Mixell
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Chlorine Solution
I've been thinking about making a chlorine solution, for short terms reaction and may be even for storage (in ground glass vessels kept at -10C)).
The problem is, I need a good and inert solvent. All I have at the moment is ChCl3 and C6H14, but I see those being chlorinated in the long term,
especially the latter one (the chlorine production will be carried outside, exposed to the sun light).
I do not have access to CS2, and I recall that CCl4 has not been available at my local chemical supply company.
Are there any other inert solvents that can be used for making chlorine solutions? Preferably something not too toxic/expensive/flammable.
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elementcollector1
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Water. If you bubble chlorine through it without the presence of UV rays, it will simply dissolve in the water.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Nicodem
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Water is not inert to chlorine. It reacts with chlorine in its disproportionation equlibrium to give chlorous acid (HClO) and HCl (mostly
dissociated). HClO can further dissociate to ClO2, ClO2-, and HCl (see DOI: 10.1021/ic50060a013 and DOI:
10.1021/ic00043a011). These are all reversible equlibriums, but the photodecomposition of aqueous chlorine, which gives oxygen and HCl as products, is
irreversible (see DOI: 10.1039/CT9252700822).
Furthermore, water is a nucleophilic solvent and might not be suitable for whatever the intended use is. Water competes as a nucleophile in the
chlorinations of alkenes and alkynes to give chlorohydrines and chloroketones correspondingly. In fact, these are commonly the main products when
chlorination is performed in the presence of water.
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woelen
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Forget about C6H14 as solvent for chlorine, it reacts quickly, especially when under light conditions. CHCl3 is better, but it also reacts, albeit
slowly. You can use it as a solvent for Cl2 if you use the liquid a short time after preparation, but keeping it for more than a few days will not
work.
Only fully chlorinated alkanes are suitable as long term storage solvents and the only liquid one I know is CCl4. C2Cl6 can be obtained much more
easily than CCl4 (it can be purchased at pyrotechnics suppliers), but this is a solid at room temperature with a low melting point. It might be that
it liquefies when sufficient chlorine is dissolved in it, but that is something you should try.
Certain freons or similar compounds also may be suitable, mixed florinated/chlorinated hydrocarbons, but they may be hard to obtain, probably even
more so than CCl4. Just check at your chemical supplier.
[Edited on 29-5-12 by woelen]
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weiming1998
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I just had an idea, what if you dissolve Cl2 in CHCl3 until saturated, then expose to the sun in a reasonably-sealed vessel? The CHCl3 will chlorinate
to CCl4, and then you can use this home-made CCl4 to use as a solvent. Of course, this is just speculation, and this might not work in real life (the
speed of chlorination).
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woelen
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In practice this can work, but you need to make quite some Cl2 before you have chlorinated all CHCl3 and you also have to be very patient. At an
industrial scale such a process can (is?) interesting, but for the home chemist it probably is too tedious. Just have a look at the short calculations
below.
Suppose you have 100 ml of CHCl3, which is a well over 100 grams of CHCl3, making up for appr. 1 mole of CHCl3. You then need appr. one mole of Cl2
gas, to convert all of it to CCl4. At standard pressure and at room temperature a rough estimate tells me that you'll need somewhere between 20 and 25
liters (!!) of pure Cl2 gas to convert all of this to CCl4. Dissolving all that Cl2 in your solvent must be done slowly, probably in multiple stages.
Maybe a slowly working chlorine generator, with the gas dried through a column of anhydrous CaCl2 and bubbling it through CHCl3 on a warm and sunny
day does the job, but it will be quite an endeavor for the average home chemist with minimal glassware and tubing. I'm not sure about the speed of the
reaction.
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AJKOER
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One answer is to prepare Chlorine hydrate (Cl2. 8 H2O). To quote:
"If chlorine gas is passed into a dilute solution of CaCl2 at about 0 °C, greenish, feathery crystals appear that can be removed from the solution,
dried, and kept indefinitely at room temperature. If these crystals are dissolved in water, chlorine gas is liberated. If they are heated at one end
of a closed tube, the other end of which is immersed in ice water, liquid chlorine appears at the cold end. These crystals are chlorine hydrate,
discovered by Sir Humphrey Davy in 1811 shortly after proving that chlorine is an element, and studied in more detail by Michael Faraday in 1823. The
formula of this substance is Cl2·H2O."
