Sciencemadness Discussion Board - Project: Making thionyl chloride from easy-to-get precursors

Project: Making thionyl chloride from easy-to-get precursors

Monoamine - 14-12-2021 at 12:28

Warning: SO₂, HCl and Cl₂ are toxic gasses. SO₃, SCl₂ and SOCl₂ are all extremely corrosive and react violently with water.

This is a set of recipes/notes on how to synthesis thionyl chloride (SOCl₂) from easy to get materials. The project will be organized into a seperate post for each separate step/experiment all within this thread. Participation and discussion is welcomed.

Ever since I learned about it, the reagent thionyl chloride SOCl₂ has fascinated me. It’s very useful as a chlorinating agent, can dry many reagents and is very “clean” to use, in the sense that when it reacts, the byproducts leave the reaction as gases. That being said, it’s also somewhat hazardous because said gases are sulfur dioxide SO₂ and hydrogen chloride (HCl), which is why this reagent is particularly hard to get for the amateur scientist.

Therefore there exists a need for a home-synthesis of thionyl chloride which uses precursors that should be in the reach of most folks as well as being relatively inexpensive. Fortunately, the need for expensive, hard to get equipment and chemicals can be circumvented by anyone with patience and an honest interest in science. In this thread we will carry out a series of experiments aimed at showing that this is indeed possible and how it can be done. The hardest to get chemical in the entire synthesis is probably concentrated H₂SO₄ and - as we shall see - only a catalytic amount is needed.

SOCl₂ is formed when sulfur trioxide (SO₃) reacts with sulfur dichloride (SCl₂). Since we are making everything from easy to get chemicals, the SOCl₂ is synthesized in three steps:

Step 1: Synthesis of SO₃
SO₃ is produced by heating Na₂S₂O₈ with a small amount of catalytic H₂SO₄ to 300C. (Without the catalyst, the temperature required is over 460C).
2Na₂S₂O₈ -> 2Na₂S₂O₇ + O₂
Na₂S₂O₇ -> Na₂SO₄ + SO₃

Step 2: Synthesis of SCl₂:
SCl₂ is produced by generating Cl₂, passing it over elemental sulfur and distilling of SCl₂ as it forms.
Ca(OCl)₂ + 4HCl -> CaCl₂ + 2H₂O + 2Cl₂ (Your method of Cl₂ generation may differ).
2S + Cl₂ -> S₂Cl₂
S₂Cl₂ + Cl₂ <-> 2SCl₂

Step 3: Synthesis of SOCl₂ from SO₃ and SCl₂:
SO₃ + SCl₂ -> SOCl₂ + SO₂

Ingredients and where to get them (the amounts are only approximate):

200ml concentrated sulfuric acid (H₂SO₄). Can usually be ordered online or distilled from some brands of drain cleaner. Note that only 30ml-50ml are actually used up in an ideal synthesis, the rest can be reused.
2kg Sodium persulfate (Na₂S₂O₈). Used in etching copper and electronic circuits. Look in stores where they sell etching supplies.
150g elemental sulfur (S). Sold in some pharmacies and online.
1.5kg calcium hypochlorite Ca(OCl₂). Used as a swimming pool chlorinating agent. 2-3L
conc. hydrochloric acid (HCl). Sold as muriatic acid in swimming pool supply stores.

Essential equipment:
Distillation setup.
Safety equipment: Gloves and eye protection. Gas mask with military grade filter, or fume hood.
Reflux condenser.
Two necked flask.
Hoses.
Heating mantle or hotplate capable of reaching at least 300C.
Heatgun.
Glassware or other material needed to make chlorine.

Recommended equipment:
Drying tube.
Glass funnel.
NaOH for apparatus drying and quenching of acidic fumes.
Full face protection.
Protective overcoat.

Cl2 generator (if Cl2 is made in-house):
Gas wash bottle.
Pressure equalizing addition funnel.
Suckback trap.

Let the fun begin!

[Edited on 14-12-2021 by Monoamine]

Synthesis of sulfur trioxide (SO3)

Monoamine - 14-12-2021 at 13:10

This procedure is adapted from the procedure described by user “Garage Chemist”. His original notes can be found here.
In brief: Sodium pyrosulfate (Na₂S₂O₇) can be formed by heating sodium persulfate (Na₂S₂O₈). Na₂S₂O₇ decomposes above 460C and gives SO₃. The temperature required for this decomposition can be lowered to around 300C by employing a catalytic amount of H₂SO₄.

Previously we attempted this synthesis with a few mishaps. Consult this post for a discussion of what can go wrong and potential pitfalls.

To a 1L round bottom flask (flask A) was added 800g of sodium persulfate (Na₂S₂O₈). To this was added some broken glass pieces as boiling chips and no stir bar, since the SO₃ can actually dissolve the stir bar (learned this during the first trial…). Finally, 15ml of 95% H₂SO₄ was also added to flask A.

This was then set up for simple distillation. But the condenser was NOT cooled with water. The joints of the setup were “greased” with 95% H₂SO₄, by using an old nitrile glove dipped in H₂SO₄ to apply it. This is highly recommended and greatly helps to reduce fumes escaping the apparatus. Flask A was placed into a heating mantle and wrapped in aluminum foil. It was found that not wrapping the flask and still head in aluminum foil makes the process agonizingly slow.

The vacuum takeoff adaptor was connected to an inverted funnel over a 95% H₂SO₄ bath in a beaker in order to dissolve SO₃ fumes and protect the apparatus from moisture. One neat feature of this is moisture trap is that any SO₃ that is dissolved in the H₂SO₄ is converted to H₂SO₄, so when you later go to retrieve the H₂SO₄ in the beaker, you’ll probably have purer H₂SO₄ then you started with!

Flask B was filled with about 50ml of 95% H₂SO₄ to give the SO₃ something to dissolve into. In retrospect, this is probably not too great an idea, because it decreases the yield because the SO3 first has to dehydrate the H₂SO₄.

A thermometer probe that was housed in a glass pipette whose tip had been melted shut was inserted into flask A. The SO₃ will attack any thermometer not made of glass, and non-digital thermometers are usually not able to read temperature in the 300C-350C range, needed here. That being said, using a thermometer isn’t really necessary for this synthesis. Turning the heating mantle to maximum and wrapping the flask in aluminum foil held the temperature at around 320C in the still head. This will however also distill off your catalytic H₂SO₄. The catalytic H₂SO₄ will do its job before distilling away.

The heating mantle was turned to maximum.

At first, at around 100C, a lot of bubbling was observed in the H₂SO₄ bath. This is from evolving O₂ via the reaction:

2Na₂S₂O₈ -> 2Na₂S₂O₇ + O₂

As the temperature in the still head rises to 300C the contents of flask A melts and is converted to SO₃ and Na₂SO₄ via the reaction:

Na₂S₂O₇ -> Na₂SO₄ + SO₃

At this point vapours filled the apparatus and small clear drops of SO₃ began to come over. Some vapour even came out of the upturned funnel in the H₂SO₄ bath. There was also some brown discolouration in the SO₃ that comes over which is due to the SO₃ oxidizing residual organic matter in the glassware.

Occasionally there was a build up of solid SO₃ in the apparatus. If this clogs the setup it can be very dangerous. It never got to that point in this run, but even so, a heat gun was used to warm the glass around frozen SO₃ to melt it into flask B.

After about 3-4 hours the drip rate had become vanishingly slow and the distillation was stopped.

Flask B was stoppered and stored in a, glass jar with a lid, on a bed of dry Na₂CO₃.

The apparatus was not washed with water to avoid the glass from shattering due to residual SO₃ reacting with H₂O and instead was simply left outside under an umbrella to allow the residual SO₃ to slowly react with moisture in the air overnight. It was then properly rinsed out the following morning.

