Sciencemadness Discussion Board - The short questions thread (3)

The short questions thread (3)

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tetrahedron - 15-10-2012 at 13:42

methyl benzoate

benzyl formate

phenyl acetate

phenyl acetate would be a closer match for methyl benzoate, because it has both a phenyl and a methyl extremity, with an ester in between. inverting the direction of the carboxylic group can result in similar physiological activity: compare the opioids pethidine and MPPP:

OTOH a carbonyl extremity would probably impart a different behavior.

[Edited on 15-10-2012 by tetrahedron]

chemrox - 16-10-2012 at 13:27

This seems a little more involved than a short question

Hexavalent - 30-10-2012 at 12:29

I've just seen a redox equation involving manganese in varying oxidation states, and mid-way through the equation "Đ" is seen. In this instance, what does it represent?

watson.fawkes - 30-10-2012 at 13:17

Quote: Originally posted by Hexavalent

I've just seen a redox equation involving manganese in varying oxidation states, and mid-way through the equation "Đ" is seen. In this instance, what does it represent?

Example? It might be an abbreviation.

Boron Trioxide - 2-11-2012 at 12:50

I have a few short slightly related question, answers to any or all would be appreciated.

-In the thread regarding MMO Electrodes it is said that MnO2 on its own is " oxygen selective anode coating" and will not produce chlorates, only oxygen, is this true?
-Is chrome attacked by nitric acid?
-When producing nitrates, of specifically manganese or cobalt, will lower concentrations of nitric acid work or does it need to be 70%?

And less importantly
-Are there any other solvent aside from MEK that will dissolve PVC?

Thanks for any answers

Mailinmypocket - 2-11-2012 at 14:11

Quote: Originally posted by Boron Trioxide

And less importantly
-Are there any other solvent aside from MEK that will dissolve PVC?

Thanks for any answers

I scanned this but I don't know how much this will help you- I find it to be very useful... There is obviously a difference in attack vs. dissolution but it might be helpful.

Attachment: Chemical Resistance Chart.pdf (1.8MB)
This file has been downloaded 955 times

As for the chromium? No reaction with nitric acid. It passivates. Just for fun I did a quick little demo here

70% HNO3, along with a dumped out container of pure Cr.

Threw a bit of Cr in the HNO3.... No reaction.

And of course, to avoid wasting the HNO3, who could resist....Copper turnings!!!!(you can see the nitrogen dioxides running over my skin, so nice!)

[Edited on 3-11-2012 by Mailinmypocket]

[Edited on 3-11-2012 by Mailinmypocket]

chemrox - 16-11-2012 at 02:20

Can someone tell me how to use url shortcuts of the type: /i5jthom /4gb1zjt ?
@kmno4 ?

Thanks
CRX

Conrad - 16-11-2012 at 06:58

Hi!
I recently bought a small bottle of phenylacetaldehyde.

Now, after 2 weeks in the chemicals cabinet, the bottle has grown half full of crystals.
I know my benzaldehyde loves to do this and turns into benzoic acid quite fast when not stored properly.
So, now comes the question:
Do I have a bottle half full of phenylacetic acid?

Does PAA build up so quick?

watson.fawkes - 16-11-2012 at 07:47

Quote: Originally posted by chemrox

Can someone tell me how to use url shortcuts of the type: /i5jthom /4gb1zjt ?

Read the instructions at http://tinyurl.com/. They're not the only service in this category, FYI.

I recommend against using these if your links have any kind of archival value. By concealing the actual URL, you prevent people from using archive.org or other time machines to find the content if either the site itself or the redirector dies.

Hexavalent - 16-11-2012 at 13:15

Would it to be possible to have an aqueous solution of just one cation or anion - e.g. Na⁺ or Cl^- with no other cations or anions present?

I have no use for it nor can I imagine one, this is simply out of curiosity.

[Edited on 16-11-2012 by Hexavalent]

watson.fawkes - 16-11-2012 at 14:49

Quote: Originally posted by Hexavalent

Would it to be possible to have an aqueous solution of just one cation or anion - e.g. Na⁺ or Cl^- with no other cations or anions present?

Charge balance. Electrostatic force.

Hexavalent - 17-11-2012 at 14:09

http://courses.chem.psu.edu/chem36/Experiments/Exp84.pdf

Can EtOAc be substituted for ether during the washing??

Vargouille - 23-11-2012 at 09:17

Yes, but it's suboptimal. EtOAc is more polar, more soluble in water, and has a higher boiling point. I do not know the solubility of methyl benzoate in EtOAc, but I do know that it is miscible with diethyl ether.

EDIT: And, from this site, water is more soluble in ethyl acetate than in ether.

[Edited on 23-11-2012 by Vargouille]

Solubility of compunds and it's behaviour

maxpayne - 24-11-2012 at 12:53

I have a 2 questions, and would appreciate if someone answer:

1. If I have two compounds dissolved in water, and one of it is soluble in non-polar solvent, will this other compound 'almost' completely go to that non-polar solvent if I add it? Please explain.

