Sciencemadness Discussion Board - The short questions thread (3)

The short questions thread (3)

Pages: 1 2 3 4 5 .. 8

Panache - 7-3-2011 at 18:20

Are amides considered proton/hydrogen acceptors, specifically i was surprised that on this page, http://www.drugbank.ca/drugs/DB00711, diethyl carbamazine was listed as having a hydrogen acceptor count of two?
Obviously one is from the tertiary amine however the other one? Perhaps i am misinterpreting something

Morgan - 16-3-2011 at 06:32

Anybody made any of this material?
http://www.csmonitor.com/Business/Latest-News-Wires/2011/031...

mr.crow - 16-3-2011 at 09:37

Can Mg ribbon be used for Grignard reagents or are they vastly inferior to turnings?

Magic Muzzlet - 16-3-2011 at 11:20

Quote: Originally posted by mr.crow

Can Mg ribbon be used for Grignard reagents or are they vastly inferior to turnings?

Yes it can, if the metal is of adequate purity. All you need to do is chop it up a bit, then wash it with dilute acid etc etc and it will work. I have done this.

mr.crow - 16-3-2011 at 11:48

Good! Turnings seem like a pain to get shipped but Mg ribbons are very common

Waffles SS - 18-3-2011 at 10:16

I decide to produce mercury naphthalene sulfonate from sodium naphthalene sulfonate
I want to react mercury nitrate with sodium 1-naphthalene sulfonate(this is good way for produce it?)
I need solubility of sodium 1-naphthalene sulfonate and mercury-naphthalene sulfonate in water
Someone can help me?

Morgan - 18-3-2011 at 14:03

Anybody familiar with this boring process?

Interesting Facts About Platinum (Scientific American article from the 1800's)
'How are they bored?' Ah, sir, you must excuse me that I do not tell you that."
http://chestofbooks.com/crafts/scientific-american/sup3/Inte...

[Edited on 18-3-2011 by Morgan]

12332123 - 19-3-2011 at 18:04

Does anyone have an opinion on whether alcoholysis of aluminium sulphide is a viable procedure for the synthesis of mercaptans (thiols)?

madscientist - 20-3-2011 at 07:16

You'll get aluminum alkoxide and hydrogen sulfide.

watson.fawkes - 20-3-2011 at 09:35

Quote: Originally posted by Morgan

Anybody familiar with this boring process?

Not specifically, but it does remind me of techniques used in horology (watch-making) for making the holes at the center of tiny gears.

redox - 20-3-2011 at 12:07

Does anybody know of a minimum-boiling azeotrope that boils at STP? (As if, one liquid is poured into another and the mixture starts boiling)

Morgan - 20-3-2011 at 14:08

Quote: Originally posted by watson.fawkes

Quote: Originally posted by Morgan

Anybody familiar with this boring process?

Not specifically, but it does remind me of techniques used in horology (watch-making) for making the holes at the center of tiny gears.

I ran across this article, another very old one.
Invisible Platinum Wire
http://books.google.com/books?id=5KTmAAAAMAAJ&pg=PA193&a...

And this ...
The Fabrication of Ultrafine Platinum Wire
http://www.platinummetalsreview.com/pdf/pmr-v30-i1-027-027.p...

[Edited on 20-3-2011 by Morgan]

Panache - 24-3-2011 at 00:17

perhaps this thread should be renamed the short question and answer thread, as i feel people are unaware that questions are posted in this thread so other people can post the answers.

Morgan - 24-3-2011 at 08:15

As I was unfamiliar with the platinum wire process, I hoped someone might have some first-hand knowledge of how they do it today, the state of the art. I mentioned a few things I was able to dig up a few days later after posting the question, but it's still lacking. I was hoping to edge the topic forward, a thousand pardons.

Panache - 24-3-2011 at 14:54

Quote: Originally posted by Morgan

Anybody made any of this material?
http://www.csmonitor.com/Business/Latest-News-Wires/2011/031...

My response to that is wtf are you doing on that site?????
lol just kidding but christian and science are as oxymoronic as it comes.

Morgan - 24-3-2011 at 18:41

I was interested in knowing if the new plastic was as good at holding pressure as PET. Every March marks another year my two liter bottles have held pressure. In 1992 I pressurized these bottles and still today they are hard enough to stand on, still quite taunt - 19 years. At the time I was planning to make a boat.
As for reading the Monitor, maybe they are starting to come around. ha
"Miller may have gotten the primitive atmospheric composition wrong. But he was right about volcanic plumes. Once again, some of his classic experiments offer reason to believe that lightning flashing through primordial gases may have provided the energy to form chemical precursors of organic life."
http://www.csmonitor.com/Innovation/Tech/2008/1028/a-second-...

Making Sulfuric Acid,Hydrochloric Acid and copper (I) chloride

symboom - 2-4-2011 at 17:09

I think maybe this more belongs here it is kinda short question
http://www.sciencemadness.org/talk/viewthread.php?tid=15942

12332123 - 7-4-2011 at 05:18

Would DMSO be a suitable alternative to THF for reductive amination of amines with STAB and formaldehyde?

kuro96inlaila - 9-4-2011 at 04:24

These make me confuse:

Yesterday,I decided to conduct a couple of experiments with my sperrylite (PtAs2) mineral.

