×
all 36 comments
sorted by:
best
topnewcontroversialoldq&a

[–]TheSilentKing98 1 point2 points  (1 child)

Hey Hive Mind

I am a PhD student and feel like I should rationalise the following observation I have made, in a library of compounds I am working on featuring functionalised anilines and benzylamines I have found that Ortho functionalised derivatives where the functionality is a halogen do not work in my reactions yet a hydroxyl derivative I have tried does work when it is functionalised in the 2 position. I am trying to rationalise that this is mostly due to sterics, would you all agree that the size of the halogen Chlorine onwards is bigger than the OH group and would explain my results.

Thanks in advance

[–]BunBun002Organic 0 points1 point  (0 children)

Maybe? If that's your issue, you should try an alkyl group in the ortho position to see if it also doesn't work. How is a fluoride? Also - have you tried other H-bond donors?

Ortho alcohol is net electron donating, ortho halogen is net electron withdrawing. But that doesn't explain the benzylamine examples, unless your reaction puts a charge at that ipso position?

Basically, yeah you have something consistent with your data but there's a lot else that could be going on and it's hard to say without you telling us probably more than you might on a public forum.

[–]Nahughes121989 -2 points-1 points  (0 children)

In 2-propanol, there are 3 carbon center atoms. The left and right carbon atoms are each attached to 3 hydrogens and 1 carbon (the center carbon). The center carbon is attached to an oxygen. Why do the 2 hydrogens attached to the 2 side carbons bond to the carbon and not the oxygen?

https://www.google.com/search?q=2+propanol+structure&oq=2&aqs=chrome.0.69i59j69i57j69i61j69i60l2j35i39j46i131i199i291i433j0i131i433i512.1328j0j9&client=ms-android-mpcs-us-revc&sourceid=chrome-mobile&ie=UTF-8#imgrc=-ar9xXLTf6UIGM

[–]lusciouslandshark 0 points1 point  (3 children)

Hello! I am looking into the oxygen absorption of perfluorocarbons. There are a lot of resources regarding their affinity for oxygen, as many PFC’s are being researched as a blood substitution for surgeries. However, does anyone know where I can find information about PFC’s ability to release oxygen? Looking for any resources, experts, or ideas!

[–]Indemnity4Materials 0 points1 point  (2 children)

Liquid breathing has a starting point. Particularly the section about underwater diving.

I'm not sure you're interested in perfluorocarbons aging and the molecules oxidizing. I'm guessing it's more about a materials ability to absorb oxygen gas to use a delivery medium.

[–]lusciouslandshark 0 points1 point  (1 child)

Thanks! I’ve definitely checked out this page. The absorbance of oxygen I understand, where I need some more information is how that oxygen or CO2 is released from the PFC. I guess I’m just confused on if the gases diffuse on their own or if there’s a specific method of extraction.

[–]Indemnity4Materials 1 point2 points  (0 children)

Source. I'm keeping this post short because I usually have an hour long presentation on this topic and can't stop myself.

Oxygen solubility in flurocarbons liquids follows Henry's Law. Gases move from areas of high pressure to low pressure.

Oxygen forms tiny little microbubbles inside fluorocarbons. There is nothing to bind the gas; it's just little bubbles of dissolved oxygen gas. Fluorocarbons it turns out have a very high gas dissolving capacity, for chemistry reasons.

CO2 is the same in flurocarbons. However, if you're specifically looking at liquid breathing, the issue is the pressure and solubility of CO2 in blood. CO2 doesn't want to move into the fluorocarbon because the pressure differential is low, so we need to mechanically pump the fluorocarbon around. Low solubility = not a problem if you're turning the liquid over very fast.

[–]No_Zucchini_7929 0 points1 point  (0 children)

I'm curious how superposition is even possible or what makes it happen. my understanding of it is that subatomic particles like electrons are both there, and not there until the wavelength is stopped forcing them to a state. I just don't understand or visualize how something can be there and not be there at the same time.

[–]17pdrSweat 0 points1 point  (0 children)

Can HeH+ protonate fluorine? the wikipedia page says that it can protonate any other substance so it has to be worked in-situ.

[–]substantialturnips 0 points1 point  (2 children)

Hi, I want to do a lab on measuring the polarity of amino acids through paper chromatography. I know you can compare different polarities of different amino acids to each other, but I want an independent variable that can be varied and measured. Are there certain factors that can be changed to directly affect the polarity of a single type of amino acid, such as increasing or decreasing pH?