Notes:
Using a simple glass column rather than a condenser is recommended, since it makes heating the column to melt solid SO₃ much easier, and water cooling was never needed.
I only noticed the need for greasing the joints with H₂SO₄ when the distillation had already started, but needless to say, this should be done before hand.
Unless you want to store the SO₄ very long term, I would recommend against putting any H₂SO₄ in the receiving flask, since it decreases the yield quite a bit. (In fact, I think if your aim is making SOCl₂, then just distilling the SO₃ directly into SCl₂ is probably the way to go).
If the synthesis of SCl₂ was performed before this synthesis, then it’s probably a good idea to use the H₂SO₄ used for drying the Cl2 in the acid bath under the upturned funnel. That way the more wet H₂SO₄ from the Cl₂ drying can itself be dried by escaping SO3 and then recycled for later batches. This way the only H₂SO₄ that is ever used up in this synthesis is the catalytic H₂SO₄ used to make SO₃.
Something that is very strange in all of this is that the content of flask A is a liquid even at temperatures as low as 250C and when almost nothing more distills over. At this point the only thing that should still be in the flask is Na₂SO₄ and maybe unreacted Na₂S₂O₇ and a very small amount of H₂SO₄. Considering that there is only a very small amount of sulphuric acid present and the melting points of Na₂S₂O₇ and Na₂SO₄ are all above 400C, this does not make any sense. The content of the flask should be solid… Any ideas why it isn’t?
This procedure also seems to be a relatively easy way to make 100% sulfuric acid.

[Edited on 14-12-2021 by Monoamine]

[Edited on 15-12-2021 by Monoamine]

SWIM - 14-12-2021 at 15:03

I had no idea the pyrosulfate would decompose at such a low temperature with a little added H2SO4.

This looks a lot simpler and less destructive to equipment than the dehydration by metaphosphoric acid.

Synthesis of sulfur dichloride (SCl2)

Monoamine - 15-12-2021 at 10:12

This synthesis is adapted from a youtube video by Extraction&Ire (found here). Also consult the wiki for this synthesis.

This is the synthesis which was maybe the easiest to carry out and produced the largest amount of product. It’s also the step which requires the most pieces of apparatus, but much of it can be home built. (I made two of the gas wash bottles out of rubber stoppers and glass straws and the pressure equalizing addition funnel out of a separatory funnel, a Claisen adapter and some hoses).

The reaction proceeds by passing Cl₂ into sulfur and distilling off the SCl₂ that forms.

To generate chlorine (Cl₂) gas, a Cl₂ generator was put together as follows: To a separate two necked 500ml round bottom flask (flask A) was added about 100g of calcium hypochlorite (Ca(OCl)₂), and to this was connected a 250ml (very clumsily improvised) 250ml pressure equalizing addition funnel that was filled with 32% hydrochloric acid.

The two necked flask was then connected to a suckback trap (flask B) with a hose, since the Ca(OCl)₂ has a risk of foaming over when reacting with HCl. For this reason it is also recommended to use at least a 1L two necked flask to house the Ca(OCl)₂, but I just don’t happen to have one. The suckback trap was then connected to a gas wash bottle (bottle C) containing about 50ml of 95% H₂SO₄, for gas drying.

Bottle C was then connected with a hose to a 500ml round bottom flask (flask D) in a heating mantle. Flask D was filled with 170g of elemental sulfur and a magnetic stir bar. Flask D was then connected to a water-cooled Liebig condenser and a thermometer was inserted into flask D.

Flask E is used to collect our SCl₂, and the gas takeoff adaptor was connected with a hose to a (again very makeshift) bottle containing dry Na₂CO₃ which serves to form a moisture barrier and can also neutralize some of the excess Cl₂ that escapes.

Finally, Flask E was connected with a hose to an upturned funnel in a NaOH bath to further neutralize Cl₂ and any other acidic fumes that are produced in the process.

SCl₂ was then produced as follows:

1) HCl from the addition funnel above flask A is dripped onto the Ca(OCl)₂ in flask A. This generates Cl₂ via the reaction:
Ca(OCl)₂ + 4HCl -> CaCl₂ + 2H₂O + 2Cl₂

2) This generates a pressure gradient towards the “exit” at F, and the Cl₂ is passed through the H₂SO₄ in bottle C, which strips it of moisture and dries it.

3) The dry Cl₂ is then passed into flask D where it reacts with the sulfur to form disulfur dichlorideS₂Cl₂, and – if an excess of Cl₂ is available – sulfur dichloride SCl₂. This happens in two steps via the reactions:
2S + Cl₂ -> S₂Cl₂
S₂Cl₂ + Cl₂ <-> 2SCl₂

4) Note that the formation of SCl₂ is an equilibrium reaction and that it slowly decomposes back into chlorine and disulfur dichloride, especially if sulfur is still present.

5) Luckily the boiling point of SCl₂ is 59.6C and that of S₂Cl₂ is 138C, so SCl₂. Therefore, we distill off the SCl₂ as it forms and collect it in flask E.

6) In the end, this produced about 200ml of SCl₂.

This was carried out, and we re-filled the flask A with fresh Ca(OCl)₂ whenever no more Cl₂ was generated, until no more SCl₂ distilled over. All in all, we used about 800g of Ca(OCl)₂ and 2L of 32% HCl.

We found that adding only a small amount of Ca(OCl)₂ to the Cl₂ generator greatly increases the efficiency of the transformation to Cl₂ and also prevents overfoaming. Unreacted Ca(OCl)₂ can be made to react with HCl in the flask by gently shaking the flask.

Despite trying to quench much of the excess Cl₂ that was made, a great deal still escaped and I stank like a swimming pool for a day. The Cl₂ also completely oxidized the connecting hoses as well as the metal lab stands and Keck clips. So just to re-emphasise, a good way to protect yourself from toxic gasses is essential here.

Discussion: Sulfur dichloride

Monoamine - 15-12-2021 at 10:18

SCl₂ slowly decomposes back into S₂Cl₂ and Cl₂. For this reason, we did not want to seal it into a completely tight container to avoid pressure build up. On the other hand, SCl₂ also react violently with water and fumes as soon as it gets in contact with the moisture in the air. As a compromise, we put it into a glass bottle with a glass stopper and only gently stoppered it, so that pressure could still be released. The bottle was then put into an open beaker on a bed of Na₂ CO₃. This way, if Cl₂ escapes the bottle it will stay in the beaker and be neutralized by the Na₂CO₃ (since Cl₂ is heavier than air). We have now stored the SCl₂ for a week in this fashion and have not had any problems.

Question: I couldn’t find a clear answer as to how SCl₂ reacts with alcohols and carboxylic acids. Does anyone know if it can act as a chlorinating agent in of itself already?

Since this reaction involves a lot of acidic fumes, we surrounded the apparatus with a few concentric rings of litmus paper. This way, as acidic fumes escaped the apparatus, the distance that they travelled and the amount could be visually gauged.

Above you can see that the litmus paper closer to the apparatus is already quite red.

Great! Now we have all the precursors to make thionyl chloride.

clearly_not_atara - 15-12-2021 at 13:35

Very nice work! Especially the SO3 generation -- that's a much more convenient setup than any other I've heard described.

Quote:

I couldn’t find a clear answer as to how SCl2 reacts with alcohols and carboxylic acids.

If it seems obvious that two chemicals react, but nobody describes the reaction, there's a good chance the reaction produces a "complex mixture".

Not stoichiometric:
SCl2 <> S2Cl2 + Cl2
S2Cl2 + ROH >> RCl + S2O (yes this is a real compound)
S2O + Cl2 >> SOCl2 + SCl2 etc

S2Cl2 is more stable and less likely to oxidatively chlorinate the substrate, which I think is why it is preferred for deshydroxy-chlorination reactions.

[Edited on 15-12-2021 by clearly_not_atara]

woelen - 15-12-2021 at 14:04

Very nice result, both the making of SO3 and the making of SCl2!
Heating a glass vessel to well over 300 C for making SO3 still is quite scary, but indeed, this is much more doable than heating dry Na2S2O7 for getting SO3 and can be done with (good quality and reliable) glassware.

