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Molecular Manipulations
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Certainly
some, but for most purposes it's not a big deal. Run a magnet over the shaving to be sure.
xfusion, I've heated potassium nitrate above that temp before. No bang to speak of (hard to speak with one's face covered in molten oxdizers and
glass shards [JK]). In large amounts and/or confined, I'd stay away.
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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Insanus
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delta H/G for absorption of water by concentrated sulfuric
Does anyone know where to find quantitative data for enthalpy/free energy changes from absorption of water by sulfuric acid? This would be useful for
predicting reactions such as
CH3CHOHCH3 <----> CH2=CHCH3 + H2O
in acidic media.
Edit: NIST has reaction energetics for both sulfuric acid and water. Calculations would need to take into account concentration and the exact ionic
species involved, but the data is there.
[Edited on 6/1/2015 by Insanus]
[Edited on 6/1/2015 by Insanus]
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Funkerman23
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Does anyone know what diameter capillary tube I need to use for a theile tube/ non mel-temp machine melting point determination? I don't have a
mel-temp machine so it has to be the older fashioned sample tied to a thermometer in a theile tube/ oil bath route..I'm stumped and I don't know if
the diameter of the capillary tubing is important to the accuracy or not. I really appreciate any advice.
" the Modern Chemist is inundated with literature"-Unknown
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j_sum1
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Silica gel from cat litter
There are brands of cat litter claiming to be 100% silica gel (according to MSDS). However, not all of the crystals are clear -- there is perhaps 5%
blue crystals in the mix. The label says that it is scented for some of the brands.
Does anyone have any experience with cat litter as a source for silica gel? It seems like it might be a useful thing to have on hand and rather cheap
and simple to obtain.
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diddi
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ok as a spill kit perhaps...
Beginning construction of periodic table display
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Crowfjord
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@Funkerman23: diameter shouldn't matter much, as long as it is fairly thin. I was unaware that melting point capillaries even came in different
diameters. I think the ones I use (also with a Thiel tube and thermometer) have an inner diameter of 1.0 mm.
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Zombie
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Quote: Originally posted by j_sum1  | Silica gel from cat litter
There are brands of cat litter claiming to be 100% silica gel (according to MSDS). However, not all of the crystals are clear -- there is perhaps 5%
blue crystals in the mix. The label says that it is scented for some of the brands.
Does anyone have any experience with cat litter as a source for silica gel? It seems like it might be a useful thing to have on hand and rather cheap
and simple to obtain. |
Ali baba... $600.00 a ton.
The blue crystals you see may be the color indicator. WallMart sells it as a flower desiccant, and it turns blue as it absorbs moisture. It's about 5
bucks a pound there.
I never saw it in cat litter.
They tried to have me "put to sleep" so I came back to return the favor.
Zom.
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PHILOU Zrealone
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But wel with chemophages  ...
Just tell chemophobes that:
Proteins are polyaminoacids, polypeptides very similar to polyamide plastic...they wont eat anything but starch...
--> Starch is a polymer of glucose similar to cellophane and cellulose.
--> They die starving and Darwin wins 
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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The Volatile Chemist
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When chlorinating solutions of, say, dyes, how necessary is it to absorb excess chlorine? I don't have a fume hood.
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Loptr
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Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it
decomposing?
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blogfast25
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| Quote: Originally posted by Loptr | | Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it
decomposing? |
No. The dihydrate has an MP of only 41 C (Wiki). On melting it becomes a saturated solution of KF in water.
Heavily dissociated into K<sup>+</sup> and F<sup>-</sup>, the fluoride ions reacts with water, because HF is a weak acid
(pK<sub>a</sub> = 3.17, Wiki):
F<sup>-</sup>(aq) + H<sub>2</sub>O(l) < === > HF(aq) + OH<sup>-</sup>(aq). Water soluble fluorides do indeed
yield alkaline solutions.
Since as HF is volatile, on heating it will partially 'boil off'. You end up with KOH, or at least KF heavily contaminated with KOH.
[Edited on 11-6-2015 by blogfast25]
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Loptr
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by Loptr | | Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it
decomposing? |
No. The dihydrate has an MP of only 41 C (Wiki). On melting it becomes a saturated solution of KF in water.
Heavily dissociated into K<sup>+</sup> and F<sup>-</sup>, the fluoride ions reacts with water, because HF is a weak acid
(pK<sub>a</sub> = 3.17, Wiki):
F<sup>-</sup>(aq) + H<sub>2</sub>O(l) < === > HF(aq) + OH<sup>-</sup>(aq). Water soluble fluorides do indeed
yield alkaline solutions.
Since as HF is volatile, on heating it will partially 'boil off'. You end up with KOH, or at least KF heavily contaminated with KOH.
[Edited on 11-6-2015 by blogfast25] |
Yeah, that's what I was afraid of, and I don't feel like working with HF at this time.
Thanks blogfast!
[Edited on 11-6-2015 by Loptr]
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greenlight
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Quick question, just concentrated some 70% Nitric acid via distilling with 98% Sulphuric acid and continued until no more nitric came over. The
sulphuric acid obviously has a large quantity of water in it now, is it fine to boil the water back off on heat so its reconcentrated for re-use
again.