Link: http://mysite.du.edu/~jcalvert/econ/hydrates.htm
Also, see for physical properties:
http://books.google.com/books?id=-2ObmTZTq2QC&pg=PA272&a...
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Mixell
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I am going to use the chlorine for chlorination and making compounds such as anhydrous SnCl4, I do not want water to be involved...
But that article is very interesting indeed, I might try making some chlorine hydrate as a novelty synthesis.
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woelen
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Chlorine hydrate is a nice curiousity and a similar compound can be made from bromine and water. Nice to play with, but not of any practical value.
The problem is that these compounds are only stable at low temperatures. Do not put a lot of chlorine hydrate in a tightly sealed glass vial, which is
taken out of the fridge. The compound easily decomposes, giving gaseous Cl2 and water and when it is stored in a tightly sealed bottle the pressure
may become so high that the bottle explodes.
Even though you cannot store chlorine hydrate safely in a sealed vial, it is nice to play with it and do interesting experiments with it. So, I
encourage you to make some and study its properties, but do not regard it as a useful tool to keep chlorine at hand at high concentration in
non-gaseous form.
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Endimion17
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For this, and other interesting details on the properties of elements and their isolation, and synthesis of basic compounds, I encourage you to take a
look at Mellor Chemistry. It's a golden chest of really interesting informations. Wikipedia can kiss its ass, believe me.
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AJKOER
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Quote: Originally posted by woelen  | | Chlorine hydrate is a nice curiousity and a similar compound can be made from bromine and water. Nice to play with, but not of any practical value....
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First, more background on Chlorine water from an older source:"Chlorination of water" by Joseph Race, page 103:
"Although chlorine water appears to be of little value because of its instability there appears to be no reason why chlorine hydrate should not be
successfully employed. The hydrate was first prepared by Faraday [9] by passing chlorine into water surrounded by a freezing mixture. A thick yellow
magma resulted from which the crystals of chlorine hydrate were separated by pressing between filter paper at 0° C. .......
Chlorine hydrate separates into chlorine gas and chlorine water at 9.60 C. in open vessels and at 28.70 C. in closed vessels. "
Another source: "Treatise on general and industrial inorganic chemistry", Volume 1, by Ettore Molinari, page 122:
"The system chlorine-water exhibits the following interesting cases: saturation of water at 0° with gaseous chlorine leads to the separation of
greenish crystals of chlorine hydrate (Cl2 + 8H20), these being stable only up to the temperature 9.6°, above which they are decomposed into gaseous
chlorine and water saturated with chlorine; this hydrate is, however, stable at higher temperatures if the pressure is raised and, on the other hand,
decomposes below 9° if the pressure diminishes. As a rule there are three phases for chlorine hydrate, but at a certain definite point, namely, —
0.26°, four phases are possible, i.e., ice, chlorine hydrate, aqueous chlorine solution, and gaseous chlorine. The equilibrium is, however, easily
altered, and at the least rise of temperature the ice disappears, whilst lowering of the temperature causes transformation of the aqueous chlorine
solution into ice and chlorine hydrate and the consequent suppression of one of the four phases; this equilibrium depends also on the pressure (24.4
cm.), and if this increases, the gaseous chlorine disappears, whilst if it diminishes, another phase (the ice or the hydrate) disappears."
Now, with respect to Chlorine water versus Cl2.8 H2O, I believe I can make an argument to the contrary (it is not a novelty). Apparently, Chlorine is
not very soluble in water and even if you consider all available sources of Cl (including Cl2, HOCl and hypochlorite), for sea level at 10 C only 9.65
grams of total Chlorine are available per liter of water (see page 71 at http://books.google.com/books?id=9_2idzksARMC&pg=PA71&am... ). This is less than 1%.
Now with respect to Cl2. 7.3H2O (which is apparently the more precise formula) my calculation is that the total available Cl as Cl2 constitutes 35% by
weight. So for 1,000 ml of Chlorine hydrate (density 1.23) the total Chlorine is 431 grams versus less than 10 grams for Chlorine water.
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woelen
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Yes, if chlorine hydrate would be a more stable compound, then it would be a really valuable concentrated source of chlorine. Unfortunately it is
unstable and can only be kept, either at low temperature, or at high chlorine pressure.
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