1.I put about half a gram of PtAs2 in nitric acid.Then I heat it.Sperrylite react with warm concentrated nitric acid producing a yellow solution.Then I add few milimeter of ammonium chloride solution,hoping that ammonium hexachloroplatinate will be precipitated.But nothing happen.

I guess ammonium hexachloroplatinate only will be precipitated if ammonium compounds solution is added to chloroplatinic acid.While I might have platinum nitrate in the yellow solution.So,I add magnesium ribbon,hoping that platinum metal precipitated out.It first bubbled,then a brownish orange compound start to precipitated. Here is the picture of it:

I can't figure what compound it is,so I set it aside and conduct other experiment:

2.I react PtAs2 with aqua regia,then I add ammonium chloride.But no ammonium hexachloroplatinate precipitated.Then I add sodium bicarbonate to the solution.It first bubbled,releasing carbon dioxide.Then a brownish orange compound start to precipitate.The colour is similar to what was precipitated in the other experiment.Here is the picture of it:

So,what is these thing actually?I guess it would be either platinum dioxide monohydrate or sodium platinum chloride.But I also doubt it is an arsenic compound.

And,if this is platinum's compound,roasting it should turn it to platinum metal right?

Nicodem - 12-4-2011 at 09:19

Just so you don't complain that questions don't get answered.

Quote: Originally posted by Panache

Are amides considered proton/hydrogen acceptors, specifically i was surprised that on this page, http://www.drugbank.ca/drugs/DB00711, diethyl carbamazine was listed as having a hydrogen acceptor count of two?
Obviously one is from the tertiary amine however the other one? Perhaps i am misinterpreting something

What you are misinterpreting is H-bonds donors/acceptors and confusing them with protonation sites (property of bases). A tertiary amide group, or in this case a N,N,N',N'-tetralkylurea, is a acceptor of the hydrogen bond. A secondary and primary amide group (or tri-, two- or mono- alkylureas, or urea itself) are both, donors and acceptors of H-bonds. But unlike amines, they are not bases as far as aq. chemistry is concerned (most amides are only just about as basic as water, thus protonation in water is nearly negligible).
Similarly, tertiary amines are only H-bond acceptors, but not donors. Ammonium ions, secondary and primary amines, and ammonia, are both, donors and acceptors (though weaker donors than sec. or prim. amides).
Thus, diethylcarbamazine has two acceptors sites for H-bond formation, but zero donors.

UKnowNotWatUDo - 13-4-2011 at 06:38

Speaking of amides, maybe you could answer this question also Nicodem (or anyone else I suppose). There is a synthesis I came across that involves taking a carboxylic acid and a primary or secondary amine and refluxing them in toluene with a catalytic amount of boric acid added (about 0.01mol eq.). A dehydration reaction follows and the water is collected in a Dean-Stark trap, giving excellent yields of the amide. My question is: in the case of a low boiling point amine, a gas at STP, how would you modify this reaction to keep the amine in the reaction vessel? Would the carboxylic acid protonate it to the ammonium salt?Or would you simply be limited by the solubility of your amine in the solvent at reflux temperature? I don't have any particular examples in mind, this is just generally speaking.

And here is the paper I'm referring to:

http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=v81...

Nicodem - 19-4-2011 at 12:53

I don't think using volatile amines is much of a problem in that type of amidations. Methylamine might be somewhat problematic, but not due to volatility (you can always put a septum with a baloon on top of the reflux), but because it partitions much better in water compared to toluene, so lots of it might end up concentrating in the few mL of water collected in the Dean-Stark. This can be surpassed by using some NaOH pellets in the Dean-Stark to absorb the water as it collects. But then again you don't really need an excess of the amine in that procedure so the partial pressure of any amine, including the volatile ones, will be very low in the vapour phase (ammonium carboxylates do dissociate easily, but the equilibrium in most cases does not allow for much free amine).
I used these boric acid and phenylboronic acid catalysed amidations a few times and I must say they work like a charm in most cases. I highly recommend these methods over any coupling reagent whenever applicable (when using robust substrates). There are plenty of other references, besides that Org. Synth., by the way.

UKnowNotWatUDo - 19-4-2011 at 19:50

Thanks for the two cents, it's much appreciated.

Will pyrogallol reduce copper?

bbartlog - 24-4-2011 at 17:25

Will pyrogallol in aqueous solution reduce Cu++, and if so, will it reduce it to Cu+ or all the way to elemental copper? Is it pH dependent?

Chlorination with CuCl2

experimenter - 30-4-2011 at 03:27

Trying to perform chlorinations using CuCl2 in microwave oven, I stumbled across the following observation:

sodium acetate in water, when reacted with cupric chloride (CuCl2) gave a white precipitate (CuCl I suppose). This means that a reaction happened but I highly doubt that it chlorinated the acetate (in order to give the sodium monochloroacetate).
The ph of the solution after reaction was around 5.