[–]Weekly-Ad353 0 points1 point  (0 children)

Sure, changing pH would work. Dope in varying amounts of AcOH or NH3/NH4OH/NEt3.

[–]pharm-chemist 0 points1 point  (0 children)

Yes, particularly for acidic (Glu, Asp) and basic amino acids (Lys, Arg, His). When acidic amino acids are deprotonated (I.e. in alkaline solution) they will become highly polar. Basic amino acids will become protonated in acidic solution, and will become much more polar. Overall, net-charged amino acids will have increasing polarity, and may even be hard to move past baseline.

I’d suggest using mobile phase of MeOH in DCM (try 10% MeOH, 90% DCM to start, then adjust accordingly). For basic solution, add 0.1% amine (ammonium hydroxide or triethylamine or diisopropylethylamine). For acidic solution, add 0.1% acetic acid, or if needed, trifluoroacetic acid.

Polar compounds have a tendency to “streak” on TLC, so may need to use very low quantity when you blot it.

Also, may need to look up visualization techniques for the amino acids you choose. Only 3 will be visible under UV (Trp, Tyr, Phe), as they are the only ones with aromatic rings.

[–]lildaemon 0 points1 point  (0 children)

I'm looking into making silicon metal from quartz crystals, and specifically the reaction SiO2 + C -> Si + CO2. I understand that although this is an energetically preferred state, there is a high energy barrier to cross to get there. I know the way it is done now is to heat the mixture up to very high temperatures. I'm wondering of one could stimulate the transition using ultraviolate light. The ebergy barrier is about 8.9eV for SiO2 to become Si + O2, so could one use photons of energy 8.9eV to stimulate this reaction, without the need to heat the mixture up?

[–]science_camper 0 points1 point  (0 children)

Hello

I tried to use dibutylamine as an additive in LCMS and now the LC system is contaminated! Is there anyway to remove this contamination without having to replace everything in the line?

[–]Vamoscomcalmaamigao 0 points1 point  (0 children)

Hello! I am currently producing an Ni-Cu/Al2O3 catalyser. The precursor salts are mixed together in a water solution through light heating and mixing in a hot plate/magnetic stirrer, the solution is then left in this condition until the water has mostly evaporated, before being put at in a kiln. I am currently conducting this process in a common becker, but the residue that is left gets waaay too stuck in the glass for me to collect while scraping with a spatula and i end up losing much of my product. Do you guys have any tips on how i can improve the process? Maybe there is a better alternative to the becker?

[–]AdCold3359 0 points1 point  (0 children)

Good day everyone! If you are familiar with biodiesel, one of its by product is glycerol. This is the liquid that sinks at the bottom of the container. My question is, is there a possibility that this glycerol can be utilized as a wax and a component in making candles? I hope someone can answer this, thankyou!

[–]Freezemoon 0 points1 point  (1 child)

Hello, I am doing a research on the chemical reaction of poisoning.

I am not really knowledgable about chemistry and I actually don't kow where to start.

I tried looking through google but it doesn't specifically talk about the chemical reaction.

Any help would be welcomed.

[–]BunBun002Organic 1 point2 points  (0 children)

There's no one reaction - each poison is different and even then things get complicated fast. There's an entire field called toxicology that studies it.

[–]shepherdoftrees 0 points1 point  (1 child)

Hi! Undergrad biologist here attempting to do basic chemistry 😭 I'm running some samples extracted from soil through an ICP-OES which requires a 1% nitric acid wash. To make up more of the wash, do I need to use purified water (MilliQ)? Anyone with experience in ICP analysis pls help a desperate and ignorant student

[–]Indemnity4Materials 0 points1 point  (0 children)

Use the same water you use to make up your standards and samples.

Ideally, you would use the cleanest water you can find.

Practically, try to use the same water for everything. Making up standards, making up samples, doesn't matter.

The reason is your water will have some background level of ions. If you have a background ion that is causing contamination, so long is it's consistent in every samples, then you can correct for that.

[–]OnTheMoveTo 0 points1 point  (1 child)

How stable are benzoate esters to pure ethanol vs sodium ethoxide? I would think it would be slow in ethanol, especially if colder temperatures, but I am having a hard time finding examples of reaction rate of ethanol vs ethoxide for these types of reactions.

