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Detonationology
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Quote: Originally posted by blogfast25  | | You're conflating purity and concentration. A solution could contain almost only, say, B and solvent and yet B's concentration could be low and
variable. Worksheets can be adjusted for concentration of reagents, usually very easily. |
How could the concentration of a solution be "low and variable" if the solution contains "almost only" a salt (I would assume) dissolved in a
solution? I do not follow this "B" analogy. Could you please explain with a relative example how purity and concentration are not related?
[Edited on 12-16-2015 by Detonationology]
“There are no differences but differences of degree between different degrees of difference and no difference.” ― William James
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blogfast25
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I didn't say purity and concentration are not related, I said you're conflating them, i.e. they're not the same.
Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some
impurity (obviously not water, and not ammonia either!). Those impurities could be expressed with respect to the ammonia (e.g. '99.9 % ammonia'),
while also expressing the concentration of ammonia with respect to the solvent (e.g. '25 w% NH3' or '2M NH3').
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MolecularWorld
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| Quote: Originally posted by blogfast25 | | Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some
impurity (obviously not water, and not ammonia either!). |
This is misleading. A highly pure ammonia solution left to evaporate in air may leave a residue from the reaction of ammonia and atmospheric carbon
dioxide. If you meant "...in a stream of dry air scrubbed of carbon dioxide" or "...in a desiccator with [lots of] sodium hydroxide" you should have
said so.
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JJay
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| Quote: Originally posted by MolecularWorld | @OP: Good job!
We all know what the product is, and that it's purity is acceptable for the intended use, but my experience requires me to ask the following pedantic
questions:
How do you know the CuCl2 is anhydrous? What tests did you perform to prove this? Did you test for decomposition after heating? You didn't even state
the temperature or the length of time spent in the "oven".
You mention sourcing reagents locally. I interpret that to mean your reagents are OTC. Everybody knows that all OTC products are highly contaminated,
enough to throw off the results of most experiments (I didn't think so, but I was wrong). You must repeat the experiment with reagent grade chemicals,
and/or verify the product spectroscopically, to ensure contaminants didn't give a misleading result.
What was the purity of the product? How did you determine this?
How do you know your product is anhydrous copper(II) chloride? Don't tell me you judged this based mostly on color, that would be "beneath contempt."
Edit: Also, those one-piece plastic mason jar lids are far from air-tight (I also use them for chemical storage). You'd be better off using a regular
gasketed lid, otherwise your "anhydrous" product soon won't be (unless you store the jar in a desiccator).
"Talk about an aversion to evidence..."
[Edited on 16-12-2015 by MolecularWorld] |
These are fair questions. I have actually not tested to see if the CuCl2 is anhydrous, but it would be possible to do so by measuring the specific
gravity of a solution of known concentration. I baked it for several hours, at around 150 C until it was dry (some green crystals formed) and then at
300 C with stirring every half hour until everything turned brown and the clumps stopped sticking together, indicating that no more water was being
emitted. The stirring is important to prevent the formation of a fused mass sticking to the crystallization dish.
The storage container could probably be sealed a little tighter with a gasket or PTFE tape, but I don't think CuCl2 is very hygroscopic, so I haven't
worried about it. (And come to think of it, I have stored hygroscopic compounds with those lids for months with no issue.) Your point is well taken,
though - I would use a better container for storing aluminum chloride. I think the cupric chloride purity is actually pretty good, but it would be
possible to recrystalize the CuCl2 as the dihydrate salt from water and then dehydrate it... or it may be feasible to recrystalize the anhydrous form
from an alcohol.
The reagents are OTC. The calcium chloride is not food grade but appears to be reasonably free of impurities and consistently behaves as expected in
other experiments. This is the first time I have used this copper sulfate, but it also performed exactly as expected; I have little reason to believe
that it contains a lot of impurities.
I am not sure how to best verify the CuCl2 spectroscopically with equipment available to the average hobbyist, but I like that idea. I wonder if ACS
publishes standards for CuCl2....
I don't know what the purity is... I don't think it is analytical grade, but I do suspect it is quite good.
[Edited on 17-12-2015 by JJay]
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JJay
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Say, CuCl2 is a little more hygroscopic than I thought... it turns blue-green as it absorbs water. Pretty interesting stuff... mine is showing no signs of changing color in its current container.
