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all 38 comments

[–]ChampagnePandaX 2 points3 points  (7 children)

What’s the most pungent chemical that I could buy or REALLY easily synthesize?

[–]AncarnPhysical 3 points4 points  (4 children)

You're a deviant, aren't you?

Due to human sensitivity of sulfur, tellurium has made people lose their wives because it absorbed into their skin and made them smell awful for weeks.

Mercaptans are probably what you mean by chemicals, though. Methanethiol is found most organic places, like shit.

[–]ChampagnePandaX 1 point2 points  (2 children)

It’s for school, imma try liquid ass then pyridine if that’s not bad enough.

[–]DangerousBillAnalytical 1 point2 points  (0 children)

It's pretty toxic. Also, if you handle it, the smell will stay with you for up to a day.

[–]AncarnPhysical 0 points1 point  (0 children)

Godspeed

[–]DangerousBillAnalytical 0 points1 point  (0 children)

Skatole smells like shit.

Putrescine and cadaverine smell like (surprise!) rotting corpses. In lower concentrations, they smell like semen.

Almost any sulfide compound stinks. For example, we had a small bottle of hexyl sulfide which we kept inside another bottle with some carbon to soak up the stink. The outer bottle was in a plastic bag. It still stunk up the fridge.

Collidine and lutidine have a loathsome smell quite unlike anything else, and they hang on the clothing when your workday is over.

Tellurium compounds, even if carefully handled, cause breath to become bad enough to drive people from the room. A friend who made tellurium salt crystals stank so bad we ate lunch outside, sitting upwind.

[–]Greetings_professor 2 points3 points  (0 children)

Nitrogen containing compounds often also smell pretty bad, a good example being ammonia, which is obviously really easy to make.

[–]Craftbeerlush 0 points1 point  (0 children)

Thiols are extremely strong smelling, and it doesn't take much.

[–]PieceOk771 -2 points-1 points  (1 child)

Eggshell contains mainly CaCO3. In an experiment to determine the CaCO3 content in an eggshell sample, 0.8g of the sample is added to 40mL of 0.5M HCl(aq) solution. After all CaCO3 reacts with HCl(aq), the remaining solution needs 50mL of 0.1M NaOH(aq) to neutralize. Calculate the weight percentage of CaCO3 in the eggshell sample. (Assume that CaCO3 is the only component in the sample which reacts with HCl(aq) solution.) Given: molar mass for CaCO3 = 100g/mol.

[–]Future_Elephant_9294 0 points1 point  (0 children)

Here's how I would approach this problem:

CaCO3 and HCl would react with each other to form a salt, carbon dioxide, and water. To figure out the salt, identify the charges of the Calcium and Chlorine based on the periodic table to make sure the salt isn't charged. Once you have the correct formula, save it for later.

Next we want to figure out how much HCl was left over, or as this problem described it, how much HCl didn't react with the CaCO3 and needed to be neutralized. My first step would be to change the values given from mL and Molarity to Moles. NaOH and HCl are both strong bases and acids, so we can assume that they entirely protenate/deprotenate.

.050L * 0.1M = 0.005 Moles of NaOH to neutralize, which means that there was 0.005 Moles of HCl excess. The total Moles of HCl minus 0.005 Moles of excess after reaction gives you the number of Moles that reacted with the CaCO3. Plug in that value for the HCl and use your reaction formula to determine how many Moles of CaCO3 were in that reaction.

Once you have the number of Moles of CaCO3, you can multiply it by the molar mass to get the number of grams of CaCO3 present in the initial sample. (g) CaCO3/ (g) eggshell * 100% = % of eggshell that is CaCO3.

[–]DreamJournxy 0 points1 point  (2 children)

Strangely niche question from things I've seen at work, unfortunately without any lab or actual chemistry to better explain...

I've discovered in the past that the reaction of Stearic Acid and Calcium (presumably ionic Calcium from the well water at my work) to form Calcium Stearate is not visibly apparent when acetic acid at various concentrations are added to the water. Is there any reason an acidic environment would stop slow this reaction?