Thumbs up!

Monoamine - 15-12-2021 at 23:18

Thank you clearly_not_atara for the tips about reactions of SCl2 with alcohols. Maybe a future project might be trying to see if it can be used to make alkyl chlorides.

DraconicAcid - 15-12-2021 at 23:29

I'm very interested to learn if SCl2 could be used to convert an acid into an acid chloride....

Lionel Spanner - 16-12-2021 at 00:39

Quote: Originally posted by Monoamine

Question: I couldn’t find a clear answer as to how SCl₂ reacts with alcohols and carboxylic acids. Does anyone know if it can act as a chlorinating agent in of itself already?

It mostly seems to be used for making chlorine-substituted thioethers from alkenes, S/N heterocycles from nitriles, and various polymeric species. There needs to be a higher oxidation state at sulphur (IV or VI) to transfer chlorine without transferring the sulphur as well.

Thionyl chloride synthesis (SOCl2) from sulfur trioxide (SO3) and sulfur dichloride (SCl2). Phase I.

Monoamine - 16-12-2021 at 01:30

To make SOCl₂, we followed the synthesis of Georg Brauer on page 382 in “handbook of preparative inorganic chemistry”. A pdf of this wonderful book can be found here.

In brief:
SO₃ is distilled into SCl₂, where it reacts and forms SOCl₂ and SO₂ via the reaction:

SO₃ + SCl₂ -> SOCl₂ + SO₂.

A round bottom 250ml flask (flask A) containing all of our oleum (from experiment 1 and some earlier trials) was set up for distillation. As the receiving flask, we used a 500ml two necked round bottom flask (flask B) connected to a Graham condenser that was fitted with a drying tube containing NaOH. Flask B was filled with all of our SCl₂ (a little over 250ml). (In retrospect, far less SCl₂ should have been used.)

Neither the Graham condenser nor the Liebig condenser were cooled with water. Feeling them periodically, however showed that they were always cold to the touch.

Flask A was heated slowly. At about 135C, a white vapour started blowing into flask B. The formation a kind of light curst in flask B was observed.

At about 150C the first drops of liquid started collecting in the condenser. At this point I realized that I should have greased the joints of the distillery with H₂SO₄ before starting the experiment, because some fumes started to leak. So I quickly pulled the apparatus open and applied some panicky sulfuric acid to the joints to seal them. This worked to stop the fumes from escaping, but I probably also lost some SO₃ in the process, because every time the apparatus is opened, thick white vapours billow out (this is not recommended).

At 175C, a quick and steady drip rate had been achieved. Yellowish solidification was observed in flask B and the temperature in flask B rose to being `warm to the touch’ and some gentle refluxing occurred in flask B.

The reaction seems to be locally exothermic. As soon as a drop of SO₃ hits the SCl₂, it fizzles and a small amount of solid forms (not sure why solid forms??)

When the distillation was done, there was a great deal of yellowish solid in the red liquid. Using a heatgun, the solid could easily be melted into the SCl₂, where it seemed to react with the SCl₂.

The melting point of the solid seems quite low, so maybe some of this is sublimed SO₃? On the other hand, this does not explain why there are also chunks of solid submerged in the SCl₂?? That being said, Brauer also comments that solidification of the consent of flask B is frequently observed in the beginning, but he does not say why this happens or what it is.

Also, at this point the whole shed I was working in completely reeked of fireworks and rotten eggs, which I think is the smell of sulfur dioxide?

Once all of the yellow solid had been melted into the solution in flask B, the colour turned to a darker red and a chunk of solid remained at the bottom of flask B.

In phase two we attempted to separate the SCl₂ and SOCl₂ in flask B by fractional distillation.

References:

Brauer G, editor. Handbook of Preparative Inorganic Chemistry V2. Elsevier; 2012 Dec 2.

Thionyl chloride synthesis (SOCl2) from sulfur trioxide (SO3) and sulfur dichloride (SCl2). Phase II.

Monoamine - 16-12-2021 at 01:38

Next, we stoppered flask B and set it up for distillation to try and fractionally distill off the SCl₂ and any SOCl₂ it contained.

At 35C, a significant amount of fumes smelling like rotten eggs and fireworks (SO₂?) evolved from the drying tube above the collection flask. Holding a solution of 25% ammonia close to the opening of the drying tube showed that these fumes are quite acidic as the ammonia neutralized them and thick grey vapour was created.

The solid in the SCl₂ + SOCl₂ flask also disappeared, so maybe there was somehow SO₃ locked up in the solid in the flask, which then melted into the solution and reacted with SCl₂ to make more SOCl₂?

The boiling point of SCl₂ is 59.6C and the boiling point of SOCl₂ is 74.6C.
Orange distillate started coming over at 50C already (which is 9.6C lower than the boiling point of SCl₂), but judging by the colour, it surely must be SCl₂.

The temperature was increased to above 65C and then varied between 73C and 57C while the orange droplets of SCl₂ continued to come over.

A lot of yellowish solid began to build up near the end of the condenser. I’m actually not sure what this is but it seems to be due to a process in the SCl2 alone since the same build up was also observed at the bottom of the flask in which the SCl₂ was stored. It reacts exothermically and violently with water.

When the distillation was nearly finished it was clear that we couldn’t properly extract the SOCl₂ out of the SCl₂. This is most likely due to several reasons: First of all, since the boiling points of these two liquids is so similar, that separating them is difficult, and we would need a larger amount of SOCl₂ compared to SCl₂ to fractionally distill them and collect pure SOCl₂. The mistake earlier was to use all of our SCl₂, we should have just used a very small amount of SCl₂ and a stoichiometric excess of SO₃.

So this is what I will do next: Use 1kg of Na₂S₂O₈ to produce SO₃ and distill it directly into 50ml SCl₂. In terms of stoichiometry, this should easily produce enough SO₃ to turn all of the SCl₂ into SOCl₂.

Question: Does anyone have an idea why this solidification in flask B happens when SO₃ first drops into the SCl₂, and what this solid might be?

[Edited on 16-12-2021 by Monoamine]

Monoamine - 16-12-2021 at 01:46

Quote: Originally posted by Lionel Spanner

]
It mostly seems to be used for making chlorine-substituted thioethers from alkenes, S/N heterocycles from nitriles, and various polymeric species. There needs to be a higher oxidation state at sulphur (IV or VI) to transfer chlorine without transferring the sulphur as well.

I see. Thank you for the insight Lionel Spanner!

Successful 2 in 1 synthesis: Optimized home-synthesis of SOCl2. Phase I.

Monoamine - 16-12-2021 at 14:43

It worked!

The previous attempt to make SOCl₂ didn’t give a pure product, most likely because we had used far more SCl₂ than SO₃, we will try this again but with a few improvements from the lessons we learned:

1) Distill the SO₃ into the SCl₂ directly. Don’t store it first and definitely don’t distill it into H₂SO₄ first.

2) Use the 100% H₂SO₄ from the previous trials as catalyst. This is optional of course, but it should help to improve the SO₃ yield.

3) Grease all the joints with H₂SO₄ before starting the experiment!

4) Fill the receiving flask with much less SCl₂ than the theoretical yield of SO₄.

We will make SO₄ from 1000g of Na₂S₂O₈. If we use 100% H₂SO₄, then theoretically, one mole of Na₂S₂O₈ should produce one mole of SO₃. So, since the molar mass of Na₂S₂O₈ is 238.03g/mol, we would expect to make 4.20 moles of SO₃.

The density of SCl₂ is about 1.62g/ml and its molar mass is 102.97g/mol. Therefore, 50ml of SCl₂ should contain:

50ml * 1.62g/ml / 102.97g/mol = 1.57mol of SCl₂

In other words, we should expect to have plenty of SO₃ to convert all of the SCl₂ into SOCl₂.

