energetic tetraamminecopper(ii) salts

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

Tetraamminecopper(ii) sulphate

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

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

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

Tetraamminecopper(ii) bis(permanganate)

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

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

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

TACPM powder

Tetraamminecopper(ii) persulphate

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

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

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

TACPS crystals

Tetraamminecopper(ii) dinitrate

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

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

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

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

TACN crystals

Sulphur monochloride, revisited

Sulphur monochloride is a reactive chemical used as a precursor to organosulphur compounds. Under standard conditions, it exists as the dimer disulphur dichloride (S2Cl2) which is a yellow-orange fuming liquid with an absolutely horrific smell. It is closely related to sulphur dichloride (SCl2) which is less stable and exists as a cherry red liquid.

Sulphur monochloride is made fairly simply by bubbling chlorine gas through molten sulphur. I did try making some sulphur monochloride in a previous post, however the results were far from satisfactory (link). As a warning, sulphur monochloride is a schedule 3 chemical weapons precursor due to its use in the manufacture of sulphur mustard (Cl-C2H4-S-C2H4-Cl).

Setup

To a two necked round-bottom flask, I added 48.32g of trichloroisocyanuric acid (TCCA). I attached 250ml separatory funnel to the center neck in which I (stopcock closed) poured 73ml of 31% hydrochloric acid. To the second neck of the round-bottom flask, I attached a vacuum adapter containing anhydrous calcium chloride with a stoppered end. To the vacuum inlet on the vacuum adapter, I attached one end of a PVC tube. The other end, I lead into a 1000ml conical flask containing 20g of sulphur. The final setup looked like this:

Reaction

I slowly heated the sulphur in the 1000ml flask, until it melted into a yellow-orange liquid. Then, very carefully I turned the stopcock on the separatory funnel a few degrees so the hydrochloric acid was introduced drop-wise to the flask. yellow-green chlorine gas was produced which quickly filled the flask. I allowed the chlorine to bubble through the molten sulphur in the conical flask at fairly rapid rate. As more and more chlorine was absorbed, the mixture became cherry red in colour due to the formation of sulphur chlorides.

When all the hydrochloric acid in the separatory funnel had been consumed, I allowed everything to cool to room temperature.

Workup and purification

I added the mixed sulphur chlorides to a two necked 500ml flask containing 6g of sulphur to which, I attached a liebig condenser circulated with cold water to the centre neck and a stopper to the side neck. I then refluxed this mixture for 20 minutes. The crude product is a mixture of sulphur monochloride and dichloride, refluxing with sulphur converts the dichloride to monochloride.

 After reflux, I allowed everything to cool to room temperature, then removed the condenser and attached a stillhead to the flask. I reattached the condenser to the stillhead and attached the rest of the equipment nescesery for simple distillation. The receiving flask was well sealed to the vacuum adapter to limit contact of the product with the air. I distilled off the sulphur monochloride till the distilling flask reached a state of near-dryness. I ended up with 10ml of fairly pure sulphur monochloride as a orange-red liquid.

Side note: all equipment that contacted the sulphur chlorides forms a dense coating of sulphur upon washing with water. This is very hard to clean, but hot toluene helps a lot.

C3N3O3Cl3 + 3 HCl ==> 3 Cl2 + C3N3O3H3

S8 + 8 Cl2 ==> 8 SCl2

S8 + 4 Cl2 ==> 4 S2Cl2

Tin(ii) oxide preparation from pewter, failed attempt

Tin(ii) oxide, or stannous oxide, is a amphoteric oxide use in the maufacture of cranberry glass. It also finds some use as a catalyst for esterifications, however this is not common. It exists in three forms, a hydrated form which is a tan-cream coloured powder, a blue-black powder, and a metastable red powder. In the future, I intend to use tin(ii) oxide in a thermite reaction to produce tin metal.

The inspiration for this experiment came from this patent (link). First, a solution of tin(ii) chloride is prepared. Oxalic acid is then added which precipitates tin(ii) oxalate. The tin(ii) oxalate is then heated with ammonia to give the tin(ii) oxide as the blue-black powder form.