2. If I have a compound A dissolved in a solvent X, and then I add another compound B which is more soluble in given solvent X, will compound A get out of the solvent X if I saturate it with compound B? (salting out??)

Vargouille - 24-11-2012 at 13:43

To answer the first question, it won't be that simple. If you have a substance that is soluble in water and in a non-polar solvent, and it is more soluble in the non-polar solvent, you will have to preform multiple washes, with vigorous mixing, mind you, to extract the majority of the substance. With each wash, a little more of the substance goes into the organic phase, and eventually, with enough washes, there will be very little of the substance left in the aqueous phase.

To answer your second, only if compound A is less soluble in a solution of compound B, or through the common ion effect. For example, isopropanol can be salted out by addition of sodium chloride, which raises the polarity of the solution so that isopropanol is less soluble in the solution. If it works by the common ion effect, compound B is a salt that shares a cation or an anion with compound A. Then, the addition of whichever is shared will force the equilibrium of the dissolution of compound A to the left. For example, assume that A is NaCl and B is HCl. The addition of HCl increases the concentration of Cl^-, shifting the following equilibrium to the left, according to Le Chatelier's Principle:

NaCl(s) <=> Na⁺(aq) + Cl^-(aq)

Thus, NaCl will precipitate if it is concentrated enough. Read up on the solubility product (Ksp) for the math behind it.

[Edited on 24-11-2012 by Vargouille]

Hexavalent - 26-11-2012 at 09:30

I recently ordered some copper sulfate from the internet, and when it arrived, I noticed that it had many small, off-white particles in it. I recrystallized it successfully (including a hot filtration, as there was some insoluble junk at the bottom of the vessel), and decided to make some tetraamine copper (II) sulfate with what remained in solution.

So, I added concentrated ammonia until no more copper hydroxide formed, and then continued adding until it all complexed back into solution. I then added ethanol until no more precipitated out, and then filtered out the lovely blue crystals.

What have I made, though? Tetraamine copper (II) sulfate, or tetraamine copper (II) hydroxide? Or a mixture of the two, perhaps?

Any help would be greatly appreciated.

Vargouille - 26-11-2012 at 11:20

Tetraaminecopper(II) sulfate. Tetraaminecopper(II) hydroxide (Schweizer's Reagent) decomposes in alcohol. I made Schweizer's Reagent once before, by creating Cu(OH)2 and dissolving in NH3. If I recall correctly, this looked somewhat different from a mere tetraaminecopper(II) solution. In any case, this is a quote from my thread on the topic:

Quote:

The supernatant is drawn up in a pipette, and added to an excess of Crown Denatured Alcohol (a mixture of 65-75% MeOH, 20-30% EtOH, <10% IPA, and <10% methyl isobutyl ketone), to precipitate the Schweizer's reagent in the same way as an amount of tetraaminecopper (II) sulfate was isolated previously. This resulted in a white precipitate with a slight blue tinge. More of the supernatant is added, and the precipitate begins to go yellow, eventually becoming a tan color. An exotherm may have occured, but it is unclear due to high ambient temperature. Some white fumes were noted to be evolved. The solution is filtered, and the tan precipitate allowed to dry somewhat. It quickly becomes darker brown, and when added to an amount of water, it dissolved very slightly, creating a yellow solution.

I don't recall ever producing blue crystals, though. The sample of tetraaminecopper(II) sulfate monohydrate I have is in the form of a purple powder. The more gradual addition you attempted may have something to do with that.

smaerd - 27-11-2012 at 14:48

Speaking of ligands I have a question.

Say one has a zwitterionic ligand such as an amino carboxylate. Say it is chelating some arbitrary metal ion, silver for example. Is it possible to protonate the (primary)amine of the ligand with a strong acid? If so would this turn the bidentate ligand to a monodentate ligand? I don't know much about inorganic chemistry or metalorganic complexes(just started learning about them this week), so pardon my ignorance.

DJF90 - 27-11-2012 at 15:03

@Smaerd: This question reminds me of the template approach to macrocyclic ligand synthesis. Great, so you assemble your multi-dentate ligand around your nice template ion. So how do you free your ligand in order to use it on stuff you want to? One common trick is protonation (of amine based ligands). The amino groups can then no-longer bind the central metal ion. Another trick is to adjust oxidation state. This can cause change in coordination geometry and make the macrocycle a poorer ligand for the metal ion (in different oxidation state). I'm sure you can find examples if you look into maybe templated cryptand synthesis or (I know for certain) there are examples from catenane synthesis (Cu or Au template, IIRC). Rotaxanes are similarly interesting.

I'm not really sure how (or if) this answers your question, but I thought it kinda relevant. If the carboxylate complex is desired, I'd guess you'd have to be pretty careful with your acid addition (titration...). I also wonder if the high density of positive charge (from protonating the amino groups) around a similarly positively charged cation would be a strong enough driving force for the complex to fall apart.