[Edited on 17-12-2015 by JJay]
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blogfast25
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| Quote: Originally posted by MolecularWorld | | Quote: Originally posted by blogfast25 | | Take an ammonia solution. Highly pure it would evaporate to literally nothing. But if some solids are left on evaporation that would point to some
impurity (obviously not water, and not ammonia either!). |
This is misleading. A highly pure ammonia solution left to evaporate in air may leave a residue from the reaction of ammonia and atmospheric carbon
dioxide. If you meant "...in a stream of dry air scrubbed of carbon dioxide" or "...in a desiccator with [lots of] sodium hydroxide" you should have
said so. |
To illustrate a simple principle it isn't necessary to provide a complete methodology, MW. And evaporate could mean 'boil off' too...
Context is everything. I was explaining something by simple example, not putting up a method of analysis.
[Edited on 17-12-2015 by blogfast25]
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blogfast25
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| Quote: Originally posted by JJay |
I don't know what the purity is... I don't think it is analytical grade, but I do suspect it is quite good.
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Contaminants in your starting reagents will partly have been carried over.
One contamination you can bet on is CaSO4 because it has limited solubility, so your CuCl2 solution will have contained some.
BaCl2 would have been better in that respect, because BaSO4 is far more insoluble.
You could use BaCl2 or Ba(NO3)2 to test your product for sulphates.
CaSO4 solubility at 20 C, about 21 g/L (water, Wikipedia).
[Edited on 17-12-2015 by blogfast25]
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JJay
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I am going to try dissolving the CuCl2 in acetone and reacting it with aluminum in an ice bath. I suspect that CaSO4 contamination is quite low due to
its low solubility and the tendency for it to get salted out by more soluble compounds. I don't remember its solubility in water offhand, but 21 g/L
is about 10x higher than I had thought.
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DraconicAcid
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| Quote: Originally posted by JJay | | I suspect that CaSO4 contamination is quite low due to its low solubility and the tendency for it to get salted out by more soluble compounds.
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You can't salt out an ionic compound from solution, unless the salt you're using is a calcium salt or a sulphate (common ion effect). If you increase
the ionic strength of the solution without a common ion, you'll just make it more soluble.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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JJay
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| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by JJay | | I suspect that CaSO4 contamination is quite low due to its low solubility and the tendency for it to get salted out by more soluble compounds.
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You can't salt out an ionic compound from solution, unless the salt you're using is a calcium salt or a sulphate (common ion effect). If you increase
the ionic strength of the solution without a common ion, you'll just make it more soluble. |
That is simply not the case
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DraconicAcid
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Oh? Explain your reasoning.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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mayko
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Wikipedia lists the water solubility of CaSO4 as:
0.21g/100ml at 20 °C (anhydrous)[1]
0.24 g/100ml at 20 °C (dihydrate)[2]
By my math, that's ~2g per liter.
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AJKOER
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Per Atomistry on AlCl3 ( http://aluminium.atomistry.com/aluminium_trichloride.html ):
"Aluminium trichloride, AlCl3, was originally made by heating an intimate mixture of alumina and carbon to redness in a stream of chlorine (Oersted's
method). It may be more readily prepared by heating aluminium in a wide glass tube in a rapid current of dry hydrogen chloride, or in a stream of
chlorine. If it is required to prepare the chloride from the oxide, a neater method than Oersted's is to heat the oxide in a current of chlorine and
sulphur chloride: -
4Al2O3 + 3S2Cl2 + 9Cl2 = 8AlCl3 + 6SO2.
Instead of chlorine and sulphur chloride, carbon tetrachloride vapour or carbonyl chloride may be used. A simple method of preparation is said to
consist in heating crude alumina or clay to redness in a current of hydrogen chloride and carbon disulphide vapour, and purifying the aluminium
chloride so obtained by sublimation over iron filings. "
Per Atomistry on COCl2 ( http://carbon.atomistry.com/carbon_oxychloride.html ):
"Carbonyl chloride may also be prepared by the oxidation of chloroform by chromic acid, when 20 parts of chloroform, 400 of sulphuric acid, and 50 of
potassium dichromate are heated together on the water-bath:
2CHCl3 + K2Cr2O7 + 4H2SO4 = 2COCl2 + K2SO4 + Cr2(SO4)3 + 5H2O + Cl2"
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blogfast25
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| Quote: Originally posted by mayko | Wikipedia lists the water solubility of CaSO4 as:
0.21g/100ml at 20 °C (anhydrous)[1]
0.24 g/100ml at 20 °C (dihydrate)[2]
By my math, that's ~2g per liter. |
Ooopsie. Also by my math. 