A 1L round bottom flask (flask A in the previous schematic) was filled with 1kg of Na₂S₂O₈ and 20ml of 100% H₂SO₄. The receiving flask (flask B in the previous schematic) was filled with 50ml of SCl₂. A Graham condenser was fitted to flask B and to the end of the condenser was attached a hose leading into a pickle jar filled with NaOH (since my drying tube had become clogged after the previous experiment). Neither condenser was cooled with water. Flask A was then wrapped in aluminum foil and the heating mantle was cranked to the max and heated to around 300C.

As SO₃ started coming over, the coloration of the SCl₂ in the receiving flask became more and more translucent and approached a much lighter orange.

This is to be expected as SO₃ converts SCl₂ into SOCl₂, which is pale yellow to colourless. However, it was clear that we weren’t generating enough SO₃ to complete the reaction. This was almost certainly due to forming the SO₃ at a temperature that was too high (above 300C) and therefore distilling H₂SO₄ out of flask A, thus depriving the Na₂S₂O₇ of its catalyst.

Therefore, over the course of the experiment, 3x30ml of 100% H₂SO₄ were very slowly added to flask A, whenever we had cooled it down enough that the temperature in flask A permitted the addition of more H₂SO₄. This did increase the drip rate, and also helped to generate some additional SO₃, but it was still not enough to fully convert all the SCl₂ into SOCl₂.

Instead, something very strange happened: Rather than becoming more and more translucent, the solution in flask B became darker and darker, going from a dark orange to a blood red burgundy to almost black/brown/violet!

Nice coloration, but not what we want. Or is it...?

In the end, the process was stopped and we collected the dark liquid in the receiving flask.

At this point I thought that the experiment had failed and so I left everything there for two days. But when I went back in a few days later, the air was thick with a delicious biting HCl smell – a very good sign!

Successful 2 in 1 synthesis: Optimized home-synthesis of SOCl2. Phase II.

Monoamine - 16-12-2021 at 14:50

To my surprise, the colour in flask B from before had changed from a dark purple/brown to a very pale translucent orange!

A quick sniff with the nose confirmed that a strong reek of HCl (the telltale smell of SOCl₂) was coming out of the flask . At this point I strongly suspected that there was actually a decent amount of SOCl₂ in the flask, and so I decided to give purification by fractional distillation another try:

Fractional_distillation_setup.png - 2.9MB

The flask with the pale orange liquid was set up for distillation (using the same setup as previously), but this time on an oil bath (both because I’m having trouble with my electricity and because it allows for controlling the temperature more accurately).

As the internal temperature in the flask increased, a huge amount of HCl fumes evolved out of the apparatus and ate through the NaOH in the drying tube in a few minutes, which was so exothermic that the small amount of moisture in the tube began to boil. I couldn’t breathe without a gas mask anymore and just to be safe I held a bottle of 25% ammonia close by so I could visualize where the fumes were.

At about 80C internal temperature, the first drops of distillate came over at a pretty slow drip rate. The temperature was slowly increased to 90C and then to 100C. Very quickly the drops changed from orange to (seemingly) clear, and so the bottles were swapped out since it was clear that no more SCl₂ was coming over.

Over the course of half an hour, the internal temperature was increased to about 150C, at which point a quick and steady drip rate had been achieved.

At this point it was also observed that no more HCl fumes evolved from the apparatus. Presumably the heavy HCl fuming earlier had been from SOCl₂ reacting with residual moisture in the apparatus, and so the lack of HCl fumes indicated that the apparatus was now completely dry.

We kept collecting distillate until no more boiling was observed in the reaction flask and no more distillate came over.

Pleasingly, only about 1-2ml of orange “stuff” (maybe S₂Cl₂?) remained in the reaction flask, and a similar amount of SCl₂ was likely in the first fraction that had been collected.

The rest was a translucent liquid with a slightly yellow tinge: SOCl₂! In the end, about 75ml of SOCl₂ were collected. Considering that we started with 50ml of SCl₂, this is a satisfactory result. More importantly it shows that an almost total conversion of SCl2 had taken place!

1kg of Na₂S₂O₈ with 110ml 100% H₂SO₄ (although this amount almost certainly can be reduced if the temperature is controlled more patiently) together with 50ml of SCl₂, can give 75ml of SOCl₂.

Discussion

Monoamine - 16-12-2021 at 14:52

Well, it seems that this method works and is pretty efficient, in the sense that almost no SCl₂ is wasted. The most resource-intensive part is probably the 1kg of sodium persulfate, but fortunately this is also the cheapest ingredient (at least where I live). I also still believe that someone with more skill than me can probably greatly increase the yield of SO₃, further reducing the quantity of sodium persulfate that is needed.

I now also think I can hazard a guess about my earlier question regarding the build up of solid in the flask B when SO₃ is distilled into SCl₂. I think that this actually is mostly unreacted SO₃. The reason for this guess is that, as was previously remarked, the SOCl₂ smell was only observed in flask B after two days and the colour had also changed dramatically from brown/purple to translucent orange. I think that SO₃ actually reacts with SCl₂ very slowly, much much slower than I thought , and that it first just dissolves into the SCl₂ and then reacts with it over the course of a day or so. At least at ~0-10C ambient temperature (my shed is pretty cold).

Finally, is this all worth it? Is this a viable way to obtain SOCl₂ if you can’t source it commercially? Well, it is a multi-day synthesis, but I think that now that I know how it’s done I can say that it can be done in about 2-3 days. Maybe 1 day for the SCl₂ synthesis, half a day for the SO₃/SCl₂ mixture synthesis, and then (maybe the following weekend) the fractional distillation to purify the SOCl₂. Also, it should not be too difficult to scale this procedure up to produce a pretty sizeable amount of SOCl₂ for only a small increase in cost, so in a few days of work I think you could definitely make a few hundred ml of SOCl₂ like this.

If you just want a few deciliters of SOCl₂ fast, then taking it out of batteries (like this) is probably the way to go, but if you want to make a few year’s worth supply, then using this procedure should be more efficient/cost effective.

[Edited on 16-12-2021 by Monoamine]

[Edited on 16-12-2021 by Monoamine]

teodor - 17-12-2021 at 08:13

Quote: Originally posted by Monoamine

I now also think I can hazard a guess about my earlier question regarding the build up of solid in the flask B when SO₃ is distilled into SCl₂. I think that this actually is mostly unreacted SO₃.

Also, in the presence of air the sulfur oxytetrachloride, ClSO2 * SCl3 could be formed (mp = 57C) and this compound is disproportionating on standing:
ClSO3 * SCl3 -> SO2Cl2 + SOCl2

[Edited on 17-12-2021 by teodor]

Monoamine - 18-12-2021 at 00:09

Hi teodor. Do you know how much air (do you mean O₂?) is needed for this to occur? The reason I ask, is because the only time there is a significant amount of O₂ present is when the sodium persulfate decomposes into sodium pyrosulfate - which is long before any SO₃ is distilled into the SCl₂. Also, while there may be some trace SO₂Cl₂ in the final product, there would only be a very small amount since SO₂Cl₂ decomposes at 100C and the distillation was carried out at 150C internal temperature.

(But maybe decomposing SO₂Cl₂ is what gave rise to the large amounts of HCl fumes that blew out of the apparatus at the beginning of the fractional distillation?)

Some afterthoughts

Monoamine - 19-12-2021 at 07:32

The thionyl chloride was sealed into glass ampules (made from test tubes) for long term storage. The reason for this is that SOCl₂ has a tendency to leak out of most bottles it's in (like bromine). If storing in glass ampules is not an option, then a bottle with a glass or PTFE cap put into a ziplock bag filled with dry Na₂CO₃ is also an option. The sodium carbonate acts as a dessicant and neutralizes escaping acidic fumes, while releasing CO₂, which blows up the ziplock bag and so gives a visual indication of when leakage has occurred.

