| Pages:
1
2
3
4
5 |
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
"bromine isolation" *chemistry nazi*
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by AJKOER  |
Now, ...over the possible loss of some of the Bromine to Bromate, one may be able to recover the loss by adding more NaBr and an excess of acid
paralleling the reaction with Iodine and Iodate. In particular, per "Synthetic inorganic chemistry: a course of laboratory and classroom study ...",
by Arthur Alphonzo Blanchard, Joseph Warren Phelan, page 232:
"Iodate and iodide ions alone have no action on each other, but with hydrogen ions present a mutual oxidation and reduction of the iodine takes place.
6H+ + 5 I- + IO3- -> 3H2O + 3 I2
No oxidation or reduction of the hydrogen occurs, but the hydrogen ion is used up, which explains why the presence of acid is necessary to make the
reaction take place." where I would assume a similar reaction based on Bromine. Link:
http://books.google.com/books?ei=E1ewT5P1IOaR6gHf9fnWAQ&...
|
For those interested in the chemistry of Bromine, I have found precisely the same reaction referenced for Br2, however, clearly expressed in an
equilibrium form. Source: "Inorganic Chemistry" by Egon Wiberg, A. F. Holleman, Nils Wiberg, page 450, equation (1):
BrO3(-) + 5 Br(-) + 6 H(+) <-----> 3 Br2 + 3 H2O
Link: http://books.google.com/books?id=Mtth5g59dEIC&pg=PA450&a...
Note, the following important comment by the author:
"Clearly, an increase of the proton concentration and decrease of pH causes the rapid protonation equilibrium (2a) to shift to the right; since this
equilibrium precedes the rate determining step, the overall reaction rate increases", which echoes Woelen comments on acid strength.
So theoretically, with the appropriate addition of excess Bromide and acid, any Bromate formation may be reduced with the recovery of Bromine. Quite
interestingly, add 1/5 of the equation above:
1/5 HBrO3(-) + Br(-) + H(+) <-----> 3/5 Br2 + 3/5 H2O
to my prior cited reaction:
NaOCl + HOAc + NaBr ---> NaOAc + NaCl + 1/5 HBrO3 + 2/5 Br2 + 2/5 H2O
produces the amazing simple and intuitive target net ionic equation:
OCl- + 2 Br- + 2 H+ ---> Cl- + Br2 + H2O
Thus, the outstanding question with the hypochlorite method (awaiting verification) is even with steps to reduce dilution (as concentrated HOBr is
more unstable with its decomposition yielding Bromine, perhaps with the help of dry acid salts like Tartaric or Critic acid and also use of conc
hypochlorite) and decreased pH (target range is a pH of about 2 to about 6.4 per previously cited Patent 5516501), and a bromide to hypochlorite molar
ratio of at least two per the target equation above; will distillation then allow the recovery of elemental Bromine.
[Edited on 18-5-2012 by AJKOER]
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
I tried making bromate by 1-cell electrolysis of the bromide, but I don't know how it went. Is this the same as the iodine synthesis of tincture of
iodine + H2O2 + HCl?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
longtermmadscientist
Harmless
Posts: 9
Registered: 12-12-2010
Member Is Offline
Mood: No Mood
|
|
OK, so I added a large excess of conc. HNO3 to solid NaBr dihydrate and the reaction started immediately- a deep red layer formed over the NaBr, and
an even deeper red substance, probably Br2, formed amongst the crystals of NaBr dihydrate. However, the reaction seemed to stall. I think that the
reason for this may be that bromine is denser than the NaBr dehydrate, and the latter forms a protective layer over the bromine. As the NaBr gradually
reacts with the HNO¬3, and forms NaNO3, this effect continues, as NaNO3 is about as soluble as NaBr. Also, even if these compounds dissolve, there
will still be a delay in the mixing of this salt layer and the HN03, and access of the latter to the NaBr. I may repeat the experiment but using a
saturated solution of NaBr.
[Edited on 19-5-2012 by longtermmadscientist]
|
|
|
BromicAcid
International Hazard
Posts: 3247
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Wait, were you trying to make nitrosyl bromide?
I would think that if you're trying to make bromine from NaBr / HNO3 obviously one gets reduced one gets oxidized.
2NaBr + 4HNO3 ----> Br2 + 2NaNO3 + 2NO2 + 2H2O
Then nitrogen dioxide can react with bromide salts:
2NO2 + NaBr ----> NaNO3 + NOBr
Just something to consider.
[Edited on 5/19/2012 by BromicAcid]
|
|
|
longtermmadscientist
Harmless
Posts: 9
Registered: 12-12-2010
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by BromicAcid | Wait, were you trying to make nitrosyl bromide?
