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Author: Subject: Everyday Chemistry (provisional)
Abromination
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[*] posted on 15-2-2020 at 10:38


My experiment using boric acid and calcium chloride to catalyze a Fischer esterification is going quite well. The reflux time has taken quite a while but the smell of ethyl salicylate is strong now and a spot test leaves an oily, wintergreen residue. Boric acid alone was unsuccessful, but with the dry calcium chloride things are looking good.

35B8F2F3-0BD2-4039-8FBA-12B9F69FD7CC.jpeg - 1.9MB4A993205-80C2-4E75-B9F2-896C4241113B.jpeg - 2MB




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[*] posted on 15-2-2020 at 11:49


Nice synth! I just saw the monohydrate of CaCl2 only loses water at 260 degrees and the dihydrate at 175... That is interesting... I guess that could be used for other esterfications.
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[*] posted on 15-2-2020 at 16:28


The idea was to remove water from the reaction. Reports of boric acid catalysed esterification have been floating around for a while, and after limited success with that alone I figured the drying agent would help. Seems to work!



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[*] posted on 15-2-2020 at 18:43


Today I prepared a buffer solution to pH 7.4 on a 4400 litre scale.



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[*] posted on 15-2-2020 at 19:41


Abromination:

Your esterification experiment is interesting but I believe demands a bit more rigor before you can conclude that the CaCl2/B(OH)3 combination is an esterification catalyst. The long reaction time that your post implies raises the possibility that you are simply reaching an equilibrium between salicyclic acid and its ester. Salicylic acid is a strong enough acid to catalyze its own partial esterification (pKa = 2.98) given enough time and high enough temperature. I suggest that this may be what you are seeing that is giving you your product. Also keep in mind that CaCl2 is not a good water scavenger in alcohol solvents as the alcohol can compete with water for coordination sites on calcium, especially when a large excess of alcohol is present relative to the water produced. Following the scientific method, I suggest the experiments outlined below:

1. React salicylic acid and ethanol without any catalyst under the exact same conditions you are using in the experiment you described (molar ratios, time, temperature). Compare ester production versus your current experiment as best you can.

2. Repeat (1) using boric acid as a catalyst alone. It appears that you have tried this but you need to do it as in (1) above.

3. Repeat (1) using CaCl2 alone.

I hazard a guess that experiments 1, 2,and 3 will give results the same as your current observation. These kinds of things make chemistry interesting and can lead to new discoveries.

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[*] posted on 15-2-2020 at 22:22


Does anyone have data about the temperature where CaCl2 starts to lose methanol / ethanol as adduct? This is interesting because if that temperature is below the boiling point of the given alcohol, it should work as a catalyst with a minimal amount of acid.
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[*] posted on 16-2-2020 at 00:19


Quote: Originally posted by AvBaeyer  
Abromination:

Your esterification experiment is interesting but I believe demands a bit more rigor before you can conclude that the CaCl2/B(OH)3 combination is an esterification catalyst. The long reaction time that your post implies raises the possibility that you are simply reaching an equilibrium between salicyclic acid and its ester. Salicylic acid is a strong enough acid to catalyze its own partial esterification (pKa = 2.98) given enough time and high enough temperature. I suggest that this may be what you are seeing that is giving you your product. Also keep in mind that CaCl2 is not a good water scavenger in alcohol solvents as the alcohol can compete with water for coordination sites on calcium, especially when a large excess of alcohol is present relative to the water produced. Following the scientific method, I suggest the experiments outlined below:

1. React salicylic acid and ethanol without any catalyst under the exact same conditions you are using in the experiment you described (molar ratios, time, temperature). Compare ester production versus your current experiment as best you can.

2. Repeat (1) using boric acid as a catalyst alone. It appears that you have tried this but you need to do it as in (1) above.

3. Repeat (1) using CaCl2 alone.

I hazard a guess that experiments 1, 2,and 3 will give results the same as your current observation. These kinds of things make chemistry interesting and can lead to new discoveries.

AvB

Just now saw your post, so sorry for posting twice.
Boric acid was tried before under the same conditions. This resulted in the smell but negative results with a spot test.

My next step was already to try calcium chloride alone as well as just the acid and alcohol for a control. I am aware of the slight reaction between the two.