[Edited on 21-6-2015 by greenlight]
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The Volatile Chemist
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At the very least, do it in a fume cub., as the dense vapours of the sulfuric acid will form on heating, and are rather unpleasent.
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learningChem
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closely related question :
How long (approx.) should it take for ~94% sulfuric acid to reach 98%? I think I tried once, and yes got nasty clouds of acid, but the concentration
didn't change much. Seems like a lot of acid evaporates, not just the water?
[Edited on 26-6-2015 by learningChem]
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greenlight
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Try what I have been doing by just heating at 110-120.C in a beaker or other glass container on a hotplate instead of boiling the acid as well by
heating strongly. This will boil the water which will evaporate off but leave the acid so no acid clouds.
You can put a watch glass or beaker bottom or anything that has a glass surface over the acid heating container to tell when all the water is gone (I
use a glass saucepan with glass lid). If the glass surface becomes fogged up with condensation, there is still water in the acid. It will also tend
to lightly fume white fumes when all the acid is nearly gone but these fumes will not fog up the glass.
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learningChem
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Thanks greenlight! I'll give your method a try. I was under the impression that, as the acid became more concentrated, more energy was needed to get
rid of the remaining water, but experiment should clear this up =)
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greenlight
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I have used this method many times usually starting with clear 60% Sulphuric acid and afterwards it is much darker yellow and thicker and will char
paper towel black on contact.
Just periodically check for water condensation and when none is visible let it heat for an additional 10 minutes to ensure good concentration
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lenner
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Whats the best way to produce AlBr3 in a solvent? Mixing Al with bromine will form AlBr3 but the reaction is very exhotermic. In a solvent water or
oxidation could be avoided as much as possible. I have doubts about what type of solvent would be adequate and what type of precaution should I take
since the reaction is so exhothermic (which is what keeps me from just trying it).
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Oscilllator
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| Quote: Originally posted by lenner | | Whats the best way to produce AlBr3 in a solvent? Mixing Al with bromine will form AlBr3 but the reaction is very exhotermic. In a solvent water or
oxidation could be avoided as much as possible. I have doubts about what type of solvent would be adequate and what type of precaution should I take
since the reaction is so exhothermic (which is what keeps me from just trying it). |
I think Dichloromethane would be ideal here as it is non-flammable, bromine is miscible in it and does not react with bromine. Just use an ice bath
and/or a reflux condensor to catch the inevitable DCM fumes.
Provided the reaction is spontaneous (It might not be) this should be a very safe way to do it, as the bromine will always be quite dilute. The only
issue you might have is adding a whole bunch of bromine at the start and then having the reaction take off and boiling away all your DCM.
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xfusion44
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Can mineral spirits (white spirit) go bad? I have (probably) pretty old can of it and it should be colourless, but mine seems to be a little bit
brown/orange (maybe there's some rust in can, which causes that?) and it's also imposible to ignite it, but it should burn... I'm confused.
Thanks for answer
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xfusion44
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| Quote: Originally posted by Molecular Manipulations | Certainly
some, but for most purposes it's not a big deal. Run a magnet over the shaving to be sure.
xfusion, I've heated potassium nitrate above that temp before. No bang to speak of (hard to speak with one's face covered in molten oxdizers and
glass shards [JK]). In large amounts and/or confined, I'd stay away. |
Thanks 
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byko3y
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lenner, see US patent 2975214 for preparation of AlBr3 in different solvents.
learningChem, concentrated sulfuric absorbs water from air as crazy. Some moisture guard is a must, the highest concentration I've obtained without
using the protection is something like 92% in a receiver and 96% in a source pot.
greenlight, 98% sulfuric acid can be dehydrated above 300°C, 90% needs more than 270°C, 80% needs at least 210°C.
You need to make at least 50% sulfuric acid content in a resulting mixture (120°C dehydration temperature for H2SO4-H2O system) to leave no way for
the nitric acid to come dilluted from the mixture. Of course, I'm supposing you have azeotropic or higher initial concentration of nitric acid.
http://www.qvf.com/qvf-process-systems/mineral-acids/concent...
[Edited on 29-6-2015 by byko3y]
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learningChem
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Thanks byko3y.
So I did a quick test. Heated a few mls of ~94% sulfuric acid in a small beaker. At ~120C white fumes can already be seen, but I don't think there's
any water in them. I tried to condense the vapor on a piece of glass and didn't see any kind of droplets. I did see condensation/refluxing on the
inside walls of the beaker, a few milimeters above the acid surface, but I suppose that's just acid (though it's a bit strange that the drops 'climb'
the wall at that low temperature?).
Anyway, just as I originally thought, removing water from already concentrated acid does require a fair amount of heat/time? (apart from some sort of
moisture guard)
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The Volatile Chemist
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| Quote: Originally posted by xfusion44 | Can mineral spirits (white spirit) go bad? I have (probably) pretty old can of it and it should be colourless, but mine seems to be a little bit
brown/orange (maybe there's some rust in can, which causes that?) and it's also imposible to ignite it, but it should burn... I'm confused.
Thanks for answer |
Red-orange stuff is probably rust, and it may not burn because it has collected a lot of water from the air over the years. Try distilling it and see
how much water's in it. The distillate should definitely a-flame.
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