I would think that if you're trying to make bromine from NaBr / HNO3 obviously one gets reduced one gets oxidized.
2NaBr + 4HNO3 ----> Br2 + 2NaNO3 + 2NO2 + 2H2O
Then nitrogen dioxide can react with bromide salts:
2NO2 + NaBr ----> NaNO3 + NOBr
Just something to consider.
[Edited on 5/19/2012 by BromicAcid] |
Ah....that’s something I didn’t think of; thanks. I would hope that not much of this would form, as this is something else which could contaminate
my bromine. I may have to use H2SO4, but I’m not enthusiastic about the yield. It seems that methods for producing bromine without involving
distillation may well be unfeasible.
|
|
|
woelen
Super Administrator
Posts: 8014
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
No, you do not need distillation. You can get a yield of e.g. 70% without distillation and the resulting bromine also can be quite pure. But you need
the right chemicals for doing so.
HNO3 + NaBr is not the right combination. I already wrote about that before. The stalling of the reaction is even something which I predicted in one
of my previous posts, due to limited solubility of NaNO3 in conc. HNO3. Your bromine most likely will be heavily contaminated with NOx or ONBr.
H2SO4 is even worse. Yields will not be better than a few percents, and for each mole of Br2 you also get a mole of SO2. This also seems like a bad
contaminant and as soon as some water is present in the system, the water, SO2 and Br2 will react to H2SO4 and HBr.
|
|
|
longtermmadscientist
Harmless
Posts: 9
Registered: 12-12-2010
Member Is Offline
Mood: No Mood
|
|
Right, this is something which is going to take more preparation than I thought. I tried the HNO3 method because this was an oxidant I had to hand and
was curious as to why there was so little mention in the literature of producing bromine from bromides using it. Now I know from both theory and
practice that this approach isn’t effective and why. If and when I manage to get any of the chemicals necessary for the more effective methods,
I’ll be giving them a try, and will report back with my results. Thanks to everyone for their feedback and help.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER |
OCl- + 2 Br- + 2 H+ ---> Cl- + Br2 + H2O
Thus, the outstanding question with the hypochlorite method (awaiting verification) is even with steps to reduce dilution (as concentrated HOBr is
more unstable with its decomposition yielding Bromine, perhaps with the help of dry acid salts like Tartaric or Critic acid and also use of conc
hypochlorite) and decreased pH (target range is a pH of about 2 to about 6.4 per previously cited Patent 5516501), and a bromide to hypochlorite molar
ratio of at least two per the target equation above; will distillation then allow the recovery of elemental Bromine.
|
The prior mentioned Patent 5,516,501 issued on May 14, 1996 ( http://www.patentstorm.us/patents/5516501/description.html ) stated "Processes for preparing a relatively concentrated aqueous solution of about
700-3000 ppm hyprobromous acid are provided. Hypochlorous acid solutions are prepared by either reacting chlorine gas with water or sodium
hypochlorite with an acid. The resulting hypochlorous acid is then reacted with an alkali metal or alkaline earth bromide in order to form the
hypobromous acid. Critical parameters are pH, Br/Cl mole ratio, and chlorine concentration. Under optimum conditions, substantially 100% conversion of
bromide to hypobromous acid can be attained."
The author also notes, "The bromide/chlorine mole ratio is also important. At pH values above about 3.0 it has been observed that as this mole ratio
approaches an optimum value of 1.5 or greater, the yield of hypobromous acid approaches 100%."
A preceding patent (Patent 5795487, Link: http://www.docstoc.com/docs/49815324/Process-To-Manufacture-... ) to the one referenced above noted that:
"a. Mixing an aqueous solution of alkali or alkaline earth metal hypochlorite with a water soluble bromide ion source;
b. Allowing the bromide ion source and the alkali or alkaline earth metal hypochlorite to react to form a 0.5 to 30 percent by weight aqueous solution
of unstabilized alkali or alkaline earth metal hypobromite; "
So one may consider changing the mixing order of the target equation above and react say, concentrated aqueous NaOCl with NaBr and then add a weak
acid to liberate Hypobromous acid (as HOBr is a very weak acid), as a path to Bromine.