[Edited on 27-11-2012 by DJF90]

smaerd - 27-11-2012 at 16:39

This is really fascinating and might explain some of my experimental results(playing around) from months back

! Speaking of pH I noticed an insoluble aminocarboxylate bidentate copper ligand upon careful acidification(H2SO4) became soluble(with no characteristic blue copper colored solution as it was prior to ligand formation). Chelation of an amino acid was initiated by dissolving Copper Sulfate in an aqueas solution. On basification/neutralization it precipitated again. When it was being dissolved into the acid solution it made these interesting cracking noises, I had never really heard anything quiet like it in a beaker. In this case, though, I doubt the ligand fell apart completely. After electrolysis with aluminum cathode and anode a brown precipitate was noted. No further analysis has been done, yet, maybe I can get one of my professors interested enough to let me have some time with a few analytical instruments.

Rotaxanes are very interesting, spent many day dreams about 'on and off' 'nano-switches' using rotaxanes via electrolysis

. Figured if I could conjure one up maybe it could be bonded or deposited onto a solid phase or something for logic gates. Again, Day-dreams.

Thanks a lot for the information. It sort of danced around and hit all of my points of interest and left me for some new things to research.

[Edited on 28-11-2012 by smaerd]

DJF90 - 2-12-2012 at 16:33

Cool, really glad I could help answer your questions and inspire you simultaneously.

Quick question of my own; Its well known that NaBr/NaBrO3/H2SO4 (aq) will yield (chlorine free) bromine, but does anyone have any experience with this system? My only experience with it was in very dilute solution for a physical chemistry experiment in kinetics; the rate equation for the generation of bromine by this reaction is actually quite complicated, IIRC, if anyone is interested I can try and find it.

AndersHoveland - 3-12-2012 at 00:24

Ethanol can be dehydrated using sulfuric acid to form diethyl ether. This can be further dehydrated into ethylene.
Can dimethyl ether be further dehydrated by concentrated sulfuric acid also? What forms?

sparkgap - 3-12-2012 at 01:06

I don't have a reference offhand, but if memory serves, dimethyl ether will only get protonated to (CH₃)₂OH⁺...

sparky (~_~)

Hexavalent - 3-12-2012 at 14:36

Vanillin can be reduced to vanillyl alcohol with sodium borohydride (Microscale Reactions of Vanillin, Rosemary Fowler, Cottey College, Nevada (see a link in the references section)), however I lack the latter reagent. Can anyone suggest any other reducing agents? NH₄Cl/Zn in acetone, perhaps?

Any help would be greatly appreciated.

sparkgap - 3-12-2012 at 15:26

Hexavalent, have you seen this?

sparky (~_~)

DJF90 - 3-12-2012 at 15:37

I will send you the requisite amount, remind me when we speak tomorrow.

Boron Trioxide - 5-12-2012 at 19:30

Another quick question

What purity is Animal Feed Grade Material, and what might its impurities be?

Thanks

Vargouille - 6-12-2012 at 03:30

From what I can tell from a cursory search, from a purity standpoint, it's somewhere between technical and reagent grade. As for impurities, it depends largely on the substance in particular and the industrial procedures used. Generally, the only impurities are those that either pose no or little harm to animals, or lack the concentration to be dangerous.

Eddygp - 21-12-2012 at 08:54

Is gluconic acid very soluble in ethanol, even cold ethanol?

Vargouille - 21-12-2012 at 12:01

http://pubs.acs.org/doi/abs/10.1021/ie50518a030

Quote:

It is extremely soluble in water, but only slightly soluble in [ethyl] alcohol, and it is insoluble in most other organic solvents.

gnitseretni - 22-12-2012 at 12:24

What do you do when the stirbars keep flying into the walls of the container? The smaller (1"-1.5") stirbars don't but they don't stir well enough. The 2" stirbar I have would stir well enough if it didn't keep flying into the walls of the container. I'm trying to stir about 400ml of mixed acids. I wanted to nitrate some methanol but it being so much lighter than the mixed acids it needs to be stirred pretty well otherwise it just floats on top. I poured the mixed acids in a wider container and that seemed to help a bit, but still not enough.
Are oval stir bars better? What works best for mixed acids?

smaerd - 22-12-2012 at 12:41

Generally stirring is a non-issue for liquid solutions of low viscosity. If it's hitting the walls of the container in my experience the stirrer is going too fast, or the container is convex. Usually I solve this by slowly raising the speed of the stirrer, hopefully this helps and aren't stupid answers or things you've already tried.

gnitseretni - 22-12-2012 at 15:05

My stirrer goes from 1 to 10 and the stir bar will hit the walls if I go past 7 regardless of how slow I raise the speed.
Anyways, I decided to do the nitration anyway. I used the 2" stir bar and a wider container. I had it stir as fast as it would go without the stir bar flying against the walls. The vortex got smaller and smaller as I added more methanol, but it still must have stirred well enough because I got methyl nitrate. Didn't think I would as it looked like the methanol stayed on top the whole time, but then again it was hard to tell as all liquids were clear as water. Oh well, guess I worried too much

Simbani - 24-12-2012 at 08:39

I have a quick question:
I wonder about the chemical stability of cyanoacrylates (glue). I really like to use this stuff to make my detonators gas-tight, but I don´t know if there could be any
side reactions with HE´s of the following familys: nitramines, nitric esters, (peroxides), azides(lead and silver) and maybe stability towards reduction (zinc, magnesium and so on).
And what are the products od decomposition?NOx?