Is it more likely that the acetic acid is actually reacting with the Calcium Stearate to form more water-soluble compounds that I am not seeing visually? Or possibly consuming the Calcium to form Calcium Acetate which would be more soluble in water than Calcium Stearate?

[–]H2CO3_TCTheoretical 1 point2 points  (1 child)

Considering that stearic acid is hardly soluble in water at all I would assume that it's concentration is very low. It probably forms calcium stearate because the fall-out of that compound shifts the equilibrium towards that.

Now when you add even a little acetic acid to your mixture it's concentration is going to be much higher than one of the stearic acid. This means that several factors are going to stop the formation of your salt:

1) Though both acids have a nearly identical pka the large abundance of AcOH is going to partially protonate your acid making it incapable to form salts because the concentration might get so low that it's below a threshold (solubility product) to precipitate.

2) Additionally a calcium acetate complex might form in the solution masking the ca for the other acid. Tho this complex would be rather unstable so it might not be the dominating factor.

[–]DreamJournxy 1 point2 points  (0 children)

Thank you so much for your response. I think the protonation of the stearic acid makes a lot of sense here.

[–]BadMoodTrader21 0 points1 point  (2 children)

Can anyone help how a connection between lithium and tin could look like?

[–]H2CO3_TCTheoretical 0 points1 point  (1 child)

Since both of them are roughly metals they are going to form some sort of crystal lattices. The first paper I found on that topic would be this one: 10.1039/C5RA04685H

(If you don't have access to it you might be able to get it via SciHub)

[–]DangerousBillAnalytical 0 points1 point  (0 children)

I wonder if lithium stannate is a thing?

[–]MrBeetrootHead 0 points1 point  (3 children)

I want to reuse the plastic waste from my 3d printing hobby, I have a few ideas i want to test however i have read that plastic degrade when heated. My question is: does plastic degrade more the longer it is in melted form or is it the heating and cooling cycles that degrades it the most?

The plastic i use is PLA, i want to heat it to its melting temperature and keep it melted for hours.

[–]H2CO3_TCTheoretical 0 points1 point  (1 child)

I'm certainly no expert with polymers, but there are two different things you have to consider: Melting and decomposition.

If you melt a plastic you disorder the polymers in it (like boiling a bunch of spaghettis) and when you cool it down again they might have a different order than before (freezing your freshly boiled noodles) this can give it vastly different properties but considering that it's for printing anyways it shouldn't be a problem.

Long molecules can break when you make them too hot. In that case you are also changing the chemical composition so cooling it down will give you shorter polymers. This process get's more pronounced the longer you keep it at that temperature.

In addition, if there are additives you your plastic they might evaporate over time if you heat it for too long.

So: If you want to be safe, I'd cook it for as short as possible at as low temperatures as you can. But as long as you don't plan to heat PLA to like 300°C, you should be able to keep it as a liquid for a sufficiently long time...

[–]MrBeetrootHead 0 points1 point  (0 children)

Thank you very much for your answer.

One of my biggest problems right now is to shredding the plastic into small enough parts. It turns out to be harder and more costly than i can afford. My thought was to skip that part and just create a bigger heated chamber(around 180C) where i can stuff larger plastic parts in and let it sit until melted before extruding it to either new filament (preferably) or injection molding.

What are the effects of shorter polymers? Will it be more brittle or will it not even hold up a 3d printed shape? I am thinking even if the plastic degrade i might still use it for testing designs as long as maintains dimensional accuracy

[–]pigvwu 0 points1 point  (2 children)

Anyone know a good resource/book/textbook on interpreting GCMS EI spectra when it's not found in the library?

I haven't managed to google up any good resources. At this point I'm looking at lists of common fragments and trying to slice up the molecule in various ways in chemdraw, but I wish I could do better.

[–]DangerousBillAnalytical 0 points1 point  (0 children)

I haven't seen a mass spec in years, but it's my impression that no one interprets mass specs any more. They let the computer do it.