Reaction example, with Bleach, NaBr and Tartaric acid:
NaOCl + NaBr ---> NaCl + NaOBr
2 NaOBr+ C4H6O6 ---> Na2C4H4O6 + 2 HOBr
5 HOBr ---> HBrO3 + 2 Br2 + 2 H2O (more so for conc HOBr with mild heating or sunlight and low pH)
With respect to the stability of HOBr and its propensity to form Bromine, a source states:
"Direct sunlight will have a negative impact on the appearance of the product [referring to HOBr] by changing the color from clear to degrees of
yellow or orange. The orange color is bromine (Br2), which is normally not a problem, but if allowed to remain in the sunlight more bromine will
develop, which will have a distinct pungent odor. The solution is still useable in the "yellow-orange‟ state, but these conditions should be
avoided by not allowing the precursor to be subjected to direct sunlight (UV) or excessive continuous heat to allow the temperature of the bulk liquid
to exceed 90° F."
The author also notes: "Hypobromous acid degradation rate accelerates with increasing concentrations. The decay rate for a 200-300 ppm solution of
available bromine would result in a half-life of about 10 days, whereas a 4000 ppm solution may only have a half life decay rate of only a few hours
or less."
And also: "When hydrogen bromide and bleach are added in the correct proportions, the pH of 6.9-7.4 will always be the result, regardless of the
hardness or alkalinity of the feed water. The decay of hypobromous acid always results in either hydrogen bromide or bromine (Br2) formation, or a
mixture of both, because the decay of hypobromous acid is an acid-releasing process. At higher temperatures this decay rate increases proportionally
with the rise in temperature. As the pH drops, the decay of hypobromous acid accelerates and it becomes an autocatalytic process. At concentrations
less than 1,000 ppm this is a slow and relatively harmless side reaction."
So forming more concentrated HOBr solutions in the presence of sunlight or mild heat in a more acidic environment should help liberate Bromine. Link:
http://www.stabilizedbromine.com/pdf/Generation%20%28Blendin...
[Edited on 22-5-2012 by AJKOER]
|
|
|
Zan Divine
Hazard to Others
Posts: 170
Registered: 3-12-2011
Location: New York
Member Is Offline
Mood: Wishing all the worst that life has to offer to that SOB Wayne Lapierre
|
|
| Quote: Originally posted by longtermmadscientist | | Right, this is something which is going to take more preparation than I thought........ If and when I manage to get any of the chemicals necessary for
the more effective methods, I’ll be giving them a try,...... |
One of my clients just ordered 300-400 g bromine. I got the H2SO4 for $10 from a hardware store. The KMnO4 came from e-Bay ($16 for 1 pound
delivered). Leslie Pools sells KBr (4 Lbs for $33). For $59 and one afternoon I'll make way more than 300 - 400 g Br and have lots of starting
materials left over.
Do any of the other methods discussed give this much interhalogen-free bromine for this low a cost with so simple a procedure? 6 to 7 g Br for $1. I
tend to doubt it. H2SO4 is the second cheapest strong acid. KMnO4 is cheaper than appropriately concentrated H2O2 by a wide margin (I don't even
consider any chlorine-containing oxidants as I don't want any BrCl).
Since the bromine spontaneously distills from the mixture due to the heat of reaction, collection of the crude product is as simple as adding a
condenser column or still head.
So, while lots of routes are available, I don't see one better than this if your goal is to obtain Br2 and the chemistry is just a tool to get there.
Considering how simple the chemistry is, why complicate it? I'd think the things you can do with Br2 would be much more interesting than making it.
When you're young, time seems infinite, but still, why waste it?
[Edited on 29-5-2012 by Zan Divine]
He who makes a beast of himself gets rid of some of the pain of being a man. --HST
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
| Quote: Originally posted by Zan Divine | All factors considered, I favor a warm, stirred mixture of NaBr & KMnO4, treated dropwise with conc. H2SO4, which will quickly warm to a temp
above the bp of Br2. The vapors are condensed, washed, dried over NaBr & distilled.
There are no interhalogens as with Cl2. This is operationally simple and it is considerably cheaper than H2O2, which is important for larger batches.
Look at Wiki under bromine for the stoichiometry.
[Edited on 14-5-2012 by Zan Divine] |
How well does the reaction proceed if the vessel is being cooled by an ice bath to prevent the fuming? I don't like the idea of immediate bromine
vapors. It seems better to me if it's left to accumulate on the bottom, from where it can be easily drawn out, dried and distilled with a small amount
of conc. sulphuric acid.
Also, wouldn't conc. hydrogen peroxide be cheaper, considering KMnO4 is hard to find in a non "analytical grade purity" state? Peroxide is easy to
find concentrated, but not extremely pure, though sufficiently pure for this application.
[Edited on 29-5-2012 by Endimion17]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Peroxide is cheaper. KMnO4 is great but 'overkill' for many reasons here.