[Edited on 24-12-2012 by Simbani]

elementcollector1 - 25-12-2012 at 19:37

No idea on the last one, but I think it wouldn't be too resistant (considering it's basically Superglue).
My question: is potash pronounced pot-ash or poe-tash?

Vargouille - 25-12-2012 at 20:17

Merriam-Webster says "POT-ash" and so do the Canucks, and I trust the Canucks.

smaerd - 29-12-2012 at 13:29

In "The Art of Writing Reasonable Organic Reaction Mechanisms" by Robert Grossman

the first problem asks which resonance structure is best. This one kind of confused me because I know boron to be an exception to the octet rule and be "satisfied" with 3 bonds or 6 valence electrons. Ex: boron triflouride, borane, etc.

I said the structure on the left was the best using this knowledge because there is no charges making it 'preferred' but the author says the one on the right is preferred because all of the atoms have a full 'octet'. Not sure who is actually right here but as you'll see even my chemical drawing utility(Marvin Sketch) doesn't force a full 'octet' on boron.

Hexavalent - 30-12-2012 at 13:17

Question: does anyone know generally how 'clean' "One Shot" drain opener/cleaner is, in terms of chemical purity? Does it usually have dyes in it, or lots of suspended carbon?

Many thanks.

@Vargouille: I've always pronounced it as 'POT-ash', and IIRC several pseudo-scientific television shows have always documented it as 'POT-ash'.

Nicodem - 31-12-2012 at 05:09

Quote: Originally posted by smaerd

In "The Art of Writing Reasonable Organic Reaction Mechanisms" by Robert Grossman

the first problem asks which resonance structure is best. This one kind of confused me because I know boron to be an exception to the octet rule and be "satisfied" with 3 bonds or 6 valence electrons. Ex: boron triflouride, borane, etc.

I said the structure on the left was the best using this knowledge because there is no charges making it 'preferred' but the author says the one on the right is preferred because all of the atoms have a full 'octet'. Not sure who is actually right here but as you'll see even my chemical drawing utility(Marvin Sketch) doesn't force a full 'octet' on boron.

The title of the book, "writing reasonable...", is not there for nothing. If you have a B-N bond where boron has a vacant valence orbital and the nitrogen a lone electron pair, it becomes unreasonable to expect no double-bonding interaction. It would be like writing the double bond of the alkenes in the form of C⁺-C^- rather than C=C. While such formality is advised when depicting ylides, nitro groups, sulfoxides, phosphine oxides and the like (in order not to break the Lewis bond conventions), it is unreasonable to do elsewhere. Nevertheless, where writing the more realistic type of interactions would make a mess, like in the case of non-localized B=N bonds, it is preferably avoided. For example, the B-N-B or the N-B-N connectivity is better left depicted as it is.

smaerd - 31-12-2012 at 06:40

That makes a lot of sense thanks for clearing that up. For some reason I forgot that just because it can form 3 bonds does not imply that the empty valence shell does not exist to interact in bonding or resonance. MO theory more than likely makes sense of that and I guess some common-sense I need to adapt. Thanks again.

Glucose Oxidase - 1-1-2013 at 12:35

Does C(NO3)4 exist?
please note that it is NO3 not NO2 so the bond between the radical and the carbon is via oxygen not nitrogen.

kristofvagyok - 1-1-2013 at 12:49

Quote: Originally posted by Glucose Oxidase

Does C(NO3)4 exist?

No.

Also there is an article where they mention it... Also they made reactions with it, but I'm skeptic, because there is no prep mentioned anywhere.

Khisamutdinov, G. Kh.; Slovetskii, V. I.; L'vova, M. Sh.; Usyshkin, O. G.; Besyrozvannyi, M. A.; Fainzil'berg, A. A.
Izvestiya Akademii Nauk SSSR, Seriya Khimicheskaya, 1970 , p. 2397 - 2399

[Edited on 1-1-2013 by kristofvagyok]

barley81 - 2-1-2013 at 13:42

I found an eBay seller in Poland who is offering 250mL of benzaldehyde for around $24 (including the shipping cost). Is this a good price for benzaldehyde? Would it arouse suspicion if I purchased it?

learningChem - 3-1-2013 at 22:40

Is it possible to go (easily) from formaldehyde to formic acid?

Hexavalent - 3-1-2013 at 22:45

In theory, yes. Formaldehyde (or the solution of it, formalin) is an aldehyde, and consequently can be oxidized to formic acid (a carboxylic acid) by refluxing with an oxidant, such as acidified potassium dichromate.