The point of the experiment is to try different reagents as a catalyst, so failure does not upset me. When each method has been attempted, I will repeat each of them again in the same conditions to stay true to the scientific method and isolate as many variables as possible. For right now its more of a matter of feasibility and curiosity.




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[*] posted on 16-2-2020 at 20:30


Aren't you worried about the boric acid forming esters with the alcohol?

This would both tie up your acid and add water to the reaction mixture.

You might deal with the water generated by using metaboric acid or boron trioxide mixed in with your boric acid ( edit:Or the CaCl you did use), but it might make the ethyl borate formation worse.

That's the thing: You want to soak up the water to drive the reaction but boric acid + alcohol makes water, and drying up that water also drives that equilibrium to the right as well as the equilibrium of the reaction you want.

If I'm totally wrong about boric acid forming esters in such conditions just ignore this.

[Edited on 17-2-2020 by SWIM]




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[*] posted on 16-2-2020 at 20:59


Quote: Originally posted by SWIM  
Aren't you worried about the boric acid forming esters with the alcohol?

This would both tie up your acid and add water to the reaction mixture.

You might deal with the water generated by using metaboric acid or boron trioxide mixed in with your boric acid ( edit:Or the CaCl you did use), but it might make the ethyl borate formation worse.

That's the thing: You want to soak up the water to drive the reaction but boric acid + alcohol makes water, and drying up that water also drives that equilibrium to the right as well as the equilibrium of the reaction you want.

If I'm totally wrong about boric acid forming esters in such conditions just ignore this.

[Edited on 17-2-2020 by SWIM]

I’m sure its possible in some quantity, but I beleive such a reaction requires a decent amount of a very strong dehydrating agent such as sulfuric acid or phosphorus pentoxide to occur. It is a factor that could also lower yield. The idea with the calcium chloride was to remove water from the reaction without taking place in the reaction itself to drive it right, but I am questioning its effectiveness in an alcoholic environment.

As stated before, this is not to find a better way to do an esterification, it is merely a project for my own curiosity and for fun. Its not particularly costly or time consuming, either.




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[*] posted on 25-2-2020 at 18:45


As I was wrapping up the ammonia project, I ran across something in Merck that I hadn't heard before... apparently concentrated ammonia dissolves copper metal? I scrubbed a bit of copper wire until it was shiny and put it in a corked test tube with a couple mL of concentrated ammonia. Sure enough, over a couple of days it turned the dark blue of an ammonia-copper complex!



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[*] posted on 25-2-2020 at 23:34


Concentrated ammonia only dissolves copper if also oxygen is present. The oxygen oxidizes the copper and the ammonia makes it possible for the copper to dissolve as the [Cu(NH3)4](2+) complex. Your solution turns strongly basic, the oxide ends up as hydroxide ion.

This effect actually is quite common. If a suitable coordinating agent is present in solution, then many metals can be oxidized much more easily than without the coordinating agent. Best example is the solubility of gold in a solution of KCN through which air is passed. A very nice example in the home lab is the dissolving of copper in dilute acid (e.g. 10% HCl) to which some thiourea is added. The copper then dissolves, with formation of hydrogen! Without the thiourea, copper does not dissolve in non-oxidizing dilute acid.




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[*] posted on 26-2-2020 at 03:50


That's really interesting! I have some silver cleaner containing thiourea, so I definitely try it.
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[*] posted on 26-2-2020 at 05:14


So that means thiourea reverse the redox potentials for the rxn?
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[*] posted on 26-2-2020 at 05:31


The thiourea does not reverse any potentials, it shifts redox potentials. In general, nearly all coordinating agents shift redox potentials, but some compounds have a strong effect, while others have less effect.

In the example of thiourea and copper, the redox potential of the reaction from metallic copper to copper(I) is shifted, such that it goes below the reaction from H2 to H(+) ion. Copper is somewhat noble, but not much so, and it is shifted to the range, where metals can dissolve in non-oxidizing acids with production of H2. The copper itself in this reaction then forms a thiourea-complex of copper(I).




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[*] posted on 28-2-2020 at 09:38


I know we've moved past the boric acid catalyzed esterification, but i do have experience with this.