What do you think, is it possible that CuCl2 can chlorinate the sodium acetate? (supposedly CuCl2 can chlorinate carbonyl compounds)

I know about the usage of LiCl catalyst, but in that case, it wasnt needed, thats why I'm suspicious that chlorination hapenned.

DJF90 - 30-4-2011 at 05:31

I don't know about acetate, but both CuBr2 and CuCl2 will a-halogenate aldehydes and ketones. If Methanol is used as the solvent then the a-chloro-dimethylacetal is furnished.

Sedit - 30-4-2011 at 12:29

What is the easiest Amine protecting group that a home chemist can get there hands on?

Its sort of besides the point but I want to attempt to protect the amino grouping on glycin in an attempt to form Ethylene Diamine thru the Kolbe electrolysis. This normally produces low yeilds but I wish to see if higher yeilds can be had by protecting the amine. However all I read gives unreasonable protecting groups that are out of my range.

bbartlog - 30-4-2011 at 15:37

Quote: Originally posted by experimenter

...sodium acetate in water, when reacted with cupric chloride (CuCl2) gave a white precipitate (CuCl I suppose).

What concentration were your solutions? Which was in excess? I would suspect NaCl rather than CuCl in this situation, if the solutions were concentrated. Did you do any testing to verify that the precipitate was CuCl? You could just set it out (still wet) on some filter paper and see if the air turns it green; the oxidation proceeds pretty rapidly. Anyway, it's just not clear to me what would be getting oxidized here to reduce the copper, which is why I have a hard time seeing CuCl as likely. Heck, in semidilute solution NaCl (which you would have) should solubilize CuCl anyway, it probably wouldn't precipitate even if present...

Sedit - 30-4-2011 at 19:05

I need to convert an amine from an electron releasing group to an electron withdrawling group so that it does not stabilize the formed radicals.

Im considering protecting the amine some how, as seen in my above post, which I feel will help this some what but conversion of the Amine by oxidation to an Nitro group, converting it to an EWG, may be a better option.

Does any one have any other suggestions as to what the best means of converting the amine to some form of EWG would be.

experimenter - 30-4-2011 at 23:55

Thanks for your answers.

Quote:

What concentration were your solutions? Which was in excess?

I can't accurately reply to this now, I must repeat the experiment.

Quote:

Did you do any testing to verify that the precipitate was CuCl? You could just set it out (still wet) on some filter paper and see if the air turns it green; the oxidation proceeds pretty rapidly.

Here is a picture of the precipitate today (one day after filtering), I remember well it wasn't green yesterday, when fresh:

It is baffling me too, how could the reaction be if this is really CuCl. Maybe copper acetate is formed? (+NaCl + CuCl)

bbartlog - 1-5-2011 at 06:56

Hm... very interesting. Looks like you have CuCl, all right! Copper acetate (or one of the copper acetates) would logically form (as well as NaCl) if you precipitated everything by removing the water, but I don't think that explains the change in oxidation state.

experimenter - 1-5-2011 at 09:37

Well, I tried it again with pure sodium acetate salt and it gave nothing.
The first time I used contaminated sodium acetate (it had a ph around of 10) and probably it had some baking soda in excess in it (I made it by neutralizing soda and vinegar).

bbartlog - 1-5-2011 at 12:13

Vinegar might have provided some organic stuff that would reduce Cu++. Excess NaHCO3/Na2CO3 should precipitate CuCO3, but that would be green from the beginning.

Panache - 2-5-2011 at 10:23

if i had a light globe (incandescent) and it was for 120V power supply, if i ganged two in series and plugged them into 240V would they work as normal, ie as i they were plugged individually or in parallel in 120V?
lazy question.

[Edited on 3-5-2011 by Panache]

Panache - 7-5-2011 at 21:01

Quote: Originally posted by Panache

if i had a light globe (incandescent) and it was for 120V power supply, if i ganged two in series and plugged them into 240V would they work as normal, ie as i they were plugged individually or in parallel in 120V?
lazy question.

[Edited on 3-5-2011 by Panache]

the answer is yes apparently, large casino's here in melbourne have themselves wired so that there are always two 120V sockets in parallel and they use globes from america, apparently there is a wattage saving in it somehow that is appreciable over so many light.

bbartlog - 8-5-2011 at 05:48

But is that what you were originally asking? It sounds like they might have cobbled together a pair of bastard 120V sockets from a 240V source, whereas your question is whether you could just put two of the globes in series plugged in to a 240V outlet.

manimal - 10-5-2011 at 14:56

Hai.

Suppose that one is chlorinating an organic liquid, as toluene or similar, with an appropriate setup.

Is it realistic to expect an absorption of Cl2 that is quantitative or nearly so? That is, can the flow rate be regulated to ensure that most all of the gas is absorbed while still maintaining a steady reaction rate, or is offgassing of Cl2 a given?

