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