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
Bleach works too, have a solution of bromine in water sitting in my garage as we speak. Making a distillation apparatus, though, is quite another
thing.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
^^no, it sucks. We're talking about yields here.
Bromine water is for testing hydrocarbon saturation. Not for making bromine.
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
Some of us don't have access to KMnO4 or H2SO4, and thus we make do.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
Zan Divine
Hazard to Others
Posts: 170
Registered: 3-12-2011
Location: New York
Member Is Offline
Mood: Wishing all the worst that life has to offer to that SOB Wayne Lapierre
|
|
I guess it's no surprise that cost is location dependent. Around here, New York, concentrated H2O2 is expensive and hard to get. Maybe because of
acetone peroxide...
None of the British suppliers will ship here anymore.
Endimion17, as a lifelong synthetic chemist, I guess I tend to go straight to the heart of a synthesis with an eye toward minimizing manipulations and
optimizing simplicity. This method requires no cooling bath, no heating bath, and absolutely no manipulations left to do after the reaction is done
except drying & distilling. To me, there is synthetic elegance in making the reaction provide the driving force for the isolation and having the
crude reaction product look like the below. You seem a bit leary about handling boiling bromine. If so, it's probably better if you don't because of
toxicity concerns. If you want a cool reaction, peroxide is superior because of the greater ease in separating layers. I'm guessing you're in Britain
so that's easy.
[Edited on 30-5-2012 by Zan Divine]
He who makes a beast of himself gets rid of some of the pain of being a man. --HST
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
I've handled chlorine before without dying, so I can deal with bromine. Should I put anything into the distillation apparatus to deal with any salt
that might have come over with the bromine?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
| Quote: Originally posted by Zan Divine | I guess it's no surprise that cost is location dependent. Around here, New York, concentrated H2O2 is expensive and hard to get. Maybe because of
acetone peroxide...
None of the British suppliers will ship here anymore. |
That's a shame. Things like that piss me off.
| Quote: | | Endimion17, as a lifelong synthetic chemist, I guess I tend to go straight to the heart of a synthesis with an eye toward minimizing manipulations and
optimizing simplicity. This method requires no cooling bath, no heating bath, and absolutely no manipulations left to do after the reaction is done
except drying & distilling. To me, there is synthetic elegance in making the reaction provide the driving force for the isolation and having the
crude reaction product look like the below. You seem a bit leary about handling boiling bromine. If so, it's probably better if you don't because of
toxicity concerns. If you want a cool reaction, peroxide is superior because of the greater ease in separating layers. I'm guessing you're in Britain
so that's easy. |
I do appreciate the simplicity during the laboratory work. That's why it seems to me it's better to mix the reagents in an ice bath and just leave
them alone in a weakly stoppered flask. Bromine essentially stays inside. It takes more than one hour before it's sufficiently progressed. I've tried
it, and the only problem I see is that not many of us have any stirring equipment. Because it's a saturated bromide solution, and effectivelly a kind
of piranha solution, all cooled down to ~0 °C, the reaction should be agitated because otherwise, layers occur, that are mixed by diffusion and
convection by small oxygen bubbles, which is too slow.
However, decent mixers can be made using an electromotor and a shaped glass rod attached to it. Old school, but effective, and there aren't any teflon
stirbars soiled with bromine.
Boiling bromine doesn't pose a problem to me because of the toxicity as I have experience with it. Lowering the yield does. The crude product has to
be distilled, as you know. If one wants a reasonably pure bromine, doing it your way is essentially distilling it twice. Given the probable scenario
that one doesn't have a semimicro destillation kit with a tight and tiny cooler, it means a significantly lowered yield.
Doing it slowly, in an ice bath, means the losses are with dissolving in the reaction mixture only, which are reasonably smaller.
It's all about the yields, because bromides are hard to find where I live (Croatia). AFAIK, there aren't any crude and cheap products, only analytical
ones.
|
|
|
Zan Divine
Hazard to Others
Posts: 170
Registered: 3-12-2011
Location: New York
Member Is Offline
Mood: Wishing all the worst that life has to offer to that SOB Wayne Lapierre
|
|
| Quote: Originally posted by elementcollector1 | | I've handled chlorine before without dying, so I can deal with bromine. Should I put anything into the distillation apparatus to deal with any salt
that might have come over with the bromine? |
Actually, bromine is somewhat more toxic/sneaky if memory serves, but your point is well taken, they are similar.
No salt should be entrained in the bromine vapors to speak of. It wouldn't matter anyway. The next step is to stir the crude bromine with some finely
ground and thoroughly dried NaBr to dry it. Conc. H2SO4 can also be used. It just depends if you prefer a sep. funnel or a filter to remove your
drying agent.
