up the boron trioxide crystals from step one into a fine powder. This was very difficult and took several hours. Then I added a roughly equal amount of magnesium powder and cuttings to the boron trioxide. I blasted the mix with a butane torch until all the mixture had turned black. I was expecting the mixture to act like a thermite and not need to be torched. I think the reason this didn't happen was because half my magnesium was not powdered. Anyway, after this I added the black residue to a beaker and added about 50ml of water. I then slowly added 12M hydrochloric acid until the acid stopped affecting the mixture. The mixture fizzed and lots of gas was produced during the addition.

I noticed a disgusting smell. This is most likely borane gas. Once the mixture had stopped bubbling, I added 400ml of water to dissolve any residual boron trioxide. Then I filtered the mixture to collect the boron product. I got 0.06g of boron as a black powder, which is a miserable 1.2% yield.

I blame the extremely low yield on the thermite not working. The butane torch probably didn't do a very good job.

2 B(OH)3 ==> B2O3 + 3 H2O  /  B2O3 + 3 Mg ==> 3 MgO + 2 B

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

Thorium dioxide from electrodes

Thorium dioxide (ThO2) also known as thoria is a weakly radioactive compound used in some welding electrodes. Thorium gives off alpha radiation which can't penetrate paper and is very short ranged, so it's only dangerous if inhaled or consumed. Thorium dioxide used to be used in gas mantles but this was stopped due to concerns about the radiation. I plan to use thorium dioxide to make Thorium metal and generally explore thorium chemistry.

Thorium dioxide can be extracted from thoriated tungsten electrodes used in welding.
These are usually 2-4% thorium dioxide.

Two years ago I brought 4 thoriated tungsten electrodes. I added them to a container of 3% hydrogen peroxide and left them there until the electrodes had corroded away.  After this there were lots of yellow crystals were left. I collected the crystals and placed them in a vial.

Yesterday I decided to purify the thorium dioxide. I realized my sample was probably a mix of thorium dioxide and tungstic acid. So I added the the yellow crystals to 5ml of saturated sodium carbonate solution and heated strongly for a few seconds. The yellow tungstic acid began to react away and a white powder (presumably thorium dioxide) remained. Once this had stopped, I filtered the mixture to collect the white thorium dioxide. Then I added the thorium dioxide to another 5ml of sat sodium carbonate solution to make sure all the tungstic acid had gone. After this, I filtered off the thorium dioxide and dried it. I got 0.8g of slightly yellow thorium dioxide.

 

 left = the thoriated electrodes (image from www.ecvv.com)  /  Right = thorium dioxide

H2WO4 + Na2CO3 ==> Na2WO4 + H2O + CO2

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

Sodium cyanate

Sodium cyanate (NaOCN) is a reagent used to prepare ureas and organic isocyanates. Despite the name, it's not very toxic. It is possible to reduce sodium cyanate to the infamous poison sodium cyanide, but I won't be doing this for obvious reasons. I'm interested in using sodium cyanate to synthesize semicarbazide.

I prepared sodium cyanate from sodium carbonate and urea.

To a metal pot I added 13.4g of anhydrous sodium carbonate. I heated the pot strongly for 10 minutes to ensure the sodium carbonate was truly dry. After this, I added 5.75g of urea with stirring.
The urea almost immediately melted and decomposed. The urea addition must be performed outside because large amounts of toxic ammonia gas are produced. Anyway, I added 2 more 5.75g portions of urea separately with stirring. After all the urea had been added (and the mixture was dry) I added 90ml of water. Glacial acetic acid was added until the mixture reached pH 6. I filtered the mixture, collecting the clear filtrate. Then I added 100ml of ethanol and stirred.

A fine precipitate of sodium cyanate formed. I collected the sodium cyanate via filtration and dried the product. 2.43g of sodium cyanate was obtained.

(NH2)2CO ==> HOCN + NH3   /   2 HOCN + Na2CO3 ==> 2 NaOCN + H2O + CO2

Sodium bromate

Sodium bromate is an interesting oxidizer used in industrial batch dyeing. Weirdly enough, solutions of it can dissolve gold so it's sometimes used in gold mining. Sodium bromate must be handled with care as it's a possible carcinogen. The easiest way to make sodium bromate is using electrochemistry. I tried this method out.

To a 300ml beaker I added 50ml of saturated sodium bromide solution and a pinch of sodium dichromate. I then electrolyzed the mixture for 3 hours and 10 minutes. The power source was 12 volts DC at 24 amps. A carbon anode and a steel cathode were used. Anyway, after the electrolysis, the mixture was filtered to remove carbon particles and then chilled down to 10 C. Nice white crystals of sodium bromate precipitated as the mixture cooled. I filtered the mixture again to collect the sodium bromate. Then I dried the product. I was left with 1.23g of sodium bromate.

left = dried sodium bromate    right = pyrotechnic with sodium bromate and Mg

anode:  2 Br- ==> Br2 + 2 e-   /   cathode: H2O + 2 e- ==> 2 OH- + H2

Br2 + 2 OH- ==> BrO- + Br- + H2O  /  3 BrO- ==> BrO3- + Br-  /   Na+ + BrO3- ==> NaBrO3