Regards

Panache - 12-5-2011 at 05:30

Quote: Originally posted by bbartlog

But is that what you were originally asking? It sounds like they might have cobbled together a pair of bastard 120V sockets from a 240V source, whereas your question is whether you could just put two of the globes in series plugged in to a 240V outlet.

i have 10 120V, 1000W projector globes that i wanted to make use of, but couldn't be assed with hasel of needing stepdown transformer everytime i needed to light something, so yes and no, but my very compact 2000W light source is pretty bright, at the moment its on my desk, when i have visitors i say, hey look at this holding my hand next to the lamp, when they lean in i turn on the light.
i think i am contributing.

manimal - 28-5-2011 at 18:38

Why are hydrobrominations frequently done with HBr in glacial acetic acid solvent? What is wrong with acetone?

bbartlog - 28-5-2011 at 19:11

I'm guessing that HBr will cause acetone to form condensation products the same way that HCl would. But there may well be other solvent effects that make specifically acetic acid the solvent of choice; I don't know anything about that.

Picric-A - 30-5-2011 at 07:57

After atempting to reduce 1,3-dinitrobenzene with NaBH4 in methanoic solution I instantly obtained a dark purple colouration, very similar to KMnO4 solution, to the point where it was nearly black.

Does anybody know what is going on here? I cant see why 1,3-benzenediamine would change to this colour?

entropy51 - 30-5-2011 at 08:43

Quote: Originally posted by Picric-A

After atempting to reduce 1,3-dinitrobenzene with NaBH4 in methanoic solution I instantly obtained a dark purple colouration, very similar to KMnO4 solution, to the point where it was nearly black.

Does anybody know what is going on here? I cant see why 1,3-benzenediamine would change to this colour?

I wonder where you got the idea that borohydride will reduce aromatic nitro to amine?? An hour in the library can save a day in the lab. To the best of my limited knowledge you will get a mixture of azo compounds. Typically colored ones.

Here's one reference.

DJF90 - 30-5-2011 at 17:24

Quote: Originally posted by Picric-A

Mixing chemicals without doing the background reading is a decent method of getting yourself damaged. Any textbook on Functional Group Interconversion will tell you that NaBH4 is incapable of reducing nitro groups to amines. If you'd have cared to do a literature search you'd find that in some cases, addition of a metal salt can reduce aromatic and aliphatic nitro groups. But I'll let you read up on that one...

Note that 1,3-dinitrobenzene isn't a particularly nice material. Also note that selective reduction to 3-nitroaniline ought ot be possible with a reagent like sodium sulfide.

Aldol Condensation

darel - 8-6-2011 at 18:17

What is the difference between a mixed aldol and a cross aldol condensation. From what I've read they can be the same aldehyde/Ketone, a different aldehyde/ketone or an aldehyde and a ketone. The only key difference I could find out was crossed aldol condensation only works if there is an Aldehyde that has no alpha hydrogens so it will never form the enol or enolate anion intermediate depending if you are catalyzing with acid or base. That was mentioned from a university undergraduate lab.

http://www.faculty.sbc.edu/jowens/Chem232_S07/Labs/Exp_5_AlC...

Waffles SS - 12-6-2011 at 20:23

What is downward distillation?

Waffles SS - 12-6-2011 at 20:25

There is way for prepare Tin(II) Chloride from Potassium Stannate?

SHADYCHASE54 - 16-6-2011 at 14:52

Hello members,
I have a quick question about ceramic tubing, typically how much would one expect to pay for a a 36'' length, an inch ID relatively thin walled. I am not looking for a supplier quote I am merely looking for someone with experience in buying a similar piece. What would be a fair price for this? I am interested in making a tube furnace kanthal wound and I would preffer a ceramic or alumina heating tube. So I guess what I am asking is, memebers who have bought this type of piece in the past what would you pay? I will thank you in advance for your help you are willing to provide.

Shadychase54

gutter_ca - 16-6-2011 at 16:07

How to activate reagent Aluminum Oxide for column chromatography?

Waffles SS - 24-6-2011 at 23:15

Does Carbonation of Grignard Reagents done at room temp by CO2 gas or it need lower temp(provide this temp by dry CO2)?
RMgX + CO2 -> RCO2H

[Edited on 25-6-2011 by Waffles SS]

UKnowNotWatUDo - 25-6-2011 at 15:21

Anyone have any solubility miscibility data for different solvents with butyl diglyme (diethylene glycol dibutyl ether)? Specifically things like toluene, benzene, methanol, and ethanol. I already know it has very limited solubility in water, as well as water in it. Even any personal experiences in lieu of solid data would be appreciated.

Mixell - 27-6-2011 at 07:45

What is the structure of copper (II) formate? Just as the oxalate (Cu2(HCOO)4)? Or is is Cu(HCOO)2?
Or a completely different structure?

Nicodem - 28-6-2011 at 11:07

Quote: Originally posted by Mixell

What is the structure of copper (II) formate?

The structure of the dihydrate has been determined:
Acta Cryst. 1965, 19, 357-362. DOI:10.1107/S0365110X65003456

barley81 - 28-6-2011 at 20:01

If I only want a vacuum pump for filtration (and maybe distillation) is it better to use an aspirator with a jet pump, the faucet, use a hand pump, or buy a 100$ vacuum pump from Harbor Freight? I am concerned that the local water supply does not have enough pressure to pull an acceptable vacuum on an aspirator.

