Every now and then during a simple distillation I'll have a substance boil but refuse to distill over. It's as if it will boil almost endlessly before
even a drop settles in the distillate flask. My first reaction was maybe since I was using a Vigreux that the vapors were condensing back into the
reaction flask as reflux so I put aluminum foil around the outside of the column. That only proved slightly effective. Could there be any other
problems that would cause this to happen? Thanks in advance.ScienceGeek - 8-3-2009 at 03:59
It's the Vigreux column!
Try taking the whole thing away from the distilling apparatus!DJF90 - 8-3-2009 at 06:31
It depends if poisoninthestain wants to fractionate whatever he is distilling. If not, then remove the vigreux from the setup, as you will only make
the job longer and more energy consuming (as you have seen by the sounds of it )
[Edited on 8-3-2009 by DJF90]Lambda-Eyde - 8-3-2009 at 07:27
Well, he says he's doing a simple distillation, right? I see no use
for a vigreux if he means simple distillation as in without fractionating column!poisoninthestain - 8-3-2009 at 15:32
the reason i say "simple" distillation is because i only have a vigreux column and nothing else. That's why I have to use it. Otherwise I'd have to
connect the boiling flask to a 105 bent vac adapter, feed that into the liebig, and then put the vigreux on the end to drip into the distillate flask.
Make sense?kclo4 - 8-3-2009 at 15:37
I've heard of people aiming heat guns, and wrapping house insulation around glassware to aid the distillation process.
Good luck! DJF90 - 9-3-2009 at 05:15
You mean your stillhead and vigreux column is one piece? If this is the case then I suggest you get a plain stillhead.Saerynide - 9-3-2009 at 08:20
Quote:
Originally posted by kclo4
I've heard of people aiming heat guns, and wrapping house insulation around glassware to aid the distillation process.
Heat guns... Hmmm... What an excellent idea!
A hair dryer should work too HydroCarbon - 9-3-2009 at 08:54
You could always use a different, less efficient, column if you have one or can obtain one. That way it would take less time, but you still would
have fractionation. Vigruex columns are extremely efficient iirc.chemrox - 9-3-2009 at 10:10
It would help a lot to know what you're working with. You could be removing a supervolatile component that will not condense in your setup. Anyway,
you need a simple still. For a columhn I would get a Hempel and pack it with glass rings or a Snyder (for simplicity and charm).HydroCarbon - 16-3-2009 at 18:38
Another question to ask is how much volume are you distilling in what size of apparatus. Distilling too small a volume of liquid in a large apparatus
will not work.
Originally posted by poisoninthestain
the reason i say "simple" distillation is because i only have a vigreux column and nothing else.
The term "simple distillation" means one with no fractionating column, no matter how you may use it.
If you have a one-piece still head plus column, then it is likely the column is fairly short. Vigreux and Synder are low holdup, in 24/40 typically
less than a ml IIRC, so you may be OK with your rig. Hemple and other packed columns usually are higher holdup.
If you want distill the highest boiling component of the mix you'll need a chaser, an inert liquid boiling enough higher than the high boiling
component to be easy to separate. I typically use kerosene, after distilling off any part boiling too low. Add some to the still pot, heat high
enough to boil up all the desired component; watch for the temperature rise as the chaser starts to reach the head.chemrox - 16-3-2009 at 22:44
Ahh yes that takes me back a bit. A couple of years ago I distilled some kerosene until I had the lower boiling materials out of the mix so I could
use it as a chaser. Anyway if the volume is small relative to your flask you might be cooling the vapor before it reaches the ditillation head. You
should wrap everything in Al foil. A heat gun could be a real help too ...Sauron - 16-3-2009 at 23:22
Lambda is right.
A SIMPLE distillation is one with NO column interposed between flask and condenser, just a suitable adapter for downward distillation.
Such distillations are only useful when the two components of the mixture have about 100 C difference in boiling point.
In a FRACTIONAL distillation, a column is used, usually with a still head and capability to control reflux ratio, and a fraction cutter on the
receiving end such as a cow.
In any factional distillation, the column must be well lagged and a long time allowed for the column to equilibrate. During that time the reflex ratio
is set to 100%, that is, reflux and no takeoff.
You will know when column is equilibrated when the tempeature at top of column is that of the lowest boiling component of the mixture. The temperature
drops no further. At that point you can slowly start reducing the RR so that slow takeoff occurs, but take care not to disturb the equilibrium by
taking off too fast or a poor seperation will result.
