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mewrox99
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[*] posted on 10-6-2010 at 17:13


Why are potassium salts less soluble than sodium salts when the anion is oxidizing



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JohnWW
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[*] posted on 10-6-2010 at 18:08


Probably because of the larger size of the K+ cation, which must make formation of crystal structures on precipitation or crystallization less enthalpy-reducing than with Na+ cations and the same anions. This would be in spite of the larger size of the K+ cation resulting it being less strongly hydrated in solution (which explains why K salts, e.g. KMnO4, are frequently anhydrous in solid crystal form, while the corresponding crystalline Na salts, including the permanganate, are hydrated, or at least have more water molecules than the K salts.).
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[*] posted on 14-6-2010 at 07:33


Quote: Originally posted by mewrox99  
Why are potassium salts less soluble than sodium salts when the anion is oxidizing

The oxidizing nature of the anion has nothing to do with the solubility (I assume you mean aqueous?). Solubility is a property that depends on lots of things. In the case of sodium vs. potassium salts it is mostly about the difference in hydration of the sodium cation vs. hydration of the potassium ion and the energy required to break/form the crystalline phase. Little can be said about the crystal structure of Na/K salts and its energies as it depends also on the anion (there are some useful estimates based on the difference in ion sizes, but I forgot the name of the theory). However, the sodium cation is more acidic than the potassium cation (due to the smaller size of Na(+) at the equal charge). Therefore the sodium cation solvates more strongly (more energetically) with a basic solvent like water. Still, this does not mean that every sodium salt will be more soluble than the corresponding potassium salt.
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[*] posted on 14-6-2010 at 12:28


Has anyone done any 'pyro' reactions with Calcium Hypochlorite

Seeing how it reacts with metallic powders, charcoal, sugar etc




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Lambda-Eyde
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[*] posted on 14-6-2010 at 16:19


I'm not sure if this is the right place to ask, but here goes anyways: Can anyone recommend an introductory book to chromatography? I'm especially interested in thin layer chromatography, not so much in modern methods requiring expensive instruments. The reason for this is that none of my basic chemistry books (two high school level, one university level) even mention TLC, and I would like to follow my organic reactions using TLC. I understand the basic principles, but I would like to learn it more in-depth and learn how to select various solvent systems, interpret results and solve practical problems. Also, it would be nice to have a physical book instead of having to resort to various online guides all the time. I wouldn't mind learning about GC, HPLC and column chromatography either, but that's not my primary interest.

The keywords are: Basic, nice price, focused on TLC and with a practical approach so I can relate it to real experiments.


I'm also looking for a book on basic practical electric "engineering", especially high voltage systems. I would love to build a Tesla coil someday, but I would also like to do so while understanding what I'm doing and not getting killed. I feel that I also need some literature on electricity that is more oriented around practical subjects, but not excluding the mathematics needed. The only knowledge of electricity I have is from a basic physics course (Chapters "Electricity" and "Semiconductor technology"), and I'm having a hard time relating this to real projects.

Again, the keywords are: Basic, nice price, practically oriented and preferably something related to HV projects.


Also, I'm going to shamelessly bump my question posted earlier:


Quote: Originally posted by Lambda-Eyde  
I made 500 ml of 0,1M (NH<sub>4</sub>;)<sub>2</sub>Fe(SO<sub>4</sub>;)<sub>2</sub> solution today by dissolving 19,61 g of the dry powder to 500 ml with deionized water in a volumetric flask. It is very old, however it had the characteristic blue color of Fe(II) it should have before I dissolved it. The solution, however, was more like a murky yellow. Now I'm seeing a brown precipitate, but not much.
Is this due to atmospheric oxidation or must it have been something in the water that caused it? Mohr's salt is said to be stable against air oxidation because of the acidic ammonium ions present in the solution.

Anyhow, would it be okay if I filtered off the fine precipitate and standardized the solution against 0,02M KMnO<sub>4</sub>? The permanganate is standardized against sodium oxalate, a primary standard.

More has precipitated and the solution is now a slight piss-yellow. Is there even any Fe(II) left in solution? I find this quite confusing as this compound is supposed to be stable with respect to aerial oxidation. If I remember I can get around to take a picture of it tomorrow.


Any input is GREATLY appreciated! :)
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[*] posted on 15-6-2010 at 09:36


.... a two part Question....
1. how can one degrade or oxidize a secondary aliphatic amine to its schiff base or imine,
2. if by using a single isomer of the secondary aliphatic amine , will the degradation or oxidation of the compound recimize it once its reduced back to the secondary aliphatic amine?

.......an interesting project, would appreciate some input, i have tried the benzaldehyde reflux , and reduced it back to the amine with NaBH4 but have noted no change ...that is chirality did not change, also tried the HI boil to destabilize the chiral center but no recimization occurred....solo




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[*] posted on 15-6-2010 at 10:59


Quote: Originally posted by Lambda-Eyde  
I'm not sure if this is the right place to ask, but here goes anyways: Can anyone recommend an introductory book to chromatography? I'm especially interested in thin layer chromatography, not so much in modern methods requiring expensive instruments.