Sedit - 29-6-2011 at 12:59

What is the simplest Styrene polymerization inhibitor that I can get over the counter?

Its oxygen that catalyses the reaction right? Would a small about of Ethanol or something of that sorts inhibit the polymerization? How about dissolving it in something like Xylene so that the concentration isn't very high?

I just want to attempt pyrolysis of PS and I want to be able to store it in order to perform various experiments with it.

Jor - 29-6-2011 at 13:20

I am not sure but isn't hydroquinone used for this purpose?

AndersHoveland - 29-6-2011 at 13:50

Does anyone know where to buy 1,3-dichloro cyclobutane? Any other 1,3- derivitives would also be suitable.

Magpie - 29-6-2011 at 20:24

Quote: Originally posted by barley81

If I only want a vacuum pump for filtration (and maybe distillation) is it better to use an aspirator with a jet pump, the faucet, use a hand pump, or buy a 100$ vacuum pump from Harbor Freight? I am concerned that the local water supply does not have enough pressure to pull an acceptable vacuum on an aspirator.

For 8 years and a lot of chemistry under the bridge I have only used a hydroaspirator. This has a lot of advantages if you have adequate water pressure. I don't know the minimum acceptable pressure, this would be equipment and task dependent, but I'm assuming 30 psig would do. Mine will draw down to a 50-60mmHg pressure, adequate for all filtrations and all the vacuum distillations to date. If the vacuum distillation procedure specifies 15-20mmHg pressure I have found that my 50-60mmHg will suffice. The product just comes off a little higher in temperature. So I lose a little more product due to decomposition.

So, unless you need a hard vacuum, or your aspirator is going to run for hours and the water consumption would be excessive, why buy a pump? If the pump is oil lubed you'll need a trap or have to change oil frequently.

I wouldn't buy a hand pump when an aspirator is just as cheap, unless water usage is a problem.

So you see, there's no direct answer. The choice very much depends on the individual's circumstances.

smuv - 29-6-2011 at 20:40

I use my fridge pump, before that I used a medical diaphragm vacuum pump. I think it is cheaper to go one of these routes than setting up an aspirator/recirculating system. On the other hand, if you just run it from your faucet, and don't care about the water usage, by all means, go for the aspirator.

The good thing about the aspirator is you don't have to worry about sucking evaporated solvent into the pump oil, but this is mostly preventable w/o a trap by not using extremely high vacuums during filtrations (you can use a bleeder valve to control vacuum level very easily).

Neil - 29-6-2011 at 21:55

I use a plastic aspirator and can easily pull 25" @ 40PSI and 27" @ 70PSI

I also installed a 1-100PSI pressure gauge on the water line so that I can check the water main pressure and plan around its flux. Works fine for me.

barley81 - 30-6-2011 at 13:11

Thank you for the information. I'll buy an aspirator and use it with the garden hose. Would a water pressure of 30 PSI and a metal aspirator work? Or would I need to buy a jet pump or plastic aspirator?

Bolt - 4-7-2011 at 11:44

Does anyone know of a preparation of sulfur trioxide from chlorosulfonic acid? Industrially, chlorosulfonic acid is produced from HCl(g) and SO3.

Mixell - 5-7-2011 at 13:56

I've encountered someone on this site claiming that titanium dissolves better in 9% HCl then in 37%.
Is it true? Because I'm planning to dissolve some titanium flitter.

Bot0nist - 5-7-2011 at 14:03

Maybe do to passivization? Here is a PDF. I just scanned over it but I think the temperature of the acid is key.

http://www.archivesmse.org/vol28_6/2866.pdf

Mixell - 5-7-2011 at 14:41

Holy crap.
Who knew that 30g of titanium flitter and 200 ml of HCl heated to about 80C will continue reacting due to the heat of the reaction. the house smells of HCl. But at least I'm pleased.

Bot0nist - 5-7-2011 at 15:17

Quote: Originally posted by Mixell

Holy crap.
Who knew that 30g of titanium flitter and 200 ml of HCl heated to about 80C will continue reacting due to the heat of the reaction. the house smells of HCl. But at least I'm pleased.

lol, careful. HCl is persistent. *cough* My throat hurts just thinking about it.

Mixell - 5-7-2011 at 15:38

Ok, the reaction is done.
But I got quite strange results:
The solution is quite concentrated and looks black, when a small amount is taken it looks very dark purple-brown.
But when I dilute it, the solution turns quite brown, resembling this shade: http://ontariohottubcovers.com/imgs/colors/color-cedar-brown... .
When I add hydrogen peroxide to it, it first turns slightly yellow (very transparent) and if I add more it turns blood red.
But what troubles me is it does not resemble the purple solution at all, as shown on Wikipedia: http://en.wikipedia.org/wiki/File:TiCl3.jpg .
What could it be (the titanium is 99.5% pure and the HCl is lab grade)?
Also, the solution turns blue on addition of concentrated HCl.