In practice, not quite. I've spoken with DJF90 previously and it seems that formic acid also has a tendency to be oxidize (it itself being a good reducing agent) into carbon dioxide and water, so one would have to be very careful when preparing it via this route and I would imagine that controlling it would be somewhat difficult and problematic.

[Edited on 4-1-2013 by Hexavalent]

learningChem - 5-1-2013 at 13:14

Thanks Hexavalent

[Edited on 5-1-2013 by learningChem]

learningChem - 5-1-2013 at 16:09

What about using the canizzaro reaction? formaldeyde -> methanol + formate?

barley81 - 6-1-2013 at 15:05

You are right. The Cannizzaro reaction would produce sodium formate and methanol. You could isolate the formate and distill it with sulfuric acid (look for an actual procedure, this may not be ideal) If you want a lot of formic acid, you can buy it from Dudadiesel.com. It's not very expensive ($12 a liter or so).

Hexavalent - 6-1-2013 at 15:44

Quote: Originally posted by barley81

You are right. The Cannizzaro reaction would produce sodium formate and methanol. You could isolate the formate and distill it with sulfuric acid (look for an actual procedure, this may not be ideal) If you want a lot of formic acid, you can buy it from Dudadiesel.com. It's not very expensive ($12 a liter or so).

In theory it should work, the equilibrium being shifted to the right due to the formation of a weaker acid.

If you don't want to buy it online, it is sold in many hardware stores as a limescale remover. I bought a sample from a local store a few months ago, and found it to be of 40% concentration, but remarkably pure (clear liquid, no colour at all, very few trace impurities).

Hexavalent - 7-1-2013 at 09:01

Question: why do filtrates tend to froth/foam during vacuum filtrations?

[Edited on 7-1-2013 by Hexavalent]

Vargouille - 7-1-2013 at 12:50

I believe that it is due to the release of gases trapped in the filtrate. The movement of the liquid is part of what releases the trapped gases, akin to how shaking a soda bottle results in foaming. While mentioning soda, recall that putting mentos in soda releases the gases because they provide nucleation sites. The filter paper or frit on your funnel will do the same thing, to a greater or lesser extent, and the application of vacuum aids the movement of the gases out of the filtrate. These gases carry some of the liquid along for a short while, making the foam.

Hexavalent - 7-1-2013 at 14:02

Why foam, though, as opposed to simply bubbling out?

learningChem - 7-1-2013 at 14:19

I tried the canizzaro reaction on ~23 grams of formalin (allegedly 40%) - I distilled ~3.6g methanol out of it and got a sludge of crystals which I'm not yet sure how to work up. I guess it can be done but yields are not going to be too good...

I'm not sure if I can buy formic acid here - I thought the local nazis had banned it for peasants like me, but now I'm not sure - I'll have to check again.

Vargouille - 7-1-2013 at 14:45

The lack of formation of large bubbles is likely jointly caused by a low amount of atmospheric gases trapped in common filtrates (as opposed to carbonated beverages) and the relatively rapidity of nucleation. The speed of most filtrations means that smaller bubbles don't have the time to coalesce into larger ones.

Bronstein - 8-1-2013 at 00:05

Does anyone know where to find the melting point for 2,4-dimethoxyphenethylamine hydrochloride? I can't seem to find it anywhere. I have looked in the printed version of Beilstein, Combined chemical dictionary and The shulgin index, and searched google. I have found some articles where they make 2,4-dimethoxyphenethylamine, but none of them isolate it as the hydrochloride.

The only thing I can seem to find is the melting point of the picrate, the hydrogenoxalate and the boiling point for the freebase.

[Edited on 8-1-2013 by Bronstein]

Nicodem - 8-1-2013 at 13:25

Quote: Originally posted by Bronstein

Does anyone know where to find the melting point for 2,4-dimethoxyphenethylamine hydrochloride?

Beilstein gives these two references for the mp of the hydrochloride:

mp 149-150 °C, Bailey et al., Journal - Association of Official Analytical Chemists 1974, 57, 70.

mp 159 °C, Kappe, Armstrong, Journal of Medicinal Chemistry 1965, 8, 368-372.

learningChem - 9-1-2013 at 13:03

How does MnO2 oxidize toluene to benzaldehyde? Do the O atoms from the Mn oxide go to the toluene or is the oxide a catalyst of sorts?

Simbani - 9-1-2013 at 19:06

New Question:
Is it possible to ball-mill silicon powder (starting with, say 100mµ)? I mean this stuff is quite hard (6,5 vs. 2,75 for Al) and I have no clue on the grindability in a (steel vessel) ballmill.

Hexavalent - 10-1-2013 at 08:36

Quote: Originally posted by learningChem

How does MnO2 oxidize toluene to benzaldehyde? Do the O atoms from the Mn oxide go to the toluene or is the oxide a catalyst of sorts?