At least one book fully dedicated to TLC is freely available in the grey zones of internet: Thin Layer Chromatography 1 (Jork et al, 1990). It is available in the DJVU format so you will need a viewer (there are two or more freeware ones available).

Quote: Originally posted by solo  
.... a two part Question....
1. how can one degrade or oxidize a secondary aliphatic amine to its schiff base or imine,
2. if by using a single isomer of the secondary aliphatic amine , will the degradation or oxidation of the compound recimize it once its reduced back to the secondary aliphatic amine?

.......an interesting project, would appreciate some input, i have tried the benzaldehyde reflux , and reduced it back to the amine with NaBH4 but have noted no change ...that is chirality did not change, also tried the HI boil to destabilize the chiral center but no recimization occurred....solo

1. Depends on the amine. If it is pretty robust toward electrophiles, the amine can be N-halogenated and transformed to the imine via elimination in the presence of a base (for example using DIPEA) or by heating. Otherwise you are left with oxidation to the corresponding ketone and reforming the imine with the proper amine.

2. Assuming that by single isomer, you mean enantiomer, it then depends on where the chiral centre is. If it is on the carbon attached directly to the amine nitrogen then yes, you get a racemic amine after reduction, provided of course you do the reduction in a symmetric reaction environment. If the chiral centre is elsewhere, then such a treatment forms two diastereoisomers in a ratio depending on the asymmetric induction from the existing chirality (usually a near to 1:1 ratio unless the existing chiral centre is close to the imine group).

Refluxing with benzaldehyde will do nothing to a secondary amine except for forming an equilibrium with the hemiaminal (and/or iminium salt if an acid is added and water removed via Dean-Stark, etc.). So I really can't see what intermediate have you been reducing with NaBH4 and how it would have given you back the starting material (reduction of the hemiaminal or iminium salt would have given you the corresponding N-benzylamine).
Refluxing in HI certainly does not racemize an enantiomer having the chiral center at the carbon attached to the amine nitrogen, except if there is a neighbouring group allowing racemization via enolization or some similar mechanism. But if that was the case such harsh conditions would have most probably decomposed the substrate - just heating up the hydrochloride salt would have already been enough for a racemization via enolization if this was viable (one such examples are alpha-aminoketones or "activated" alpha-aminoacids). If there is no such neighbouring group then treatment with acid does nothing. In some cases a nickel catalysts can be used for a one step racemization via dehydrogenation/hydrogenation mechanism but the amine must be robust enough not to succumb to side reactions.




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[*] posted on 15-6-2010 at 12:39


Lambda-Eyde: I have ten or so books on chromatography of several varieties. I will upload them and post a link after finals are over (finish on saturday, expect them sometime early next week).
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[*] posted on 15-6-2010 at 13:12


Nicodem, thanks for the recommendation. I couldn't find that book in references, though. And there was only one for sale on Amazon at 300$.

DJF90: Thanks. I'm looking forward to that.

Am I the only one on the forum who prefers books in paper form? :P
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[*] posted on 15-6-2010 at 23:24


No not at all, theres nothing better than a real book!
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[*] posted on 16-6-2010 at 00:33


Quote: Originally posted by Lambda-Eyde  
...
Am I the only one on the forum who prefers books in paper form? :P


I generally prefer physical books, but more recently have increasingly been switching to digital for simple reference materials.

There's several reasons for this. The first is that I already have around 4 thousand books and journal issues of all sorts, moving gets to be a bit challenging if many flights of stairs are involved.

Then there's the problem of I'm using more references than fit on the desk/table, which electronic media solves. Plus I can copy out the currently useful bits of references, and combine them in a single location as part of the documentation of what I am working on.

Finally:
Quote:

And there was only one for sale on Amazon at 300$.

need I say more?

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[*] posted on 17-6-2010 at 06:00


Quote: Originally posted by Lambda-Eyde  
I made 500 ml of 0,1M (NH<sub>4</sub>;)<sub>2</sub>Fe(SO<sub>4</sub>;)<sub>2</sub> solution today by dissolving 19,61 g of the dry powder to 500 ml with deionized water in a volumetric flask. It is very old, however it had the characteristic blue color of Fe(II) it should have before I dissolved it. The solution, however, was more like a murky yellow. Now I'm seeing a brown precipitate, but not much.
Is this due to atmospheric oxidation or must it have been something in the water that caused it? Mohr's salt is said to be stable against air oxidation because of the acidic ammonium ions present in the solution.

Anyhow, would it be okay if I filtered off the fine precipitate and standardized the solution against 0,02M KMnO<sub>4</sub>? The permanganate is standardized against sodium oxalate, a primary standard.