Some tests were performed on the SCl₂. First, it was found that sulfuric acid and sulfur dichloride are not miscible. (This really surprised me!)

Adding some phosphorus pentoxide (P₂O₅) to the test tube generated SO₃ in-situ and so initiated a conversion of SCl₂ to SOCl₂, which was greatly sped up when the test tube was heated a little.

The result is a bottom layer of H₂SO₄ and a top layer of SOCl₂. The brown tinge is probably contaminants (maybe also from the P₂O₅?)

Does anybody know why SOCl₂ is immiscible with H₂SO₄?

Fery - 19-12-2021 at 08:30

I store SOCl2 in its original bottle, and yes it is permanently slowly leaking. I covered the stopper with few layers of aluminium foil which neutralizes leaks of this nasty compound whicht eats through Al in few months so I have to check the bottle periodically and replace the Al foil.
Btw I suppose your H2SO4 in the bottom layer is 100% otherwise SOCl2 would react with all traces of H2O in your H2SO4. In fact you previously produced oleum so there is no H2O. AFAIK there is no reaction between H2SO4 and SOCl2. SO3 could react with HCl into chlorosulfuric acid but you do not have H2O there, so no HCl produced by hydrolysis of SOCl2.

[Edited on 19-12-2021 by Fery]

teodor - 20-12-2021 at 14:05

Quote: Originally posted by Monoamine

Hi teodor. Do you know how much air (do you mean O₂?) is needed for this to occur? The reason I ask, is because the only time there is a significant amount of O₂ present is when the sodium persulfate decomposes into sodium pyrosulfate - which is long before any SO₃ is distilled into the SCl₂. Also, while there may be some trace SO₂Cl₂ in the final product, there would only be a very small amount since SO₂Cl₂ decomposes at 100C and the distillation was carried out at 150C internal temperature.

(But maybe decomposing SO₂Cl₂ is what gave rise to the large amounts of HCl fumes that blew out of the apparatus at the beginning of the fractional distillation?)

Hi Monoamine.

No, I don't know. In one old book, it was mentioned that this compound is formed when chlorine, which is not free from O2, acts upon SCl2. But H2SO4 and SOCl2 or SO2Cl2 can also give HClSO3, which probably can react with SCl4 which probably could be formed by SCl2 disproportionation to S2Cl2 and SCl4.
It's only a guess and I believe it could be far from the reality, especially if the solid compound is yellow because SClO3 * SCl3 should be colorless.

I know that the presence of as little as 0.2% of SOCl2 prevents the polymerization of SO3. It is used as a stabilizer of the liquid (alpha) form. On the other hand, the presence of even traces of water (1 H2O molecule per 100000 SO3 molecules) cause polymerization. It would be interesting to see what will be different if one adds SOCl2 to the pyrosulfate mixture. But water comes from H2SO4. Is it possible to use chlorosulfonic acid as a catalyst for SO3 production here?

Your assumption that SO2Cl2 is decomposing could be wrong, because

2HClSO3 <-> SO2Cl2 + H2SO4

is an equilibrium at 170C. In this system, many complex reactions occur and the darkening of the mixture is some evidence of the presence of other compounds, probably
with sulfur-sulfur bonds.

But the exact composition probably doesn't matter. The important step is disproportionation which gives sulfur in the 4th oxidation state starting from 2 and 6, and I think this step is essential to get SOCl2, also it can go through some other intermediate compounds. You managed it exceptionally well, so, congratulations!

[Edited on 21-12-2021 by teodor]

[Edited on 21-12-2021 by teodor]

[Edited on 21-12-2021 by teodor]

Monoamine - 21-12-2021 at 04:53

Very interesting idea to use chlorosulfonic acid as a catalyst teodor! I don't know if it also works, but it's certainly worth a try.

Now that I think about it, I'm sure that the final product contains some, since some of the HCl produced must have reacted with some of the SO₃ or even H₂SO₄. Also, if a drop of the final product was placed on paper it immediately blackened it, like oleum.

I wonder if sulfur dichloride would also react with chlorosulfonic acid to make thionyl chloride? But I assume that it would, since chlorosulfonic acid is almost certainly an intermediate of sorts (SOCl₂ and SO₃ would produce it). As you say, there are probably many additional reactions going on at the same time, and SOCl₂ is likely the energetically most favourable product, which is why the reaction tends towards it eventually.

teodor - 21-12-2021 at 07:20

Monoamine,

I have enough quantity of SOCl2 and that is the only reason I don't try to repeat your experiment which is very interesting indeed.

I think you can try to remove chlorosulfonic acid with NaCl or follow a procedure of SOCl2 purification with quinoline or dimethylaniline which I also want to try because for some applications it is better to use a purified compound.

I have an impression, just from the theoretical study of the topic, that SOCl2 in SO3/SCl2 system is not favorable at all and this is why there are only 2 reports of successful synthesis on SM. And it looks like the synthesis has some peculiar point which I think is a disproportionation reaction when some intermediate compound results in SOCl2/SO2Cl2 mixture.
It is different when SO2 is used because it looks like in this case SOCl2 is formed in a pure form. I mean the reaction of SO2 (usually in liquid form) with PCl5, NbCl5, WCl6, and similar higher chlorides.

Fery - 21-12-2021 at 12:29

Hi teodor, I have plenty of quinoline so I can offer it to you for free. More than 10 small bottles of 100 ml then IIRC 1x 1 L bottle and also maybe 1 or 2 bottles of volumes like 250 ml and 500 ml. It is just a little yellow-brown, not colorless anymore. It would require vacuum distillation. All these bottles are somewhat old like 10-30 years. SOCl2 is so nasty that I'm quite afraid to purify it. Whenever possible I would prefer to use it directly from original bottle. Which syntheses require very pure SOCl2?

teodor - 22-12-2021 at 06:14

Yes, Fery, SOCl2 is nasty, for that reason, I will try the distillation after improving my fume hood. There are 2 things I plan to do: change the fan and make additional air inlets from the bottom and behind. I can detect SOCl2 leaks easily but for some other toxic compound the leak detection could be harder, so I prefer to practice with nasty SOCl2 first. If I will manage it I will be able to distill other similarly toxic things in my hood.
If you can give me some quinoline for the distillation it would be great, thank you very much.
I prefer to use pure compounds in many types of reactions, especially when I do experiments or synthesis which I didn't make before, to put the reference in the bottle, then I will be able to compare any other results with my references.

Bedlasky - 22-12-2021 at 10:46

Very impressive work Monoamine! I really enjoyed reading this report. Have you some plans for your thionyl chloride? I have some papers about reactions of some transition metal oxides with SOCl2 or PCl5 if you are interested. You can also make VOCl3 from V2O5 and SOCl2.

Quote: Originally posted by teodor

It is different when SO2 is used because it looks like in this case SOCl2 is formed in a pure form. I mean the reaction of SO2 (usually in liquid form) with PCl5, NbCl5, WCl6, and similar higher chlorides.

Hi, Teodor. May I ask you where did you read this? Usually this goes in opposite way (metal oxide + SOCl2 --> metal chloride (or oxychloride, depending on metal) + SO2).

teodor - 22-12-2021 at 12:55

Quote: Originally posted by Bedlasky

Hi, Teodor. May I ask you where did you read this? Usually this goes in opposite way (metal oxide + SOCl2 --> metal chloride (or oxychloride, depending on metal) + SO2).

Hi Bedlasky. Yes, this is interesting because is not mentioned very often. The first historical mention which I found is in the German book "Die Chemie in Wasserahnlichen Losungsmitteln", Berlin, 1949. It's available from z-lib. If you will not find please U2U me, I will send it to you.

But I read first about this type of reaction in the book "Synthetic Inorganic Chemistry", William J.Lolly which I have in the paper form only. Here is the scan of relevant pages:

Attachment: img-211222214652.pdf (2MB)
This file has been downloaded 283 times

But probably after evaporating SO2 the reaction goes as you said, so may be SOCl2 exists here only as dissolved molecules in SO2.