The oxidation of toluene is usually done using KMnO₄, which could produce benzaldehyde in theory, but would be exceptionally difficult to control: for this reason, this concept usually produces benzoic acid and manganese dioxide.

Perhaps this is what you mean?

[Edited on 10-1-2013 by Hexavalent]

learningChem - 10-1-2013 at 12:01

Hexavalent,

Toluene can be converted into benzaldehyde using MnO2. See for instance

http://www.google.com/patents/US613460

Reaction of toluene and MnO2 in ~60% sulphuric acid --> benzaldehyde. My question is : where does the O come from - what happens to the MnO2?

-------

I'll put this in the benzaldehyde sticky thread too...

[Edited on 10-1-2013 by learningChem]

barley81 - 10-1-2013 at 15:12

MnO2 (Mn IV) + 4H+ + 2e- ---> Mn2+ + 2H2O
Manganese dioxide is converted into manganese II sulfate.

learningChem - 10-1-2013 at 16:32

Ahh - thanks barley81

What about the 'common' equation? Something like this?

MnO2 + H2SO4 + Ph-CH3 --> MnSO4 + Ph-CHO + H2O + H2

Vargouille - 10-1-2013 at 16:42

Oh, here I was thinking you were looking for the reaction mechanism.

The balanced ionic equation, unless I've made an egregious error, is:

2MnO2 + 4H⁺ + PhCH3 -> PhCHO + 2Mn⁺² + 3H2O

[Edited on 11-1-2013 by Vargouille]

learningChem - 10-1-2013 at 16:55

2MnO2 + 2H2SO4 + PhCH3 --> PhCHO + 2MnSO4 + 3H2O

?

Vargouille - 11-1-2013 at 02:57

Yup. Only difference is that the sulfate is ignored because it doesn't take part in the reaction.

bahamuth - 13-1-2013 at 10:36

Was wondering if anyone here know where or how I might be able to do a "bulk" search of CAS numbers?

I have a long list of CAS numbers that I need the names for, and I would like to just feed them into a database search and get a text file out or similar.

I know ChemACX can do it but haven't been able to access their site today..

mr.crow - 14-1-2013 at 21:44

Quote: Originally posted by bahamuth

Was wondering if anyone here know where or how I might be able to do a "bulk" search of CAS numbers?

I have a long list of CAS numbers that I need the names for, and I would like to just feed them into a database search and get a text file out or similar.

I know ChemACX can do it but haven't been able to access their site today..

Is this for that ebay auction? I just cut-n-pasted them all from a spreadsheet into the sigma aldrich website! Its quite picked over already

bahamuth - 15-1-2013 at 13:04

Quote: Originally posted by mr.crow

Quote: Originally posted by bahamuth

Is this for that ebay auction? I just cut-n-pasted them all from a spreadsheet into the sigma aldrich website! Its quite picked over already

Got me there. But I am looking for unusual stuff, molecular biology stuff in unopened containers. Mostly buffer reagents and the like.

Saw that auction a very long time ago but only recently got my hands on some money to procure some of the stuff I need.

Anyways, did a bulk search with a trial ChemACX and opened the resulting exported file in ChemBioFinder Ultra to browse both the names and the structures.

Btw, how did you manage to bulk interrogate all those CAS numbers at Sigma A, or did you input one and one.....?

mr.crow - 16-1-2013 at 08:15

Just Ctrl-C, Ctrl-V! I tired to write a scraper program in Python but it needs cookies or something to work.

kingkey24 - 22-1-2013 at 05:06

How do I gain access to the reference section forum?

ScienceSquirrel - 22-1-2013 at 05:30

Quote: Originally posted by kingkey24

How do I gain access to the reference section forum?

Send a U2U message to Polverone requesting access.

Units for hygroscopy?

Hexavalent - 30-1-2013 at 09:13

Is there a unit for measuring/quantifying the hygroscopy of a compound? Perhaps something like moles(water)/minute?

[Edited on 30-1-2013 by Hexavalent]

Vargouille - 30-1-2013 at 13:55

Apparently there is a source (Lang's Handbook of Chemistry) that states the hygroscopy of certain compounds in terms of grams water per gram of material. There is also the concept of a "critical relative humidity", which describes the relative humidity below which a compound will not absorb water from the air.

Eddygp - 30-1-2013 at 14:33

Should copper powder react with fluorescein?

Vargouille - 30-1-2013 at 15:10

I can't imagine that it would, at least, not on a reasonable timeline. Copper isn't likely to even bat an eye at the ketone, ether, or alcohol moieties, and the carboxyl group isn't likely to do much except in the presence of atmospheric oxygen. Even then, fluorescein's pKa is only 6.4, so the reaction would be quite slow.