DOI: 10.1021/ol036123g

I have used this method to prepare the alpha hydroxy ester of threonine. Good yields, and largely chemoselective between acids.


[Edited on 28-2-2020 by Medyc]
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[*] posted on 9-4-2020 at 21:40


Stores are still largely empty of cleaning supplies here! I had a bucket of calcium hypochlorite granules in the shed so I set about converting it into chlorine bleach with sodium carbonate. Filtering the calcium carbonate was the hardest part - it is very fine and the concentrated bleach quickly destroys paper and cotton! :o Next time I might try to set up a large scale vacuum filtration system.

I measured concentration by density and wound up with 5.5 gallons of ~6% bleach. I think I could do it more efficiently next time. This was way more than I needed so I donated some to the local makerspace (which is staying open making face shields) and the rest to a mutual aid group for distribution :cool:




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[*] posted on 10-4-2020 at 12:28


Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.

The silver solution has a slight blue cast from the 10% Cu, and the nickel solution is a lovely green.




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[*] posted on 11-4-2020 at 04:06


Quote: Originally posted by Medyc  
I know we've moved past the boric acid catalyzed esterification, but i do have experience with this.

DOI: 10.1021/ol036123g

I have used this method to prepare the alpha hydroxy ester of threonine. Good yields, and largely chemoselective between acids.


Interesting paper thanks. It mentions a 1971 paper claiming it was the first paper suggesting the use of boric acid for esterification but only for phenol esters. The paper is attached below

Attachment: boric-acid-lawrence1971.pdf (69kB)
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[*] posted on 11-4-2020 at 07:55


Quote: Originally posted by arkoma  
Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.
Such a waste arkoma, shouldve save a few more for bulk rxn to save reagents
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[*] posted on 11-4-2020 at 11:25


Quote: Originally posted by fusso  
Quote: Originally posted by arkoma  
Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.
Such a waste arkoma, shouldve save a few more for bulk rxn to save reagents


Nah. I've spent $100 making something I could buy for $3 Only other silver i have is a Kookaburra from Oz, and AIN'T using it.

*edit* the HNO3 was home sourced also.

*edit* Some diethyl ether



[Edited on 4-11-2020 by arkoma]

2020-04-11-155157.jpg - 91kB

[Edited on 4-11-2020 by arkoma]




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[*] posted on 9-5-2020 at 11:24


I've been experimenting with using the paper straws that are popular these days as casings for DIY fireworks. I've had a few successes, but I think the best part so far has been lighting Zn/S at night and shining the shortwave UV Tdep gave me through the plume of phosphorescent smoke. It's like incense smoke in a sunbeam, except it's ghostly green and everything is dark. It doesn't show up well on my camera, unfortunately



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[*] posted on 11-5-2020 at 12:18


Ran some sulfur dioxide through concentrated hydrogen peroxide to restock my sulfuric acid (so inconvenient, really wish you could buy it here), was a great test for my new fumehood. I had never noticed how endothermic the reaction of HCl and sodium metabisulfite is, I had always ended up having to do it in freezing weather outside before now.

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List of materials made by ScienceMadness.org users:
https://docs.google.com/spreadsheets/d/1nmJ8uq-h4IkXPxD5svnT...
--------------------------------
Elements Collected: H, Li, B, C, N, O, Mg, Al, Si, P, S, Fe, Ni, Cu, Zn, Ag, I, Au, Pb, Bi, Am
Last Acquired: B
Next: Na
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[*] posted on 30-5-2020 at 21:10


Distilled some otc isopropyl alcohol. Seemed to be already high purity. Head temp held rock steady and I discarded only a couple of mL of heads and tails.

IPA is not quite the ubiquitous substance here that it is in the US. Nice to find such good quality.
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[*] posted on 1-6-2020 at 15:09


I'm trying my first medium-scale brew & still.... 3 gallons of what the yeast pack promises is 20% alcohol are in the fridge and I'm setting up to try distilling in the backyard tomorrow :)



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[*] posted on 1-6-2020 at 19:02


I find it quite fulfilling distilling small (gallon size batch of 55% EtOH) amounts of alcohol. And besides, I'm chipping in on the Sars-Co2 abatement. Hand sanitizer is mostly EtOH.



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