But your remark is very interesting. Probably all reaction which you have mentioned can go in opposite direction in such conditions.

[Edited on 22-12-2021 by teodor]

Bedlasky - 22-12-2021 at 14:20

Thanks, Teodor!

Maybe there is some sort of equilibrium. For example:

NbCl5 + SO2 <--> NbOCl3 + SOCl2

So when you dissolve NbCl5 in few times excess of liquid SO2, you obtain good yield of SOCl2. But when you rise temperature, SO2 become gas and equilibrium shift to the left. I wonder how they separate SOCl2 from NbOCl3 and NbCl5 without evaporation of SO2.

This sort of equilibrium would exist even at room temperature. Wiki says this:

In the laboratory, niobium pentachloride is often prepared from Nb2O5, the main challenge being incomplete reaction to give NbOCl3. The conversion can be effected with thionyl chloride.

So even at RT, conversion of Nb2O5 to NbCl5 isn't complete. These dehydratation reactions using SOCl2 or PCl5 are often quite complicated, involving various chlorides/oxychlorides at different oxidation states or some addcuts with SOCl2/PCl5/POCl3.

teodor - 22-12-2021 at 14:51

Maybe adducts like Al2Cl6 * SOCl2 could be used to shift an equilibrium during the synthesis. This is also the way to purify the result. The question is, how to break this adduct to make SOCl2 free.

Alucard - 23-12-2021 at 04:03

Quote: Originally posted by Monoamine

It's very strange cause Na2S2O7 melts about 400C, and it was not able to be molten by using gas camp stove (LPG), due to a reason camp stove gives about 350C-380C heat range. I used Na2S2O7 made from NaHSO4 though...

Maybe it's possible to add some 95% H2SO4 to the molten chunk of Na2S2O7 in the flask I have, it is just a one damn piece of salt, but not sure if this will help to melt it at about 300C, but I think I could try to do it.

Not sure it will help anyway, probably doing this will only cause H2SO4 azeotrope will be distilled over solid Na2S2O7 instead.

Monoamine - 23-12-2021 at 13:31

Quote: Originally posted by Alucard

Hi Alucard,

I would give it a try. Adding sulfuric acid to catalyze the decomposition is what's done here and a camp stove should be plenty hot. You'll know when it starts to work when the apparatus fils with dense white vapours. Let us know how it goes

Also, cleaning out the remaining Na₂SO₄ (and probably some sodium pyrosulfate) is a bit tedious. Yes, it will fuse into a solid chunk when it's cooled, so adding some water, heating it until a lot of solid dissolves in the water and pouring it into a bucket a few times will eventually get the flask clean again.

Monoamine - 24-12-2021 at 03:37

Quote: Originally posted by Bedlasky

Very impressive work Monoamine! I really enjoyed reading this report. Have you some plans for your thionyl chloride? I have some papers about reactions of some transition metal oxides with SOCl2 or PCl5 if you are interested. You can also make VOCl3 from V2O5 and SOCl2.

Hello Bedlasky, I was mostly hoping to use it to make acid chlorides and chlorinate alcohols. In terms of inorganic chemistry, trying to make thionyl bromide (by reacting SOCl₂ with dry HBr) might be something too. I don't know much about transition metals, but I'm curious to find out.

Monoamine - 24-12-2021 at 03:48

Quote: Originally posted by teodor

Monoamine,

I have enough quantity of SOCl2 and that is the only reason I don't try to repeat your experiment which is very interesting indeed.

I think you can try to remove chlorosulfonic acid with NaCl or follow a procedure of SOCl2 purification with quinoline or dimethylaniline which I also want to try because for some applications it is better to use a purified compound.

Re distilling at a lower temperature could shed some light on whether there is chlorosulfonic acid in the product. Since their boiling points differ by around 77C this should work fine.

Excess SO₃ in the mixture can be dealt with by adding more SCl₂ and then fractionally distilling off the unreacted SCl₂.

Alucard - 24-12-2021 at 04:26

Quote: Originally posted by Monoamine

To add sulphuric acid as a catalyst I will need to prepare Na2S2O7 as a single step and mill it to powder before continuing. I see no reason to add H2SO4 to a single chunk of Na2S2O7 in the flask, I think it will only cause H2SO4 to be distilled off, especially if H2SO4 will not be able to reach bottom of the flask and will stay over surface of the Na2S2O7.

I only followed garage chemist's post

http://www.sciencemadness.org/talk/viewthread.php?tid=5495&goto=search&pid=121391

but reached about 550C or maybe 600C only, so Na2S2O7 melted and gave bubbling but very rarely aside with a small quantity of white mist, and nothing was distilled off, because there was no enough of heat, as 650C or 700C needed, not 600C...

So I don't know if adding H2SO4 will even work when starting from sodium bisulfate but not from sodium persulfate, I think there might be a big difference, and persulfate is better than bisulfate.

To say the truth it's probably not necessary to deal with oleum to obtain thionyl chloride, because there might be a possibility to prepare thionyl chloride starting from sulfuryl chloride without using SO3, but using exactly SO2Cl2 and S2Cl2.

teodor - 24-12-2021 at 04:28

I am dreaming of getting SOBr2 from the reaction of SOCl2 with BBr3 which is described as good both in purity and yield. Also, I wish one day I will be able to study the group of BX3 compounds including BF3.
But what is interesting, you can accidentally fall to sleep working with SOBr2. It is exceptionally good in delivering Br2 directly to the brain when you have it in the air. And this causes this effect. At least, it is what the literature says.

Monoamine - 24-12-2021 at 08:29

Quote: Originally posted by teodor

I am dreaming of getting SOBr2 from the reaction of SOCl2 with BBr3 which is described as good both in purity and yield. Also, I wish one day I will be able to study the group of BX3 compounds including BF3.
But what is interesting, you can accidentally fall to sleep working with SOBr2. It is exceptionally good in delivering Br2 directly to the brain when you have it in the air. And this causes this effect. At least, it is what the literature says.

And by "asleep" you mean forever. Right?

teodor - 24-12-2021 at 08:37

Quote: Originally posted by Monoamine

And by "asleep" you mean forever. Right?

No, probably until the excess of Br2 will be metabolized by the body. But the danger is probably about those processes you will have going in the laboratory while you are unconscious.

[Edited on 24-12-2021 by teodor]

S.C. Wack - 24-12-2021 at 11:31

Quote: Originally posted by Monoamine

First of all, since the boiling points of these two liquids is so similar, that separating them is difficult

Refluxing with S will convert the SCl2 to S2Cl2.

Quote: Originally posted by Monoamine

I wonder if sulfur dichloride would also react with chlorosulfonic acid to make thionyl chloride?

SCl2 + HSO3Cl -> SOCl2 + SO2 + HCl

Quote: Originally posted by Monoamine

Excess SO₃ in the mixture can be dealt with by adding more SCl₂ and then fractionally distilling off the unreacted SCl₂.

SCl2 + 3SO3 -> S2O5Cl2 + SO2.

If one is going through this much trouble, I suspect that SO3 is better made from SO2, O2, a quartz tube, fire, and part of a catalytic heater from ebay (it's a shame Coleman got out of the business...used models sell for at least twice as much as they did new 15 years ago) or equivalent...would you believe that it might be easier to substitute SCl2 for SO2, soda-lime glass for quartz, and C for Pt?

Monoamine - 25-12-2021 at 03:04

Hi bibliomaster. Thank you for the insights. I'm starting to suspect that pyrosulfurylchloride may actually be a significant constituent of the final product I got.
I'm reading this publication about it atm. (Sorry, I couldn't find a translation). One reason is also that the temperature at which distillate started coming over at a fast rate is right around the boiling point of pyrosulfurylchloride.
Theoretically I would have had about 4x (molar) more SO₃ than SCl₂, I just thought I'd be very inefficient in producing it, hence the excess. Now I wonder if the pyrosulfurylchloride could be induced to decompose to SOCl₂...
But maybe it would be easier to just use 1kg of Na₂S₂O₈ to make SO₃ and distill it into 200ml of SCl₂, rather than 50ml. Then there would not be expected to be an excess of SO₃ and it would be approximately stoichiometrically equivalent.