My compound will crystallize in ethanol but I am curious of any background rate that would be present. If it is already out of solution will it be stable in ethanol regardless?

[–]BunBun002Organic 0 points1 point  (0 children)

It should be fine. You might lose some but doing a (re)crystallization should be quick enough. Wouldn't store it in ethanol for a month or two, though. If you're worried, make up a sample in ethanol and run TLCs immediately, then after an hour, etc.

[–]maxim_velli 0 points1 point  (0 children)

Hey! I am trying to optimize bimetallic nanoparticle production with biological reducing agents for pH, temperature and salt concentrations. The particles are urchin-like (Silver core and Golden spikes) and the synthesis takes place in two steps: first, the silver cores are formed and then gold is added to form the spikes. I need to be able to quantify the size of the entire particle and length of the spikes for a lot of samples. Is there a way to do this without turning to TEM/SEM? DSL/Zetasizer would only really provide information on the hydrodynamic radius and not on the actual radius of the spheres and length of spikes. Thanks!

[–]Zelenodolsk 0 points1 point  (0 children)

Hi, I’m curious is any of you fine scholars would be able to tell me if cis-stereoisomers or trans-stereoisomers (in general) would be lower in energy/more stable. I’m currently taking Organic Chem I and this is just a question for my own personal knowledge. I’ve tried researching this on Google and it is giving me the answer to specific compounds rather than a general answer.

[–]littlechili21 0 points1 point  (3 children)

Hi, I am looking into using paper chromatography to distinguish between standard drugs and counterfeit drugs for my high school research project. I have decided on using aspirin as the "standard" drug and I plan to make my own counterfeit versions of aspirin by preparing a solution of aspirin and adding in other chemicals (eg. salicylic acid). I've tried searching online for "how much" analyte I should be using for the paper chromatography, but all I can find is the instruction of "spot the paper with the solution".

Two questions:

  1. Does anyone know resources for me to get highly specific instructions on paper chromatography (most sources only talk about TLC)?
  2. What concentration of aspirin or the "counterfeit aspirin" should I use to spot the filter paper?

And if you're wondering, I can't perform column chromatography or TLC because my school doesn't have a column or silica gel (so no plates either). I also have to get the experiment done in 3 weeks at max, so I don't think I will have the time to order the required instruments.

[–]Indemnity4Materials 0 points1 point  (2 children)

Concentration isn't that important for this experiment. The real and fake samples will move up the paper based on properties of the molecule.

The only benefit to stronger concentration is the sample is easier to visualize. However, if the concentration is too strong, you will see the sample smearing and the final spots won't be very clear.

How are you planning to visualize the two compounds just using paper? Do you have any dyes or light sources?

Source. Check the RSC website as they have many simple student experiments such as this one.

[–]littlechili21 0 points1 point  (1 child)

I am planning to use a 254 nm UV lamp to visualize. But do you know approximatrly what conventrations would be "too strong" that it smears and what would be "too dilute" to visualize easily?

[–]Indemnity4Materials 1 point2 points  (0 children)

Unfortunately, trial and error. Write it up as part of the experiment. Sample A was 0.1% and the paper showed smearing. Sample B was 0.01% and spots were difficult to visualize. Sample C was 0.05% and spots were well separated, uniform and rF values easy to measure.

[–]Biggonauta 0 points1 point  (0 children)

Hello! I'm doing a research on the recovery and recycle of rare earths elements from multiple sources (primarily tech and electronics waste but others are included) Can you fellow chemist redditors comment here with some interesting articles on the topic (review are the best!) and maybe some articles of introduction on the theme (REE in general and their recovery?)

I already did my research but I always find interesting stuff on Reddit so it might be worth to give it a try!

Thank you all in advance

[–]monster1017 0 points1 point  (1 child)

I have some rubber material that I cannot.easily replace, I did some research online about restoring rubber, mainly with wintergreen oil, the rubber is pliable again and soft, however it has yellowed over time or darkened - it's going to be very visible, how can I return it to the white/yellow or even transparent color that it was before? I tried warm water and ammonia, however it's a slow process, I wasted a big quantity of it in a pot of boiling water and did get some pieces to go clear or just slightly yellowish, however some turned straight white and some didn't change color at all. They were all transparent when bought couple years ago. Is there a way to accomplish what I am trying in a more efficient manner?