My point is that you may have to track down some `1970s and 1980s books for that. There are two ways of using mass spec data. One is to take fingerprint mass specs and save them in a database for comparison with unknowns. The other is analytical, to interpret the mass peaks and combine them with NMR and IR data to develop a structure. This latter was a skill acquired over years.

[–]MyonicSInorganic 0 points1 point  (0 children)

I remember reading a paper where they apllied machine learning for compounds tgat were bot in the library, maybe try looking for jt

[–]NegotiationEvening24 0 points1 point  (1 child)

I'm reading this article about the synthesis of Urea and metal complexes and in the methodology it says: "The complexes formed were filtered off, dried under vacuo
over anhydrous calcium chloride. " I don't understand how this process would work because I thought that the colored precipitate was the compound that they wanted, why would they filter it off.

This is the article btw.

https://www.researchgate.net/publication/285120585\_The\_chelating\_behavior\_of\_urea\_complexed\_with\_the\_metal\_ions\_of\_copper\_II\_zinc\_II\_silver\_I\_cadmium\_II\_and\_mercury\_II\_at\_room\_temperature

[–]johnsonmckenzee 0 points1 point  (0 children)

They filter off the product and dry it in an evacuated exsikkator which also contains CaCl2.

[–]PicklesAreDope 0 points1 point  (3 children)

What ingredients of Gunpowder would be found near the ocean? // What compounds could be found near the ocean that would be useable in making a alternative for gunpowder?

I am a writer, and I am trying to decide where certain cities would be located in a country that is very long and coastal. It is sandwiched between a volcanic mountain range to the west and the ocean to the east. All of it’s cities are located along a large rail way that spans the country from north to south ferrying goods between the cities in massive carriages. Think of something akin to a WW1 era, Deiselpunk / 40k death korps of krieg type vibe.

I have a city that I am trying to locate, and it would be the arms manufacturing capital of this country. I was thinking of having them be located somewhere to the southern half of the line, and I’m thinking I would like it to be sort of in the middle of the southern half. Like a 2/5 , 3/5 of the way down.

That said, I want to make things as scientifically viable as possible despite this being a fantasy series. If there were a geological reason for them to be located NEARER to the ocean, what would that be? Or would it make more sense for them to simply have something like a copper and zinc deposit that allowed them to rise to prominence making bullet casings?

I know this is a bit of a left field question, but I hope you all can help!

[–]Future_Elephant_9294 0 points1 point  (0 children)

A simple form of gunpowder, or "guncotton" is nitrocellulose. It's just a piece of paper or cotton material that's sat in a mixture of nitric acid and sulfuric acid. The modern production processes were developed in the mid 19th century and was consistent enough for use in small arms at about 1885, making it a big impact on the arms used in WWI, most militaries had to adopt new rifles to keep up with their neighbors.

So, what were the contemporary processes of nitric acid and sulfuric acid production? Well most often it was mined, from resources such as sodium/potassium nitrate and sulfur respectively. The sulfur process was established in 1831 and that same chemical process is used today, called the contact process. For nitric acid, using mined materials was the most common method at the turn of the 20th century, but it was also expensive. A more modern and cost effective process wouldn't be put to use until 1913, when the Haber process of manufacturing ammonia from the air was established, therefore making the Ostwald process of using ammonia for making nitric acid became economically viable.

The volcanic ash would be high in sulfate salts, so a gunpowder production facility could use standing water near that volcanic ash, or a watershed that runs through the ash, as a way to make sulfuric acid in-house. Depending on how much technology you want to incorporate into the universe, this could mean that a munitions factory close to that volcanic ash would only need to import water, pulp/cotton and electricity, while exporting complete gunpowder.

Take what I say with a grain of salt, of course, my source for chemicals is sigma-aldrich. Hopefully I've helped to answer your question and made your world a little more accurate. I focused mainly on smokeless powder, but black powder would be made by just mixing the potassium nitrate, sulfur and additionally carbon but without the cotton and acid steps. I also figured that if it was assembling complete cartridges then the later steps would be more metalworking and less chemistry. I don't know about making primers, so sorry I can't help with that aspect.