I use a funnel with a glass wool plug, outside.
There is a patent that describes holding Br2 just below boiling for some time to dry it, but I've never tried it.
If a tiny bit of salt gets through, no problem,it stays in the pot during distillation.
If you choose this method you can just start with HOT tap water (I can hear the purists shrieks but it doesn't matter one iota). Similarly, KMnO4 can
be any purity, I've used 80 - 98% and the results are the same. The H2SO4 is just drain cleaner. Work with as concentrated solutions as you can get to
maximize yields.
All that being said, everybody optimizes their syntheses to suit their equipment and you've made good points, Endimion17. Now that I know all the
factors in play, you're probably ahead of the game with H2O2
[Edited on 30-5-2012 by Zan Divine]
He who makes a beast of himself gets rid of some of the pain of being a man. --HST
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
I choose NaBr, as I have a steady supply. 
How do I thoroughly dry this, a microwave? (Did it with MnO2 sludge to great effect)
Now, to find some glass tubing. I have no idea where to start looking. 
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
barley81
Hazard to Others
Posts: 481
Registered: 9-5-2011
Member Is Offline
Mood: No Mood
|
|
Elemental scientific has some. You could look for a neon sign shop or a glassblower in your area. They will probably have it. You could also get some
on eBay..
|
|
|
Zan Divine
Hazard to Others
Posts: 170
Registered: 3-12-2011
Location: New York
Member Is Offline
Mood: Wishing all the worst that life has to offer to that SOB Wayne Lapierre
|
|
ec1, Yes, the microwave solution is perfect. Just heat the salt in 30 sec to 1 minute intervals until it no longer gets real hot. Then cook the hell
out of it for several minutes (It's best not to leave it alone. Some microwave ovens have been known to react poorly to being driven with no real load
to soak up the microwaves. They'll arc. If you turn it off right away, no harm done).
If you have a balance you can simply weigh it between heating periods until no more weight is lost. This is called "drying to constant weight".
E-bay would seem to be the best source of glassware of all sorts for many people. In the US, there aren't many ways more direct. You'll save over
buying from distributors but for most glassware don't expect "a steal". The common things like round bottom flasks and condensers and especially
heating mantles are still moderately costly. If you can use the smaller scale 14/35 (or 14/20) glassware instead of the old standby 24/40 you'll save
money. E-bay has glass tubing (Stay away from the soft glass, it's a relic. Go for borosilicate or hard glass). I don't know how well e-Bay works for
other countries.
[Edited on 31-5-2012 by Zan Divine]
He who makes a beast of himself gets rid of some of the pain of being a man. --HST
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
I *have* a balance, but it kinda sucks. Only accurate to the nearest 50-gram interval. XD
As for the round-bottom flask, which is useful for distillation and making retorts, I was wondering if a lightbulb could withstand the heat. It's so
thin I'd expect it to shatter upon touch after heating, but maybe...
How do you "shape" the hard glass? I know you can make holes by passing through a heated piece of copper pipe or wire, but how do you bend it without
'kinking' it?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
Air pressure, usually from the lungs, hence the term glass-"blowing".
There are a number of old glassblowing manuals digitized. Search for them. I don't recall the titles offhand. I've read a number of them; they're all
pretty much equally good for the basics.
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
| Quote: Originally posted by elementcollector1 | | How do you "shape" the hard glass? I know you can make holes by passing through a heated piece of copper pipe or wire, but how do you bend it without
'kinking' it? |
You can't pass a heated wire through. No way. The heat dissipation is too great for the glass to reach sufficient temperature.
I recommend downloading one of several available flameworking (not glassblowing, that's just "making bottles") manuals on the Web.
If you want to obtain some skills, buy a handful of glass tubing of various sizes, one usual propane-butane blowtorch and one fine blowtorch with a hot, narrow flame like the ones used in kitchens.
When you get everything, start following the manual and soon you'll be able to join tubes, make small flasks, make holes through flasks, join tubes
with those flasks, etc.
But don't fool yourself - you'll never be able to recreate a real flameworker's job, simply because you don't have the neccessary tools.
It's might seem like a futile practice, but believe me, elementary flameworking is essential for a home chemist as it can save a lot of money. I've
saved quite a lot just by buying glass tubing and making my own stuff, some of them disposable. I didn't want to spend money on buying something that
will have to endure lots of heating.
Forget the light bulbs. They're for entry level home chemist, for boiling water. That's glass with high percentage of sodium. It's soft, thin and
fragile, with a too low softening point.
Just buy a handful of glass tubing.
|
|
|
| Pages:
1
2
3
4
5 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