Simbani - 1-2-2013 at 09:08

Assumed I would like to distill (heavy, 6-12 C-atoms) naphta, what would be the best way to clean my flasks and condenser afterwards? I made some experiences with a beaker full of thsi stuff and it was a pain to get it out there -.- I didn´t got it really clean till now.
Now what would be the easiest way to clean contaminated glassware, it should be easy to obtain tho. Is there some way to oxidize it, are they reactive enough?

elementcollector1 - 1-2-2013 at 09:30

Quote: Originally posted by Simbani

Assumed I would like to distill (heavy, 6-12 C-atoms) naphta, what would be the best way to clean my flasks and condenser afterwards? I made some experiences with a beaker full of thsi stuff and it was a pain to get it out there -.- I didn´t got it really clean till now.
Now what would be the easiest way to clean contaminated glassware, it should be easy to obtain tho. Is there some way to oxidize it, are they reactive enough?

Depends. For nonpolar substances, I'd recommend an initial cleaning with a nonpolar, evaporating solvent (such as gasoline, and away from any source of heat, sparks, or open flame). Acetone works well in all situations, as does water in most. A final wash with water is necessary for all glassware in my opinion, no matter the nature of the contamination.

Simbani - 1-2-2013 at 10:56

I tried acetone too, but it doesn´t work. I don´t know what is all in there accurately, but I know it is a mix of naphta (I think the heavy mix) and pentane. I got the beaker relatively clean with soap-concentrate btw.
Isn´t there a chemical way to clean a distillation apparatus? Like oxidating the little residudes which stick on the surface of the glass after washing? H2SO4/H2O2, H2SO4/HNO3, HNO3/HCl, HNO3/H2O2?

elementcollector1 - 1-2-2013 at 22:02

Piranha solution (I think it's H2O2/H2SO4 conc.) seems to be a good method, or most of the other ones you mentioned.

Ampoule codes

darel - 2-2-2013 at 13:38

I purchased a few small ampoules with a gold band where I assumed it is to be broken. I originally thought it was gold or some thin sheet of metal that when heated would have broken the glass at that point on cooling due to thermal stress. Upon research I found it is color coding.

http://en.wikipedia.org/wiki/Ampoule#Ampoule_codes

http://link.springer.com/article/10.1007%2FBF03007733?LI=tru...

That was all I could really find. Does anyone know or have a table for what these codes are?

Vargouille - 2-2-2013 at 16:01

The bands aren't ampoule codes, they're indicating where the ampoules have been prescored. See here.

smaerd - 7-2-2013 at 16:19

So today in analytical chemistry I did a basic buffer solution preparation lab. Using the henderson hasselbalch equation and a phosphate buffer. I prepared a series of pH buffers like 7.4, 7.6, 7.8, etc. Anyways When I measured the pH of these solutions they were off by like 0.20 for each. I know I should probably just ask my professor but is that acceptable? It seems pretty far off... I think a lot of it may have to do with how moist the phosphate salts were before becoming standards. Maybe also ionic strength?

[Edited on 8-2-2013 by smaerd]

DraconicAcid - 7-2-2013 at 16:59

Quote: Originally posted by smaerd

So today in analytical chemistry I did a basic buffer solution preparation lab. Using the henderson hasselbalch equation and a phosphate buffer. I prepared a series of pH buffers like 7.4, 7.6, 7.8, etc. Anyways When I measured the pH of these solutions they were off by like 0.20 for each. I know I should probably just ask my professor but is that acceptable? It seems pretty far off... I think a lot of it may have to do with how moist the phosphate salts were before becoming standards. Maybe also ionic strength?

[Edited on 8-2-2013 by smaerd]

That's pretty typical. The H-H equation is an approximation (you're assuming that the concentrations that you put in are also the concentrations at equilibrium), and ionic strength, as you say, is also a factor. If you want a particular pH, make an approximate one with your calculated buffer, then adjust with acid or base until it's correct.

gsd - 8-2-2013 at 07:33

Esterification Question

If a mixture of 1 mole Ethanol, 1 mole 2-propanol and 1 mole 1-propanol is esterified with 1 mole of acetic acid then what is the likely distribution of product (and why?)

gsd

[Edited on 8-2-2013 by gsd]

DraconicAcid - 8-2-2013 at 09:22

Quote: Originally posted by gsd

Esterification Question

If a mixture of 1 mole Ethanol, 1 mole 2-ethanol and 1 mole propanol is esterified with 1 mole of acetic acid then what is the likely distribution of product (and why?)

gsd

I'm not sure that propanol would react much slower than ethanol in an esterification reaction, but I do have to wonder what you meant by "2-ethanol."

gsd - 8-2-2013 at 09:34

Quote: Originally posted by DraconicAcid

......

I'm not sure that propanol would react much slower than ethanol in an esterification reaction, but I do have to wonder what you meant by "2-ethanol."

I am sorry. I meant 2-Propanol not 2-ethanol

gsd

DraconicAcid - 8-2-2013 at 09:57

I considered that possibility, but it could also have been 2-chloroethanol, etc., so I didn't want to assume. I would expect the product distribution to depend on the rates of the reaction, and the rate will be determined by the nucleophilicity of the alcohols, and I would expect that to depend slightly on the bulk of the organic group on the oxygen. But I'll let a *real* organic chemist offe a more definitive opinion.