[Edited on 5-7-2011 by Mixell]

Mixell - 5-7-2011 at 20:38

Well, I figured it out, the brown color is due to Ti[IV] contamination.
And the blue color is due to the TiCl6 ion? Is it titanium in its III or IV oxidation state?
Is there any good use for this solution? Or I'll just leave in on my desktop and see how it turns clear?
Shame that it will go to waste.

sternman318 - 12-7-2011 at 08:06

What does the squiggle that is bound the carbon represent?

http://en.wikipedia.org/wiki/Acyl_group

fledarmus - 12-7-2011 at 08:14

That's just a shorthand to say that anything could be connected to the other end of the bond. Rather like an engineering drawing, if you were drawing a closeup of the valve on the end of a pipe, you might make a wavy line across the end of your drawing of the pipe to show that the pipe continued beyond the edge of the page but wasn't important to what you were trying to illustrate.

In organic chemistry it's usually used when your reactions are occuring at a specific group on a rather large structure, like cholesterol. Rather than redrawing the entire cholesterol structure at each step of the reaction, you just draw the piece that is reacting in each step, and the reader can assume that the rest of the molecule remains untouched.

In your specific case, only the part of the molecule you show is considered an "acyl group" - some sort of carbon chain (R) attached to a carbonyl (C=O) attached to - something. An acyl group is just a fragment. If the "something" is an OH group, your acyl group is part of a carboxylic acid, if an OR' group it is part of an ester, if an NHR' or NR'R" it is part of an amide, etc.

[Edited on 12-7-2011 by fledarmus]

Mixell - 12-7-2011 at 19:48

What is the best way to dissolve nickel powder?
Concentrated sulfuric and nitric acid don't seem to do much, it reacts vigorously at first, but then the reaction stops.

[Edited on 13-7-2011 by Mixell]

sternman318 - 12-7-2011 at 20:21

Thank you fledarmus!

And Mixell, have you tried diluting it? I have no idea if this is the case, but I know that concentrated nitric acid doesnt dissolve copper because it forms a passivation layer. Diluting prevents the formation in that case.

Mixell - 13-7-2011 at 07:36

I tried diluting both of the acids, but nothing seem to work, may be I just didn't find that "magic concentration".

[Edited on 13-7-2011 by Mixell]

OK to boil sodium chlorate liquor electrolytically?

m1tanker78 - 15-7-2011 at 16:32

...instead of boiling after filtration? In other words, is there anything wrong with simply bumping the voltage/current up at the end of a run to boil the liquor, then filter and proceed as usual (aside from a little extra anode erosion)? The solution is maintained at 80-85 degrees C anyway and acidified throughout. I don't see how 2 or 3 extra amps would hurt but don't want to screw up this batch.

Thanks

jwarr - 18-7-2011 at 18:55

I'd like to remove carbonyl adulterants from my denatured ethanol. Would mixing it with sodium hydroxide and then distilling be sufficient?

jwarr - 18-7-2011 at 19:00

I'd like to remove carbonyl adulterants from my denatured ethanol. Would mixing it with sodium hydroxide and then distilling be sufficient?

AndersHoveland - 18-7-2011 at 19:34

Quote: Originally posted by jwarr

I'd like to remove carbonyl adulterants from my denatured ethanol. Would mixing it with sodium hydroxide and then distilling be sufficient?

Yes, not aware of any reason it would problematic. Just be aware that the sodium acetate will easily melt at higher temperatures, and you would want to stop distillation as soon as all the alcohol is out, otherwise some of the sodium acetate may decompose, giving off acetone vapor.

hkparker - 21-7-2011 at 17:52

I do believe ethanol reacts with NaOH to make sodium ethoxide...

http://www.sciencemadness.org/talk/viewthread.php?tid=2656

AndersHoveland - 21-7-2011 at 17:57

Quote: Originally posted by hkparker

I do believe ethanol reacts with NaOH to make sodium ethoxide...
tid=2656

Only if the ethanol is not diluted with water, and solid NaOH is used. Obviously, sodium ethoxide reacts with water.

hkparker - 22-7-2011 at 00:15

oh course, it reacts to reform the reactants...

solo - 22-7-2011 at 22:15

In the enclosed study it the TEA they use mean triethylamine or the triethanolamine.....i'm confused since they sometimes use the same abbreviation for both.....but i feel it may be the triethylamine since it may be removing the hydrogen chloride from the reaction....any thoughts.....solo

Attachment: An Insight of the Reactions of Amines with Trichloroisocyanuric Acid.pdf (160kB)
This file has been downloaded 3017 times

Polverone - 22-7-2011 at 23:10

Quote: Originally posted by solo

In the enclosed study it the TEA they use mean triethylamine or the triethanolamine.....i'm confused since they sometimes use the same abbreviation for both.....but i feel it may be the triethylamine since it may be removing the hydrogen chloride from the reaction....any thoughts.....solo

It is triethylamine according to the diagram of reaction scheme 2.