Monoamine - 25-12-2021 at 03:07

Quote: Originally posted by teodor

Yes, that's a real truth. Don't do chemistry while unconcious. Also, turn off all flames before becoming unconscious.

Bedlasky - 25-12-2021 at 13:32

Quote: Originally posted by S.C. Wack

SCl2 + 3SO3 -> S2O5Cl2 + SO2.

Sorry, but this reaction really doesn't work. You have 9 oxygen on the left and 7 on the right. Don't forget that this is redox reaction.

If your idea is dehydration of SOCl2, S2O3Cl2 is product. However I didn't find mention about this compound anywhere. Which isn't surprising,because tetravelent sulfur doesn't form series of condensed compounds with S-O-S bridge like hexavalent sulfur.

S.C. Wack - 25-12-2021 at 15:25

Add another SO2.

Also, 2SO3 + SOCl2 -> S2O5Cl2 + SO2, and 5SO3 +S2Cl2 -> S2O5Cl2 + 5SO2
(do not bother counting)

This difficulty can however be overcome. DE139455 may interest you:

...Es wurde nun gefunden, daß man viel bequemer und billiger zum Thionylchlorid gelangen kann, wenn man Schwefeltrioxyd auf Einfachchlorschwefel S2Cl2 bei einer Temperatur von etwa 75 bis 80C einwirken läßt.
Der Vorgang verläuft nach folgender Gleichung:
SO3 + S2Cl2 = SOCl2 + SO2 + S.
Sorgt man nun dafür, daß der bei der Reaktion gebildete Schwefel sofort wieder durch einen passend gewählten Chlorstrom zu Einfachchlorschwefel chloriert wird, so erreicht man nahezu die theoretische Ausbeute an Thionylchlorid. Durch das vorliegende Verfahren wird also das lästige Arbeiten mit dem leicht dissoziierbaren Doppeltchlorschwefel und das kostspielige Kühlen durch die Einhaltung der erforderlichen niederen Temperaturen, sowie endlich die Bildung von Nebenproducten, wie Pyrosulfurylchlorid, vermieden. Das Verfahren wird durch das folgende Beispiel erläutert.

Beispiel:
In 1000 kg Einfachchlorschwefel, die sich in einem mit Rührwerk und Rückflußkühler versehenen Kessel befinden, läßt man ebenfalls 1000 kg Schwefeltrioxyd in langsamem Tempo zutreten, indem man die Temperatur auf 75 bis 80C hält und gleichzeitig Chlor einleitet, um den bei der Reaktion gebildeten Schwefel sofort wieder zu Einfachchlorschwefel zu chlorieren und somit der Umsetzung mit SO3 zu entziehen. Nach Beendigung der Reaktion wird rektifiziert...

[Edited on 26-12-2021 by S.C. Wack]

AJKOER - 26-12-2021 at 09:54

My recent thoughts based on a source providing a guide on reaction rates relating to SO3 formation. Please review Page 209, here https://www.google.com/books/edition/Sulfur_in_the_Atmospher... by clicking on the 'Preview' key.

Based on this reference, I no longer, for example, recommend a Singlet oxygen path, but a path based on nitrogen oxide radicals that appears strong.

As to revised ideas, I once noted a seemingly strange path to SO3 cited in an old text that passing SO2 over HNO2 that can reputedly create SO3 (see citation here http://www.sciencemadness.org/talk/viewthread.php?tid=85001#...). However, more recently, I found a reference consistent with such a path. Some ideas as to how to implement follow, starting first with the proposed action of UV light on HNO2:

HNO2 + UV -> •OH + •NO

Or, perhaps better, •NO as a generated product from the decomposition on heating of HNO2, per Wikipedia https://en.wikipedia.org/wiki/Nitrous_acid , to quote:

"Gaseous nitrous acid, which is rarely encountered, decomposes into nitrogen dioxide, nitric oxide, and water:

2 HNO2 → •NO2 + •NO + H2O"

where HNO2 can be created, for example, by the careful step wise addition of H2O2 to NH3 (see https://www.researchgate.net/publication/317692348_Effects_o...) with an approximate reaction (use as a working estimate) given by:

NH3 + 3 H2O2 -> HNO2 + 4 H2O

where even dilute HNO2 here may be acceptable as it can be dehydrated (with CaCl2) or warmed to induce primarily nitrogen oxide product formation. Note: a better (less water) perhaps choice on thermal decomposition at 330 C is Mg(NO3)2 with products, per Wikipedia https://en.wikipedia.org/wiki/Magnesium_nitrate:

2 Mg(NO3)2 → 2 MgO + 4 NO2 + O2

where MgO is a photocatalyst and note the added benefit of oxygen.

Continuing with the presumed action of minor water vapor on SO2:

SO2 + H2O = H2SO3 = H+ + HSO3-

My suggested path to the apparent required radical •HSO3 (per the source link provided below) is:

HSO3- + •NO -> •HSO3 + NO-

the last ion being a precursor to unstable H2N2O2. Note, the reaction rate between HSO3- and •NO2 (from the oxidation •NO) is provided previously on Page 209, Eq 27 in the opening reference.

And finally, per Equation 10 at https://www.ncbi.nlm.nih.gov/pmc/articles/PMC7558052/ , the final reaction with air/oxygen is claimed to be:

•HSO3 + O2 -> SO3 + H+ + •O2-

A confirming source on this path, per Page 43 at https://mymission.lamission.edu/userdata%5Cpaziras%5CChem102... , •NO being apparently a known catalyst for the action of O2 and SO2 leading to SO3.

However, not sure if one can obtain a good amount of SO3 here, but the chemistry is suggestive of any path that conveniently radicalize humid SO2 creating •HSO3 which is further exposed to oxygen. For example, humid SO2 in strong light and air which actually is mentioned in Atomistry on SO3 at http://sulphur.atomistry.com/sulphur_trioxide.html , to quote:

"Sulphur dioxide, when exposed to strong illumination, undergoes partial and reversible conversion into sulphur and sulphur trioxide."

which also suggests adding a strong photocatalyst (like ZnO, as conveniently sourced from zinc acetate at https://nanoscalereslett.springeropen.com/articles/10.1186/s...) where the electron hole (h+) may promote the •HSO3 radical formation:

HSO3- + h+ -> •HSO3

Another, less friendly path, but well known:

"Sulphur trioxide is formed to a small extent together with sulphur dioxide when sulphur or compounds of sulphur are burned in oxygen or air. In an oxygen tomb, in the presence of a compound which on combustion yields water vapour and oxides of nitrogen, for example ammonium nitrate, the combustion goes completely to sulphur trioxide."

Albeit, note: "Ammonium nitrate mixed with sulfur or with metal powders can be exploded by shock. " per https://webwiser.nlm.nih.gov/substance?substanceId=453&i... .
===========================

An interestingly idea on a more speculated path (which I surmised from the opening reference) is to employ the •HO2 radical, albeit, only half as effective as •NO2, but this idea may still provide a workable and safer path.

In the current context, •HO2 could perhaps be conveniently formed from say a solid peroxide like zinc peroxide mixed with riboflavin subject to water vapor and strong fluorescent light (see Reactions (6) to (9) here https://www.tandfonline.com/doi/pdf/10.1080/10473289.1996.10... and also this source: https://www.researchgate.net/publication/6323324_Formation_o... ).