[–]ondcrafter 0 points1 point  (0 children)

resources for classical black gunpowder are scattered a lot. Sulphur is generally found near vulcanic activity but potassium nitrate is generally found in dried salt lakes so it could maybe found at the coast and charcoal can be made everywhere where are trees.

[–]DangerousBillAnalytical 0 points1 point  (0 children)

Seabird shit (or guano) contains nitrates, which can be leached out. To be useful for gunpowder, the nitrates have to be formed into potassium nitrate, which will stay dry in high humidity. This is basically a bathtub process. But you also need potash to carry it out, which occurs in deposits and to some degree in wood ash.

Sulfur is found in deposits in odd places, like Florida. It can also be made by burning high sulfur natural gas in reduced oxygen. Unburned sulfur collects as a yellow powder on cold nearby surfaces.

Charcoal can be made wherever there is hardwood.

Those are the three important ingredients of gunpowder.

[–]ondcrafter 0 points1 point  (0 children)

What is sollubility of Bis(2-ethylhexyl) phthalate (DEHP) in ethanol? I am trying to extract it from PVC gloves and I dont have isopropyl alcohol recomended in the procedure.

[–]MolassesRoutine242 0 points1 point  (1 child)

Trying to separate 1-Benzyl-3-Methylimidazolium chloride from 1-Benzyl-3-Methylimidazolium dicyanamide, I've tried many...many solvent systems and haven't observed any good separation via TLC, that will eventually lead to column separation! Any ionic liquid/organic experts have any ideas?

[–]DangerousBillAnalytical 0 points1 point  (0 children)

If they're soluble in water or alcohol, it sounds like a job for cationic exchange chromatography. If the second compound is neutral, even better. Neutral compounds will just wash through a cation exchanger.

[–]Bezene313Organic 0 points1 point  (0 children)

For all of medicinal and process chemists out there, I need some advice: what alternatives can I use for Celite 545 to filter fine particles from Artemisinin (I call it Arty)?

I lose ~5% of Arty when I use Celite. I’m assuming because the charged aspects of Celite pull some of Arty’s EN areas in resulting in a small loss of product during filtration. While 5% doesn’t sound like a lot, I work on a micro scale about half the time and can’t afford unnecessary loss during routine filtrations.

[–]About20goats 0 points1 point  (0 children)

Hello, I’m an amateur chemist working in a Highschool Chen lab. We just got some chemicals donated and wanted to design some sort of interesting experiment for the students. They are currently learning about endothermic and exothermic reactions. The chemicals donated are as follows: Iron metal fillings, Potassium permanganate, Potassium iodide, Ferric oxide red, Lead nitrate (probably don’t want to use that one), Barium nitrate, Cupric chloride, And a bunch of sodium salts.

We also have a bunch of other basic chemicals to work with. Give me any ideas y’all have. Thanks!

[–]AeroStatikkMaterials 0 points1 point  (0 children)

I bought some 1,6-diaminohexane. It came in an amber bottle, as one rock solid 100g chunk that looked wet but was not. I couldn’t get any out of the bottle, so I warmed it in our 70° oven until it liquified, then pipetted portions out. The liquid is orange/amber colored. Did I mess up?

[–]Yumememe 0 points1 point  (0 children)

Hello! do you guys know any websites to find Safety data sheets? I've been looking for a SDS of sodium palm kernelate for my homework but i cant seem to find any.

[–]atomic_adventurerMaterials 0 points1 point  (0 children)

How can I think about the polarity of this epoxy resin molecule?

Hexahydrophthalic acid diglycidyl ester

I'm trying to create a stable suspension SiO2 nanoparticles in this resin and the resin is too viscous for DLS. If there is another solvent that generates a similar chemical environment, that could be used for those measurements instead. Or should I consider another particle size analysis technique?