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blogfast25
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DA is completely correct, JJay.
But the contamination by CaSO4 will be small and won't interfere here.
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JJay
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No, he is not. But the contamination by CaSO4 will be very small indeed.
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DraconicAcid
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Citation? Or reasoning?
If you increase the ionic strength of a solution, you will decrease the activity of the ions, so they basically act as if they have a lower
concentration. This makes ionic compounds *more* soluble in solutions of high ionic strength.
It is known.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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JJay
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Let me put it to you this way: you can't look at ionic strength alone to determine solubility. You have to look at activity (which is sometimes
expressed as a function of ionic strength) and concentration. Ordinarily, the effect of increased ionic strength decreases as concentration increases.
Also, copper (II) forms complexes with water.
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blogfast25
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| Quote: Originally posted by JJay | Let me put it to you this way: you can't look at ionic strength alone to determine solubility. You have to look at activity (which is sometimes
expressed as a function of ionic strength) and concentration. Ordinarily, the effect of increased ionic strength decreases as concentration increases.
Also, copper (II) forms complexes with water. |
You sound confused, TBH. No one said 'alone' but increased ionic strength increases solubility (except in the case of common ion effects).
See e.g.:
| Quote: | | The theory of activity versus concentration is important in industrial, environmental, and biochemistry. The increase in solubility of an electrolyte
in a solution of a second electrolyte with no common ions compared with pure water is not an easy concept to grasp because it seems to be
counterintuitive. The simple experiment described here illustrates this principle visually and dramatically. Students attempt to dissolve CaSO4•2H2O
(gypsum) in pure water and in 0.25 M NaCl. The gypsum dissolves almost completely in the sodium chloride solution, but not in pure water. Students
then measure the calcium concentrations in filtered aliquots of both solutions to quantify the solubility difference they observed. Students calculate
mean activity coefficients using their measured concentrations and also from the Davies Equation, an extension of Debye–Hückel theory. The basic
principle is there are ionic interactions between the solute ions and the solvent ions, which allow for more dissolution because only free ions enter
into the expression for the solubility product equilibrium constant. From a simple mathematical point of view, in higher ionic strength solutions,
activity coefficients for calcium and sulfate become smaller, and hence the concentrations must be larger to maintain a constant solubility product at
equilibrium. |
Source.
Yes, Cu(+2) forms hexaaqua coordination complexes. What has that got to do with anything here, though? It also forms weak tetrachloro coordination
complexes but even in pure CuCl2 solution not so much.
[Edited on 18-12-2015 by blogfast25]
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JJay
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I can put it into some terms that you will understand. Let's say that Sam, Jill and Jen are ions. They don't like hanging out at the bookstore much,
especially when their annoying relatives are there. They are more likely to hang out at the bookstore when members of the opposite sex are present.
But they don't hang out at the bookstore when it is packed.
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blogfast25
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| Quote: Originally posted by JJay | | I can put it into some terms that you will understand. Let's say that Sam, Jill and Jen are ions. They don't like hanging out at the bookstore much,
especially when their annoying relatives are there. They are more likely to hang out at the bookstore when members of the opposite sex are present.
But they don't hang out at the bookstore when it is packed. |
And this is your theory? Or based on what mainstream chemistry theory?
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blogfast25
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Basic theory of activities, ionic strength and solubilities
Why believe your system would behave differently?
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JJay
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That is how those activity-concentration equilibrium equations work. It's not *my* theory.
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blogfast25
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With ions that are attracted to the opposite sex? Whatever makes you happy, I guess...
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JJay
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by JJay | It looks like perhaps the easiest way to react CuCl2 with aluminum while capturing the product in an anhydrous state is to dissolve the copper (II)
chloride in dry ethanol and then drip it onto the aluminum. I think aluminum chloride should go into solution with the copper precipitating out. Once
the reaction is complete, the reaction mixture could be filtered and the aluminum chloride precipitated out of the ethanol with chloroform. Right?
[Edited on 16-12-2015 by JJay] |
I'm wondering, just wondering, if AlCl3 with ethanol will not perhaps give Al ethoxide? Or at least partly? Hmmm... |
It does, or at least it gives a strong complex that behaves very similarly to Al ethoxide... so chloroform probably won't precipitate the AlCl3... It
looks like acetone might work instead of alcohol... I suspect that AlCl3 will form complexes with acetone too, but likely they can be cleaved with a
boiling water bath under vacuum (rather than in a stream of HCl gas at 300 C).
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