Tin(ii) chloride

To prepare the initial tin(ii) chloride solution, I used pewter, which is an alloy consisting of about 95% tin with the rest being copper, bismuth and antimony. These other components aren't a problem though as they are very unlikely to make it through the process.

To a 300ml beaker, I added 18.95g of pewter (powder and lumps). Using a graduated cylinder, I then added in 66ml of 33% hydrochloric acid and swirled the beaker. A fair amount of bubbling occurred which gradually diminished after a few minutes. I allowed The mixture to stand over night, then covered the breaker with cling wrap and heated the mixture at a low temperature no more bubbling occured, which took about 5 hours. Every now and then the mixture was stirred and a small amount of 33% hydrochloric acid was added to allow for liquid lost by evaporation. a small amount of hydrogen gas was given of at a steady rate for the first few hours, after which the hydrogen became indistigwishable from the bubbles of evaporating liquid. To remove undissolved material, I then filtered the mixture, collecting the clear filtrate in a 250ml beaker. The filtrate should be a roughly 40% solution of tin(ii) chloride. The next step is to convert this to tin(ii) oxalate.

Tin(ii) oxalate

19.12g of oxalic acid dihydrate and 45ml of water were added to a 250ml beaker and heated with stirring to around 60 C, whereupon the oxalic acid dissolved. The tin(ii) chloride solution prepared above was then added in small portions, with stirring in between additions while maintaining the temperature at around 60 C. Complete addition took around 40 minutes. With the first few additions, the mixture slowly became cloudy with fine white precipitate. The precipitate then redissolved towards the end of addition. The slightly yellow, clear solution was then taken off heat and allowed to cool for 1 hour. Beautiful needle-like crystals gradually precipitated as the solution cooled. These crystals should be the product, tin(ii) oxalate. After 1 hour of cooling, The crystals were filtered off and carefully washed on the filter with 100ml of cold water in portions and finally dried, yielding 6.05g of tin(ii) oxalate as white, needle-like crystals.

Tin(ii) oxide (failed)

The 6.05g of Tin(ii) oxalate prepared above was added to a 250ml beaker. 20ml of water was then added with stirring to form a suspension. 4ml of 25% ammonia solution was added and the mixture was heated to 60 C whereupon the tin(ii) oxalate dissolved. The temperature was maintained at 60 C for 40 minutes with occasional swirling of the beaker. After the first 20 minutes, an additional 4ml of 25% ammonia solution was added resulting in a white precipitate with a crystalline structure identical to the starting tin (ii) oxalate. After the 40 minutes of heating, no black tin(ii) oxide had precipitated as claimed by the patent and the experiment was abandoned.

Sn + 2 HCl ==> SnCl2 + H2   /   SnCl2 + H2C2O4 ==> SnC2O4 + 2 HCl

NH3 + H2O <==> NH4OH   /   SnC2O4 + 2 NH4OH ==> (NH4)2C2O4 + SnO + H2O

Ammonium permanganate

Ammonium permanganate is an intriguing salt which at room temperature exists as slightly bronze-metallic purple crystals. It is a moderately strong primary explosive but has very rarely found much use in this field due to its high sensitivity, short shelf life and low power in comparison to the more widely used explosives. I don't really have any use for the compound and mainly just made some because of its interesting composition and to test out its explosive properties.

To make the ammonium permanganate, I followed this procedure (link) which seemed to give very satisfactory results.

To a 150ml beaker, I added 6.18g (0.0391 moles) of potassium permanganate and 17g (0.3178 moles) of ammonium chloride. I then added in 116ml of water using a graduated cylinder and stirred the mixture to dissolve everything. A dark purple solution was obtained and the beaker became intensely cold to the touch. Using a retort stand, I secured the beaker over a water bath heated by a hotplate. Medium heat was then applied and a thermometer was inserted into the mixture to monitor the temperature. Most of the time the temperature of the mixture stayed at around 80 C and care was taken not to let it go much higher then this. The liquid was occasionally stirred with the thermometer. Once the volume of the liquid had been reduced to 74ml, I took the mixture off heat and rapidly filtered it while hot into a 250ml conical flask to remove brown manganese dioxide. This is formed by some of the heat sensitive permanganate ions decomposing.