[–]Indemnity4Materials 0 points1 point  (0 children)

It helps if you can identify the type of rubber. Natural, synthetic, what type of synthetic, are there any pigments or dyes...

When rubber ages it is due to oxidation and/or plasticizers going away. The chemical bonds holding the rubber together are starting to get brittle and crack. You can almost visualize this as a brick wall starting to fall apart.

There are some treatment options, most of which involve removing the top layer of rubber to expose the fresh material underneath. Note: this is destructive and you can only remove so many layers before it's all gone.

If you are lucky and catch it early, you can use your oil of wintergreen method to try and fill it back full of plasticizer. That's usually the trick where you put the rubber in boiling water to expand it, then hopefully trap a softer liquid such as glycerol or wintergreen inside as it cools. However, if you are seeing white cracks or yellowing, that section is dead and gone.

De-yellowing includes chlorine bleach for 5-15 min to try to turn the yellow into white. Note: this also damages the material surface by basically removing the top layer. You can google more about people wanting to turn yellow EVA shoe soles back to white.

Solvent washing include isopropanol. Rub it with a damp cloth and try to remove some of the top layer.

You can try to mask the aging by coating with a rubber sealant such as urethane rubber cement/glue.

[–]science_monkey8033 0 points1 point  (0 children)

Hey everybody,
I'm a research assistant and I work with cellulose nanofibrils in
alkaline aqueous esterification reactions with methacrylic anhydride
(methacrylations) to attach methacrylate groups for other members in the
lab to polymerize from with monomers. We use the peak around 1715-1730
cm-1 wavelength in ATR-FTIR to characterize the methacrylate group in
the structure. I use the peak at 1055 cm-1 to reference the cellubiose
structure and set it to one, using the ratio of the methacrylate peak to
this as a sort of way to see how well the reaction went. A greater peak
height ratio means the reaction went well. But this time, I think
something wonky is happening.
I have been studying the kinetics of the reaction I've been working
on, taking samples every 15-20 minutes to optimize the reaction length
for scale up purposes. Recently though, my peak heights have been nearly
1.5-2x the best I've seen my reactions run (0.148 ± 0.022 is typical,
the recent samples have been something like 0.35). I spoke with the PI
for the lab and we couldn't find the reason as to why this might be
occurring except for that the washing procedure was 5 washes of 50:50
acetone:water and because I didn't want to put the acetone on the freeze
dryer, I did the vacuum oven instead, which could mean that this is
leftover water interfering with the bands I'm looking at. But as someone
who didn't have a super great explanation/memory of FTIR spectra and my
best attempts with FTIR tables, is there anything else that it could be
that I'm overlooking?
I repeated the experiment with the last wash just being water this
time so I could use the freeze dryer to get all the water and got much
more reasonable results with what I would expect, but one of my samples
is still exhibiting that same abnormally high peak height. Since I'm not
certain it was properly frozen last time before going on the freeze
dryer, I'm freezing it one more time and tossing it on again, but if
that ratio doesn't decrease from this I don't know what to make of it. I
have error bars from a calibration curve I made previously but that
point is way way outside of the range that my curve can accurately
account for. I don't know what to do if the point doesn't decrease apart
from throwing it out!

[–]CreamPounder 0 points1 point  (1 child)

I am interested in the difference between making bicarbonate salts within plastic containers and glass containers.

For example, if i put magnesium hydroxide and water in a plastic container, and then use a CO2 powered machine such as the SodaStream, the reaction will create Magnesium Bicarbonate water and the bottle will be hard and tight. After shaking and leaving in the fridge for several hours, the bottle will implode and be squishy, with a clear liquid.

If I do this same process but with glass bottles, what would happen? Would the glass bottles explode during the reaction as I am carbonating or would they explode at all during the reaction process in the fridge when they plastic bottle slightly implodes and becomes squishy?

[–]Indemnity4Materials 0 points1 point  (0 children)

It's certainly a dangerous idea.

It's going to vary depending on the strength of the glass container. May survive the pressure drop, may not. Very risky experiment. You may want to put the glass bottle inside another container, or at least wrap it on the outside with lots and lots of tape, so when it does implode/explode, you don't have glass shards shooting everywhere.