Here is a picture of the solution, nine days old:




A close-up on the precipitate:

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[*] posted on 17-6-2010 at 14:00


Would a mixture of sodium bisulfite, sodium sulfite, and sodium dithionite suffice to make bisulfite adducts?



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[*] posted on 18-6-2010 at 04:21


Quote: Originally posted by querjek  
Would a mixture of sodium bisulfite, sodium sulfite, and sodium dithionite suffice to make bisulfite adducts?


Bisulfite adducts were a common technique used to elucidate structure of organic molecules prior to H1 NMR. They used many adducts back then because this was the part of the regime which was the only way to determine structure.
Not all aldehydes and ketones form bisulphite adducts easily or simply or stably, some do and you can even use the adduct formation as a crude purification step, in other examples it forms difficulty or not at all.
So in answer to your question, without the ketone/aldehyde your intending to formt the bisulphite adduct of, no-one could reliably answer your question, what is your material?




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[*] posted on 18-6-2010 at 05:38


The substrate is normal benzaldehyde. I've made its adduct before with sodium bisulfite, but I ran out, and found a product whose constituents are the three compounds I've listed above.

I know that benzaldehyde will produce bisulfite adducts, so that's not where my concern arises from. It's that, previously, I've added bisulfite solution to a solvent which I had forgotten to neutralize excess acid in, and that produced a nice lot of SO2. Although all of the compounds I'm asking about are basic salts, I figured I should ask before trying: I'm not sure what the specific basic conditions to form bisulfite adducts are.




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[*] posted on 18-6-2010 at 06:22


first post

How long is concetrated NH3-Solution 25-30%) , conc. HNO3 (65%) and concentrated HCl (37%) durable? I store it at 10 degree celsius in original and dense containers (HCl and HNO3 in glass bottles, NH3 in plastik bottles), and i want to buy stuff for the next 15 years or so...

:=)

[Edited on 18-6-2010 by Addon]
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[*] posted on 18-6-2010 at 14:57


I'm trying to make Pb acetate using excess white table vinegar and 30% H2O2. How do I get it to crystallize? When I boil it down and cool it only becomes a brown sticky mass. Also is it normal that when I diluite it becomes turbid and a white powder precipitates?
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[*] posted on 18-6-2010 at 23:24


Dilution with tap water precipitates lead carbonate, mostly. Use distilled water.

Brown is either an organic or iron impurity. Let it sit for a long time and see if any crystals form; if not, precipitate and start over.

Tim




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[*] posted on 19-6-2010 at 06:26


Quote: Originally posted by Addon  
first post

How long is concetrated NH3-Solution 25-30%) , conc. HNO3 (65%) and concentrated HCl (37%) durable? I store it at 10 degree celsius in original and dense containers (HCl and HNO3 in glass bottles, NH3 in plastik bottles), and i want to buy stuff for the next 15 years or so...

For NH3(aq) and HCl(aq) it depends only on the stability of the containers they are stored in. There is pressure building up during the warmer times in bottles containing concentrated aq. solutions of HCl or NH3, so it is best to make sure the container is robust enough and/or these are stored in cold.
For conc. HNO3 stability it depends also on the temperature and exposure to light. Generally it can last for years with minimal decomposition to NO2 if stored in a glass container, cold and dark, but even if gets saturated with NO2, there are simple means to remove it (UTFSE).
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[*] posted on 19-6-2010 at 10:11


thx a lot!

:=)
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[*] posted on 21-6-2010 at 03:53


How do you make Barium Permanganate from Ba(NO3)2 + KMnO4



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[*] posted on 21-6-2010 at 08:17


Combine solutions and filter the Ba(MnO4)2 precipitate.

Tim




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[*] posted on 21-6-2010 at 08:53


Quote: Originally posted by DJF90  
I'm not entirely sure to be honest, If I were going to try it I would make sure the phenol and zinc dust were mixed homogeneously (as much as possible) and then heat the distillation flask until benzene starts (and finishes) coming over.


Good point - surface area will be a concern for this one. It makes me think of the benzene from sodium benzoate thread where a major improvement was made by adding steel wool to the reaction vessel to increase thermal conductivity. I'm going to try a set up where I use a pressure equalizing addition funnel to feed phenol pre-heated to 165 C slowly onto red hot zinc dust inside a 3-neck flask connected to a distillation column. I think it might help by increasing the amount of phenol in contact with the zinc dust and increase the amount of time it has to react with the zinc.
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[*] posted on 22-6-2010 at 15:15


Quote: Originally posted by 12AX7  
Combine solutions and filter the Ba(MnO4)2 precipitate.

Tim


Thanks. For some crazy reason I thought Ba permanganate was soluble




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[*] posted on 22-6-2010 at 15:41


Salts of univalent oxy-anions are usually soluble, but it appears that Ba(MnO4)2 is an exception.
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