DJF90 - 8-2-2013 at 13:21

gsd: I did an esterification with isopropanol, and it took 18 hrs to reach 94% by HPLC. The methyl ester on the other hand, was formed under identical conditions (apart from substituting iPrOH with MeOH) was complete (98%+) in 3 hrs. I dont know how well this answers your question, but I hope you find it of use. I can't divulge the structure of the starting acid due to IP reasons, but its a substituted phenylacetic acid.

smaerd - 17-2-2013 at 07:44

Does anyone know how to go about doing a hydrolysis of a schiffs base that both it and it's products(aldehyde and amine) are pretty insoluble in water?

I've tried steam distillation from aq. oxalic acid, tried refluxing from aq. sulfuric and neither seem to be near complete reactions. The steam distillation carried over some of the suspected aldehyde which is great although a large amount of unreacted schiffs base remains at the bottom of the boiling flask(2g of reactant to 75mL water carried over). Any tips would be exceptional.

Boron Trioxide - 21-2-2013 at 21:52

Question about Electrolysis:

Problems with Anodes:
Carbon, both graphite and activated carbon seem to share this problem, while they are conductive enough, they do not produce nearly any gas during electrolysis as compared to the cathode which produces a fine continuous stream of bubbles, is there a reason and/or solution?

Thank you for your help

elementcollector1 - 21-2-2013 at 21:55

Reason: 2 moles of hydrogen are produced per every mole of oxygen.
Chlorine is barely visible, and only occurs at high concentrations of chloride salt.
Sulfate simply doesn't bubble out as SO2 at all.

AJKOER - 22-2-2013 at 19:54

Quote: Originally posted by d010060002

The strongest concentration of ammonia that is pure and easy to get was 10%. A lot of the recipes I'm looking at require a more concentrated form. Has anyone concentrated ammonia with good results using an easy method (boiling, freezing at a reasonable temperature).

Two ideas:

1. Use Ammonium carbonate in place of aqueous NH3.

2. CaCl2.xNH3 where x = 2, 4 or 8.

elementcollector1 - 25-2-2013 at 08:14

Can you heat a RB flask effectively on a flat hotplate? I'm thinking of doing the synth of K metal at my school, and want to use ground glass - but I have no Erlenmeyers!
Alternatively, will rubber joints affect things?

DraconicAcid - 25-2-2013 at 09:18

Quote: Originally posted by elementcollector1

Can you heat a RB flask effectively on a flat hotplate? I'm thinking of doing the synth of K metal at my school, and want to use ground glass - but I have no Erlenmeyers!?

You can put a hot water bath on the flat hotplate, or an oil bath if you need a higher temperature.

Question

Organikum - 25-2-2013 at 18:18

Would the use of formic acid salts of amines as ammoniumformate for example in the condensation of an aromatic aldehyde with a nitroalkene as benzaldehyde and nitroethane, be problematic in any way in special possible esterification?

Always acetic acid, aren't there so much more organic acids? Thats how it started and now I am up to explore the possibilities. This condensation looks like a good place to start with I thought and now I am looking for traps as my basic chemical is bad and I tend to overlook stuff which is obvious for somebody who had chemistry in school for years.

If you know about sidereactions with formic acid/formate (or other organic acid, citric, oxalic and tartaric are next) pls tell.

regards
/ORG

PS: I start with ammoniumformate for I found an old HIVE posting by Hest which indicates he used it. Regarding his english and other things it might be just a error/mismatch thats well possible.

AJKOER - 14-3-2013 at 09:30

Quote: Originally posted by LiHMDS

Hi everyone!
I have a slightly stupid problem with some household - today I found my bottle of few kg bromine is pretty closed - glass stopper is stuck in this fn bottle. I tried to carefully knock, turn etc, nothing helps. Is there any 'smart' way to open this joint taking on discount all situation?

[Edited on 1-7-2012 by LiHMDS]

Freeze the whole bottle.

Carefully pour boiling water on just the joint.

Caution: The glass could crack, releasing come Br2, so do outside and wear appropriate safety gear.

Vargouille - 14-3-2013 at 12:11

Or, for a less dangerous way: freeze the bottle, then warm up the joint with a rag soaked in hot water from the tap. Wiggle the stopper back and forth as you do so, but gently. No bromine should be airborne, but I would put a fan on just to be safe.

dasgoose21 - 19-3-2013 at 09:40

How would I go about precipitating TACN from a very saturated solution of itself? I'm not very fond of the idea of putting it in a desicator for a few days and then further drying out the paste with expensive acetone, so I am looking for a different way to pull the Tetra out of solution without going through a lot of trouble.

Eddygp - 30-3-2013 at 09:41

What is the easiest, safest and most useful way for a home chemist to test for urea in a solution which might contain NH3 too?

plante1999 - 30-3-2013 at 09:42

To test for urea, use oxalic acid, a precipitate will form if urea is present, ammonia do not interfere.

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