Panache - 27-7-2011 at 02:02

does anyone know what is added to automatic dishwashing liquids (these are rarer, as opposed to powders and tablets) to stabilise the carbonate, bicarbonate solutions, or is it something completely different?

Also downward distillations is when the condenser is lower than the still head.

Lambda-Eyde - 27-7-2011 at 12:42

Is the -CH2OH substituent activating or deactivating? I.e. will benzyl alcohol, on treatment with an alkyl chloride and FeCl3 give the meta or ortho/para product?

I'm guessing meta, since benzaldehyde and benzoic acid both have deactivated rings, but I'm asking since I didn't find the particular substituent (-CH2OH) in my organic chemistry book. Google didn't turn up anything relevant either.

If it is deactivating, i.e. meta-directing, is there any way to reversibly modify the -CH2OH substituent reversibly to an activating, ortho/para-directing one? Would converting the alcohol to a benzyl ether change the reactivity of the ring? Also, would this reaction be a competing side reaction in a Friedel-Crafts alkylation? The Williamson ether synthesis is conducted in a basic environment, to what extent does this reaction occur in the conditions of a F-C alkylation?

Hardness of Fe++

smuv - 3-8-2011 at 22:43

Why is Fe(II) considered a borderline acid?

It has lower Electron affinity, Ionization energy, electronegativity and chemical hardness than both Cd(II) and Hg(II), both of which are weak acids.

Also anyone have any quantitative hardness data for thioethers? At this point I am just extrapolating from H2S.

--Thanks

p.s. data: http://pubs.acs.org/doi/abs/10.1021/ja00364a005

@ lambda, somewhat deactivating. The electronegativity of oxygen pulls electron density from the ring. There is also no way for the oxygen to contribute its lone pairs to the ring. So slightly meta directing. I have no data for product distributions but I would expect a lot of para product as well.

fledarmus - 4-8-2011 at 04:18

Quote: Originally posted by Lambda-Eyde

Is the -CH2OH substituent activating or deactivating? I.e. will benzyl alcohol, on treatment with an alkyl chloride and FeCl3 give the meta or ortho/para product?

I'm guessing meta, since benzaldehyde and benzoic acid both have deactivated rings, but I'm asking since I didn't find the particular substituent (-CH2OH) in my organic chemistry book. Google didn't turn up anything relevant either.

If it is deactivating, i.e. meta-directing, is there any way to reversibly modify the -CH2OH substituent reversibly to an activating, ortho/para-directing one? Would converting the alcohol to a benzyl ether change the reactivity of the ring? Also, would this reaction be a competing side reaction in a Friedel-Crafts alkylation? The Williamson ether synthesis is conducted in a basic environment, to what extent does this reaction occur in the conditions of a F-C alkylation?

This is a slightly activating, ortho-para director. It acts very much as a similar alkyl substituent would act (eg n-propyl). Because the oxygen can't actually contribute electrons directly into the ring, it only acts as an electron withdrawing group on what would otherwise be a methyl group.

A Japanese patent (JP08176049) measured the products of iodination on the methyl ether and obtained a 1:0.5:2.6 ratio of ortho, meta, and para substituted compounds.

smuv - 4-8-2011 at 19:33

Generally it is considered a moderately deactivating meta director. I cannot read the patent you post, but it may be a special case.

DDTea - 5-8-2011 at 06:30

When taking an object out of a muffle furnace and allowing it to cool in a desiccator, should the vacuum be applied as it cools?

I don't think that it should. Objects cool by two mechanisms: conduction and blackbody radiation. By removing the air from the desiccator, much of the conductive medium is being removed. So the object can only conduct heat away via the desiccator platform or radiate heat (less efficient).

My colleague believes that the vacuum increases the rate of cooling. I feel like doing an experiment to prove my point would be a dick move, and I may have forgotten some thermodynamics. Who is right?

Mixell - 5-8-2011 at 08:16

Well, If the object is indeed a solid that would not sublimate under vacuum, applying more vacuum means less contact with particles that could take heat away from your object, thus lowering the rate of cooling. Although if it does sublimate/evaporate, then the vacuum should increase cooling rate by promoting some of the more agitated particles to escape the object (or liquid) and take away a good amount of heat with them (by lowering the force applied by the air pressure, thus lowering the boiling point/increasing the sublimation rate).

[Edited on 5-8-2011 by Mixell]

watson.fawkes - 5-8-2011 at 08:23

Quote: Originally posted by DDTea

When taking an object out of a muffle furnace and allowing it to cool in a desiccator, should the vacuum be applied as it cools?

"Should" is relative to your goal. If your only goal is speed of cooling, then eliminating a fluid cooling medium (gas or liquid) will make it cool more slowly. That's not nearly the whole story, though. Since thermal emissivity goes up as the fourth power of temperature, radiation completely dominates thermal loss at high temperatures. Fluid cooling will add something, but it may not be worth it.

The best reason to keep it under vacuum is to prevent from reacting with the atmosphere while it's hot. For some materials, this isn't a concern. If you've got materials that will oxidize, or more rarely nitridize, vacuum is perfectly warranted.