Anyway, I then chilled the flask containing the filtrate in an ice bath to 5 C whereupon a fair amount of ammonium permanganate precipitated as dark purple felted, short needle-like crystals. The mixture was then immediately filtered to collect the product, pressing it on the filter to remove as much liquid as possible. After drying, I obtained 1.68g (0.0123 moles) of ammonium permanganate which if pure, corresponds to a yield of 31%.

Left = dry ammonium permanganate  /  Right = mildly energetic decomposition of  ammonium permanganate (results in large cloud of manganese dioxide)

The reaction is a relatively simple double displacement with ammonium permanganate being forced out of solution by exploiting its radically reduced solubility in ammonium chloride solution at low temperatures. Hence the large excess of ammonium chloride.


KMnO4 + NH4Cl ==> NH4MnO4 + KCl

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

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

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

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

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

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

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

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

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

Chromium chlorate

Chromium(iii) chlorate is a slightly soluble inorganic salt with the formula Cr(ClO3)3. I chose to make some chromium chlorate because I could find almost no information on it and I was curious about what it would be like.

 I knew copper(ii) chlorate was very soluble, so I thought chromium(ii) chlorate would be as well. I tried adding a solution of chromium(ii) sulphate to a solution of calcium chlorate. I hoped that calcium sulphate would precipitate and I would be left with a chromium(ii) chlorate solution. A precipitate did form, but it was a blue-green color so I knew it couldn't be purely calcium sulphate.

I decided to give chromium(ii) chlorate a break and try making chromium(iii) chlorate.

I prepared a chromium(iii) chloride solution and to it I added a solution of calcium chlorate. A navy blue precipitate of chromium(iii) chlorate formed! The same precipitate formed when sodium chlorate was used in place of calcium chlorate. Anyway I filtered off the chromium(iii) chlorate and dried it. Success!












left = dried chromium(iii) chlorate / right = pyrotechnic with chromium(iii) chlorate and Mg
2 CrCl3 + 3 Ca(ClO3)2 ==> 3 CaCl2 + 2 Cr(ClO3)3

Copper carbonate

Copper carbonate is a nicely colored copper salt used as a pigment. Its formula is often stated to be CuCO3 but Cu3(OH)2(CO3)2 and (Cu2(OH)2CO3) are closer to the truth. I plan to use it to catalyze the decarboxylation of niacin to pyridine, as well as to make calcium copper silicate.

Copper carbonate is relatively easy to make. I made some.

To a 500ml beaker I added 20g of copper sulphate pentahydrate and 150ml of water. The mixture was stirred to give a clear blue solution. With strong stirring, I slowly added 20g of sodium bicarbonate. As I added the sodium bicarbonate, the mixture started foaming and a wonderful green-blue colored precipitate of copper carbonate gradually formed. Once all the sodium bicarbonate had been added, the mixture was filtered to collect the copper carbonate product. After drying, I was left with 31g of copper carbonate.

NaHCO3 + CuSO4 ==> Cu(HCO3)2 + Na2SO4  /  Cu(HCO3)2 ==> CuCO3 + H2O + CO2

Calcium hydroxide

Calcium hydroxide, commonly called slaked lime or caustic lime, is a compound used in cement and to treat sewage. It is slightly soluble in water, and saturated solutions are called limewater. I am interested in using calcium hydroxide as a catalyst for the production of diacetone alcohol.

 Calcium hydroxide can be easily made from sodium hydroxide and calcium chloride. I tried this out.

To a 500ml beaker I added 27g of calcium chloride, then just enough hot water to dissolve it. I set the beaker aside. Then to a 250ml flask I added 20g of sodium hydroxide and a sufficient amount of water to dissolve it. Once the calcium chloride and sodium hydroxide had dissolved into solutions, I added them together. A dense white precipitate of calcium hydroxide formed. I filtered the calcium hydroxide off and dried it in the sun. This yielded 22g of calcium hydroxide in the form of a white powder. The photo was taken before the calcium hydroxide was fully dry.

CaCl2 + 2 NaOH ==> Ca(OH)2 + 2 NaCl