Monitor the temperature at still head to follow progress and change fractions at collector when temperature rises at still head.
This is a slow tedious procedure and cannot be rushed. There is no such thing as a good fast fractionation.bbartlog - 23-1-2010 at 21:14
I have a different question about distillation: is it normal, when distilling various fractions one after another, to see a temporary *drop* in the
head temperature if there is a particularly large gap in boiling points from one fraction to the next? Just to provide some background: I was
fractionally distilling something of a mess I had made while doing an experimental synthesis of 2-chloropropane. While distilling various fractions I
saw what I assume is a normal stepwise progress of temperature over time, something like
33-35C step 1 36g 2-chloropropane comes over, then temp rises fairly quickly to
46-49C step 2 13g 1-chloropropane(?, best guess) comes over, then another rise to
55-57C step 3 11g of some unknown compound comes over, then another rise to
67-69C step 4 15g of some isomer of hexane(?, best guess) come over,
...and then the temperature in the head drops back down to 59C, despite no particular changes to the setup (i.e., heat input remained constant). My
first intuition was that this shouldn't happen, but after thinking about it more I thought that perhaps once one heat transferring fluid has been
driven into the receiver, and the next one is not yet boiling, there can be a situation where heat transfer to the head slows enough for the
temperature to dip. In this case the next step would probably have been some isopropyl alcohol and water mix coming over somewhere above 80C.
So, anyway... dips in temperature at the still head when fractionating? Normal?
[Edited on 24-1-2010 by bbartlog]smuv - 24-1-2010 at 08:49
Yes this is normal, as you assumed, thermometer cools as no/little vapor is reaching the still head.
This being said, if you made the chloropropane from ZnCl2/HCl/Isopropanol, your fractions look a little weird. Did you use a fractionating column, or
are these temps from a simple distillation?
[Edited on 1-24-2010 by smuv]bbartlog - 24-1-2010 at 09:11
The fractions are weird because I didn't use ZnCl2 (and as a result didn't make much 2-chloropropane). The run in question was done using
CuCl/HCl/Isopropanol, just to see what would happen. Short answer: not much (starting mass of IPA was 180g, so you can see that the yield of anything
is puny), though the production of what looks like a small amount some hexane isomer via dimerisation is interesting.
I'm somewhat skeptical of the 'catalytic' role of ZnCl2 in this synthesis, in part because the prep in Vogel uses so much of it, and because the
procedure described there adds anhydrous ZnCl2 to a mix that ends up containing a bunch of water. I have to wonder whether it doesn't function mainly
to withdraw water from the system, in which case other compounds that do so would also work, albeit less well if they are not as acidic. I did a trial
on a much smaller scale with IPA/HCl/CaSO4 which seemed to support this idea, but there are obvious mechanical problems with using CaSO4 as it doesn't
form a hydrate melt and will turn into a solid block in your flask if you add too much of it. I plan several more runs with other variants but since
they take a couple of days each it will be a while before I have a set of results to post...
(edit) - oh, and yes I have a fractionating column - I use a 300mm Hempel column packed about 2/3 full with ~8mm lengths of glass tubing I cut.
[Edited on 24-1-2010 by bbartlog]smuv - 24-1-2010 at 20:46
That 'hexane isomer' is probably isopropyl ether (though I am a little surprised by its formation), the middle fractions could be azeotropes between
isopropyl ether/water/isopropanol/2-chloropropane, based on these structures a few of these are good candidates for forming azeotropes.
The ZnCl2 is not really catalytic, while it acts as a simple Lewis acid it also acts as a dehydrating agent and keeps the concentration of Cl- high.
I also suspect that you need a large excess of ZnCl2 because if you add ZnCl2 to conc. HCl most of it probably hydrates to form Zn+2 and Cl-. This
hydrated ZnCl2 is a much weeker lewis acid. Probably if you add enough ZnCl2, the high zinc and chloride concentration allows some covalent ZnCl2 to
exist.
Actually your results are pretty surprising, I tried a prep of isopropyl chloride from sulfuric acid/hydrochloric acid/isopropanol and only recovered
propylene. The same system with ethanol gave some chloroethane but mostly ether. You should more rigorously purify your isopropyl chloride to make sure this is what you have isolated. I would
recommend washing with dilute base, drying over CaCl2 and a fractional distillation. Burn a small amount of the stuff, it should burn with the
evolution of HCl, as seen by fumes of HCl hydrating in air. If the vapor is ignited (away from anything that will color the flame) the compound
should burn with a blue/green flame.