Mixell - 5-8-2011 at 08:25

But on a more practical note, I doubt very much that if the object does not sublimate, that the air particles in a closed vessel could relieve much heat from it. The lion share of the heat will be transfered to the desiccator platform. Although the air particles might transfer the heat to the glass casing of the desiccator, allowing some non negligible heat transfer.
Could be quite interesting to check this out.

DDTea - 5-8-2011 at 08:54

Quote: Originally posted by Mixell

...The lion share of the heat will be transfered to the desiccator platform. Although the air particles might transfer the heat to the glass casing of the desiccator, allowing some non negligible heat transfer.
Could be quite interesting to check this out.

That was my reasoning: the air would transfer energy to the glass casing which, in turn, would heat the surroundings and result in a net loss energy (i.e., cooling) in the system (desiccator and its contents). However, watson.fawkes brought up an interesting point about the importance of radiative energy loss at higher temperatures.

In this procedure, we're using the muffle furnace to burn off residual organics on some weigh boats (they appear to be sand--they're used in LECO C/S analyzers) at 1000*C--i.e., raging orange glow.

If, at this temperature, blackbody radiation is the primary means of cooling, then the vacuum issue may be moot--at least until the weigh boats cool below a certain temperature.

watson.fawkes - 5-8-2011 at 09:03

Quote: Originally posted by DDTea

If, at this temperature, blackbody radiation is the primary means of cooling, then the vacuum issue may be moot--at least until the weigh boats cool below a certain temperature.

I'll add only to keep them out of air drafts while cooling. This can thermally shock ceramics and glass enough to break them.

AndersHoveland - 5-8-2011 at 09:25

Quote: Originally posted by Lambda-Eyde

Although chlorination of nitrobenzene with FeCl3/Cl2 leads to the meta isomer, methyl groups are said to be 2,4-directing. So while nitration of nitrobenzene lead mostly to the 3-isomer of dinitrobenzene, when toluene is nitrated only about 5% of the product is 3-nitromethylbenzene.

Quote: Originally posted by fledarmus

The above answer would seem true.

Dibenzyl Ketone

Picric-A - 8-8-2011 at 03:30

Dry distillation of calcium phenylacetate with a current of dry N2 yields dibenzyl ketone in good yields.

I am having problems with what i thought would be a simple procedure of making the calcium salt. Here is what a did;
-1 mole of cacium hydroxide was suspended in 200ml of hot water.
-To the suspension, 2 moles of phenylacetic acid was added and the solution was boiled for 5 mins
-the water was then evaporated. This left me with what seemed to be the phenylacetic acid crystals (smelt like it) and calcium hydroxide powder.

Does anybody know why the two didnt react? What am i doing wrong? I thought it would be siple acid-base chemistry :s

smaerd - 10-8-2011 at 12:35

edit - nevermind I see now

[Edited on 11-8-2011 by smaerd]

redox - 10-8-2011 at 13:49

Quote: Originally posted by Picric-A

Dry distillation of calcium phenylacetate with a current of dry N2 yields dibenzyl ketone in good yields.

I am having problems with what i thought would be a simple procedure of making the calcium salt. Here is what a did;
-1 mole of cacium hydroxide was suspended in 200ml of hot water.
-To the suspension, 2 moles of phenylacetic acid was added and the solution was boiled for 5 mins
-the water was then evaporated. This left me with what seemed to be the phenylacetic acid crystals (smelt like it) and calcium hydroxide powder.

Does anybody know why the two didnt react? What am i doing wrong? I thought it would be siple acid-base chemistry :s

In a different thread, Vogelzang made calcium phenylacetate from phenylacetic acid and calcium oxide instead of the hydroxide. Maybe you can try this.

Mixell - 13-8-2011 at 05:08

Does iodination of hexane is possible at RT? and if yes, how fast the reaction is and what is the favorable species?

theflickkk - 13-8-2011 at 06:25

I don't think that hexane would under iodination at RT. Free-radical sub is the mechanism by which it would probably go by. However, with I2, this is likely to be practically non-existent at RT.

Mixell - 13-8-2011 at 07:03

Bromination proceeds quite fast at RT as I've seen on youtube, I'm afraid that iodination proceeds similarly (the iodine will be dissolved in hexane and will stay like that for a few weeks, so I need to be sure that it won't react at all).

[Edited on 13-8-2011 by Mixell]

theflickkk - 13-8-2011 at 07:39

Bond energy of :
C-H bond approx. 397KJ/mol
Br-Br bond approx. 193KJ/mol
I-I bond approx. 151KJ/mol
C-Br bond approx. 280KJ/mol
C-I bond approx. 240KJ/mol
H-Br bond approx. 366KJ/mol
H-I bond approx. 299KJ/mol

Calculating energies of bonds broken + bonds formed:

For bromination= 397 + 193 - 280 - 366 = -56KJ/mol
For iodination = 397 + 151 - 240 - 299 = +9KJ/mol

Since bromination is exothermic, its likely to be energetically favorable while the opposite would likely apply to iodination.

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