Expected reactions leading to •HO2 are:

ZnO2 + H2O (vapor) = ZnO + H2O2

H2O2 + UV = •OH + •OH

H2O2 + •OH -> H2O + •HO2

and thereby theoretically also some SO3 in the gas stream per cited Reaction (31), which per the cited text may also be photochemical based again with riboflavin as the photocatalyst under strong fluorescent light:

•HO2 + SO2 -> •HO + SO3 (31)

which all occurs apparently at around a reaction rate of 1/2 of the cited reaction rate with •NO2, namely:

•NO2 + SO2 -> •NO + SO3 (27)

where the SO3 would hopefully appear as a fine white deposit with cooling.

Kind of an interesting idea as one would be working without exposure to a very toxic acidic gas or high temperatures used to induce radicalization, albeit all likely slower with a reduced yield.

Note: This approach can be viewed as cited above "Sulphur dioxide, when exposed to strong illumination, undergoes partial and reversible conversion into sulphur and sulphur trioxide" with the added benefit of a peroxide (ZnO2) and also the photocatalyst ZnO created from the ZnO2, all in a very amenable path.

[Edited on 27-12-2021 by AJKOER]

Monoamine - 28-12-2021 at 21:43

Quote: Originally posted by S.C. Wack

Add another SO2.

Also, 2SO3 + SOCl2 -> S2O5Cl2 + SO2, and 5SO3 +S2Cl2 -> S2O5Cl2 + 5SO2
(do not bother counting)

This difficulty can however be overcome. DE139455 may interest you:

...Es wurde nun gefunden, daß man viel bequemer und billiger zum Thionylchlorid gelangen kann, wenn man Schwefeltrioxyd auf Einfachchlorschwefel S2Cl2 bei einer Temperatur von etwa 75 bis 80C einwirken läßt.
Der Vorgang verläuft nach folgender Gleichung:
SO3 + S2Cl2 = SOCl2 + SO2 + S.
Sorgt man nun dafür, daß der bei der Reaktion gebildete Schwefel sofort wieder durch einen passend gewählten Chlorstrom zu Einfachchlorschwefel chloriert wird, so erreicht man nahezu die theoretische Ausbeute an Thionylchlorid. Durch das vorliegende Verfahren wird also das lästige Arbeiten mit dem leicht dissoziierbaren Doppeltchlorschwefel und das kostspielige Kühlen durch die Einhaltung der erforderlichen niederen Temperaturen, sowie endlich die Bildung von Nebenproducten, wie Pyrosulfurylchlorid, vermieden. Das Verfahren wird durch das folgende Beispiel erläutert.

Beispiel:
In 1000 kg Einfachchlorschwefel, die sich in einem mit Rührwerk und Rückflußkühler versehenen Kessel befinden, läßt man ebenfalls 1000 kg Schwefeltrioxyd in langsamem Tempo zutreten, indem man die Temperatur auf 75 bis 80C hält und gleichzeitig Chlor einleitet, um den bei der Reaktion gebildeten Schwefel sofort wieder zu Einfachchlorschwefel zu chlorieren und somit der Umsetzung mit SO3 zu entziehen. Nach Beendigung der Reaktion wird rektifiziert...

[Edited on 26-12-2021 by S.C. Wack]

Oh wow! This is interesting info. Especially since S₂Cl₂ is somewhat easier to make (and more stable to store) than SCl₂! Thank you for sharing!

Alucard - 29-12-2021 at 03:32

Quote: Originally posted by AJKOER

•NO2 + SO2 -> •NO + SO3 (27)

where the SO3 would hopefully appear as a fine white deposit with cooling.

Kind of an interesting idea as one would be working without exposure to a very toxic acidic gas or high temperatures used to induce radicalization, albeit all likely slower with a reduced yield.

[Edited on 27-12-2021 by AJKOER]

I think you can probably burn sulfur in NO2 atmosphere with no O2 or air present, so :

2S + 3NO2 = 2SO3 + N2

If you really like NO2 to be involved...

S.C. Wack - 29-12-2021 at 16:49

Quote: Originally posted by Monoamine

This is interesting info.

BTW len1 was aware of the patent and made SCl2 anyways.
https://www.sciencemadness.org/whisper/viewthread.php?tid=10...

Colleen Ortiz - 25-1-2022 at 01:11

Hello,
This is such an insightful post. I have always found thionyl chloride fascinating myself. SOCl 2 is the chemical formula for thionyl chloride, an inorganic compound. It's a colorless, fairly volatile liquid with a terrible acrid odor. The reaction between sulfur trioxide with sulfur dichloride is the most important commercial synthesis. By heating oleum and gently distilling sulfur trioxide into a cooled flask of sulfur dichloride, this synthesis may be adapted to the laboratory.
SO3 + SCl2 → SOCl2 + SO2
Syntheses from phosphorus pentachloride (The chemical compound PCl5 stands for phosphorus pentachloride. One of the most significant phosphorus chlorides, along with PCl3 and POCl3, is POCl3. As a chlorinating reagent, PCl5 is used. Although commercial samples can be yellowish and polluted with hydrogen chloride, it is a colorless, water-sensitive, and moisture-sensitive solid), chlorine and sulfur dichloride, or phosgene (The organic chemical compound phosphogene has the formula COCl2. It's a colorless gas with an odor similar to freshly cut hay or grass in low quantities. Phosgene is a valuable industrial component, notably for the manufacturing of polyurethane and polycarbonate plastic precursors.) are some of the other ways.
SO2 + PCl5 → SOCl2 + POCl3
SO2 + Cl2 + SCl2 → 2 SOCl2
SO3 + Cl2 + 2 SCl2 → 3 SOCl2
SO2 + COCl2 → SOCl2 + CO2

You should also try new reactions with thionyl chloride further. Thionyl chloride is primarily utilized in the manufacture of organochlorine compounds, which are often employed as intermediates in medicines and agrichemicals. It is frequently favored over other reagents like phosphorus pentachloride because its by-products (HCl and SO2) are gaseous, making product purification easier.

basement - 25-1-2022 at 02:26

I was just pointed to this thread.

Your method of producing SO3 was really neat. I was expecting this whole synthesis project to be based on Leonid Lerner's book and I'm surprised this wasn't mentioned at all.
A very valuable source which also describes this process very similar in detail

I thought it has to be mentioned!

SMALL-SCALE SYNTHESIS of LABORATORY REAGENTS with Reaction Modeling
Leonid Lerner
2011
https://www.routledge.com/Small-Scale-Synthesis-of-Laborator...

[Edited on 25-1-2022 by basement]

[Edited on 25-1-2022 by basement]

Leonid Learner's Book

Monoamine - 29-1-2022 at 14:01

Hi Basement,

Props for mentioning Leonid Lerners book: "Small scale synthesis of laboratory reagents". I have a copy, but I was out of country at the time so I couldn't use it. But I also highly recommend it. It's a really fun book and even goes into some of the mathematics behind reaction modelling.

Now that you mention it. I just saw a very relevant reaction in Lerner's book: It's how to make sodium pyrosulfate from the much easier to obtain sodium hydrogen sulfate (NaHSO₄). THis is from page 180:

2NaHSO₄ -> Na₂S₂O₇ + H₂O, at T = 310^oC.

So you can make sodium pyrosulfate at the same temperatures as the SO₃ from pH down swimming pool additive. This is kind of neat, it means that thionyl chloride can entirely be made from swimming pool supplies. (The things people put in their pools... yikes)

bluamine - 27-8-2024 at 10:40

Quote: Originally posted by Monoamine

Step 1: Synthesis of SO₃
SO₃ is produced by heating Na₂S₂O₈ with a small amount of catalytic H₂SO₄ to 300C. (Without the catalyst, the temperature required is over 460C).
2Na₂S₂O₈ -> 2Na₂S₂O₇ + O₂
Na₂S₂O₇ -> Na₂SO₄ + SO₃

300° C would be hot for borosilicate glass though. And sulfuric acid is reactive to metals. What kind of containers should be used for step one (making SO3)? Did you use quartz glass?

[Edited on 27-8-2024 by bluamine]

[Edited on 28-8-2024 by bluamine]