I don't think CuSO4 would appreciably help this reaction, if you can obtain anhydrous ferric chloride, you might be able to use that in place of
ZnCl2. Check ebay for a supplier. bbartlog - 24-1-2010 at 22:08
I already washed with 0.5M Na2CO3 (following the weighing, so I have less mass now) and have for now stored the first three fractions in one bottle
and the last fraction in another. Putting them back together may be stupid but I figure that if they're really separate compounds I can easily
redistill them to separate them. I may try the additional tests you suggest for the isopropyl chloride though to be honest I don't think there are any
other candidate compounds that both have the right boiling point and could plausibly be produced by this system.
Thanks for the comments on the role of ZnCl2. I figured it was mostly just a convenient combination of mid-strength acid and dehydrating agent in one.
The mini-trial with CaSO4 (not CuSO4) was just to see what mere dehydration without additional acidity would do.
I see what you mean about the highest-boiling fraction I mentioned possibly being the ether. I think it's worth mentioning that all of this involved
two hours of reflux followed by four hours of painfully slow distillation, so there was quite a lot of time for reactions to take place. Anyway I have
no particular notion of what mechanism would have produced ether here, there is nothing to remove water and thereby favor its production, though I
imagine that CuCl can do various strange things by donating an electron.
I could obtain anhydrous ferric chloride, but obtaining anhydrous anything for this reaction seems pointless when we consider that the addition of our
HCl solution is just going to add a bunch of water back into the system. And I already have a bunch of ZnCl2 in the form of an 80% solution. If I want
to achieve the proportion of water present in Vogel's prep, I can mix IPA and HCl and then partially dehydrate it before adding it to the ZnCl2; no
need to render the ZnCl2 anhydrous by complex means.smuv - 25-1-2010 at 21:38
The mechanism producing ether would be the same mechanism as say H2SO4 producing ether, you protonate the alcohol and that is attacked by another
alcohol, forming the etherate, losses the proton gives the ether. The loss of water drives the equilibirum, but still, some always forms. I am
surprised however it did not simply dehydrate to propylene.
If you mix IPA with HCL and 'dehydrate' it you will simply form a bunch of propylene mixed with HCL(g).bbartlog - 26-1-2010 at 08:48
I'm talking about removing some of the excess water that comes with 31% HCl, not trying to aggressively strip all possible H2O from the system.
Anyway, it does sound like I might want to set things up so that I can determine whether any of my alcohol is escaping as propylene gas. I would have
thought that propylene in HCl would just undergo hydrohalogenation as it was produced; in fact depending on conditions, forming a bunch of propylene
mixed with HCl might be a fine intermediate step towards the final isopropyl chloride.
The small trial I did with CaSO4 was primarily intended to see whether the IPA would dissolve HCl reasonably well, or whether I'd be producing a lot
of HCl vapors. Obviously if I take aqueous HCl and bind up most or all of the water, I'm going to drop the bp of the HCl/water mixture quite low and
could have some really unpleasant gas production to deal with. I figured that the IPA, with its -OH group, would probably hold the HCl well enough,
but better to find out with 30ml instead of 500.
To be honest though I'm still not seeing *why* dehydration would be important. Given that the reaction does proceed even with a fair bit of water
present, it doesn't seem like the presence of water interferes per se, which would make it appear that the removal of water is just intended to drive
a reaction equilibrium. But since 2-chloropropane is escaping as a gas under the conditions I maintain, I shouldn't need to drive the equilibrium in
other ways; it's the rate I would be concerned with. Thinking about it, maybe it's the equilibrium concentration of some intermediate specie that I'm
driving, and this determines the rate... but it can't be propene as that would also be gassing off. If C3H8O gets protonated, I guess the H2O+ leaving
group that is formed to takes an electron with it (leaves as plain old H2O) and that leaves a secondary carbocation C3H7+. You know, I think that the
problem with dilute sulfuric acid in this context is that HSO4- may not form a stable compound with the carbocation, whereas Cl- will. Anyway, if the
concentration of C3H7+ determines the reaction rate then I think I see why I need something that will steal H2O.
)
