I'm going to take an extreme case here but bear with me.
Lets assume that you have Borax that is 10% impure, or onlly 90% pure! Lets also assume that the impurity is Sodium Chloride.
A saturated Borax solution will have less Borax in it than expected due to the common ion effect. This causes the less soluable of two sodium salts to precipitate more than predicted.
The 10% Sodium Chloride would become MORE saturated in solution and therefore the impurity level would be increased.
So, your saturated solution is less pure than the starting example.
This hypothetical example shows what you can do to yourself if you don't know the basic guidelines of what you are purifying or how you are purifying it.
Now, if there were a less soluable sodium salt, such as Sodium Aluminate (IIRC), it would concentrate in the solids that remained behind. So the precipitated Borax would end up with more of the alumiate salt.
I'm going to take an extreme case here but bear with me.
Lets assume that you have Borax that is 10% impure, or onlly 90% pure! Lets also assume that the impurity is Sodium Chloride.
A saturated Borax solution will have less Borax in it than expected due to the common ion effect. This causes the less soluable of two sodium salts to precipitate more than predicted.
The 10% Sodium Chloride would become MORE saturated in solution and therefore the impurity level would be increased.
So, your saturated solution is less pure than the starting example.
This hypothetical example shows what you can do to yourself if you don't know the basic guidelines of what you are purifying or how you are purifying it.
Now, if there were a less soluable sodium salt, such as Sodium Aluminate (IIRC), it would concentrate in the solids that remained behind. So the precipitated Borax would end up with more of the alumiate salt.
Either way, you can make things worse.
PE
The basic guidelines of Technical grade borax powder are well known, and the worst case I have seen is guaranteed to have less than 700 ppm of chloride. Under what circumstances could that amount of chloride NOT be dissolved in the first solution? I am depending on the total chloride, sulphate and iron which total 1330 ppm max in the original charge of borax to be dissolved in the first charge of water, 95% of which will be removed by decanting it. That solution certainly would yield a more contaminated borate powder than the original if it were evaporated. I would in fact be happy if it had only the soluble impurities and no borax. Instead, it will be discarded or used for drain cleaner or one of several other household uses. How can what remains in the undissolved borax not have less chloride than it had before? What if it does have more insoluble impurity? I am not using the insoluble impurity that remains. I am making a second saturated solution from it, leaving insoluble impurities behind. The second saturated solution removes doubt about hydration of the powder and MUST have smaller percentages of chloride, sulphate and iron. Why, in the name of good science, don't you read thoroughly what I wrote?
Patrick, using the simple principle of solubility and the common ion effect, I have explained to you how you can either concentrate an impurity in the mother liquor or in the solid. It is that simple. To do what you are describing, you must know what you are doing. If I cannot explaine some simple facts of chemistry to you, then I'm sorry.
You have to know what you are about, before you undertake it.
And, I said that what I had given was hypothetical, but nonetheless, your procedure as outlined will concentrate all halide and phosphate impurities along with other soluable sodium salts in the concentrated Borax solution.
Here is an example. Let us assume compund X is 90% pure and dissolves in water at the rate of 50 g/l. It contains Y which is soluable at 100 g/l as an impurity.
Therefore, 50 g of X contains 5 g of Y.
If you take 100 g of XY mixture and place it in 100 ml of water, you will dissolve 50 grams of X and ALL OF Y. This means that the solution is now a saturated solution of X which is now 10% impure. By using bad procedure, you have doubled the impurity.
If you are after the solid, it is more pure than that in the liquid, but if you assume a third ingredient Z which is less soluable than X, then you have concentrated the Z in the solid as well.
Patrick, using the simple principle of solubility and the common ion effect, I have explained to you how you can either concentrate an impurity in the mother liquor or in the solid. It is that simple. To do what you are describing, you must know what you are doing. If I cannot explaine some simple facts of chemistry to you, then I'm sorry.
You have to know what you are about, before you undertake it.
And, I said that what I had given was hypothetical, but nonetheless, your procedure as outlined will concentrate all halide and phosphate impurities along with other soluable sodium salts in the concentrated Borax solution.
I'm sorry, but that is a chemical fact.
PE
And it is also a chemical fact that the solution I am keeping is not the solution you are talking about. I am using the sediment from it to make a second saturated solution which cannot, according to your own lecture, be as impure as the first because the first has dissolved and carried away 95% of the soluble impurities. You apparently do not know what I am doing. Read it again or forever hold your peace on this subject. I gave a numerical example illustrating exactly what tou are preaching. The first saturated solution carries with it the soluble impurities in the 200 grams of borax, while dissolving only about 40 grams of borax. The remaining 160 grams of sediment contain only about 5% of the original soluble impurities. The first saturated solution is discarded or used around the house. The 160 remaining grams of borax are sufficient to make at least 3 liters of saturated borax solution. Have I said anything wrong so far? I'm beginning to think you are the one who flunked out of chemical engineering. This method of using differential solubilities to purify is very old. If I had a solvent that would dissolve everything but borax, I could eventually get very pure borax without losing any, but it would not be in the sediment. It would be in the last saturated water solution I made.
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Author: American National Standards Institute.
Title: American national standard specification for photographic grade sodium tetraborate pentahydrate and sodium tetraborate decahydrate.
Published: New York, N.Y. : American National Standards Institute, c1982.
Physical Description: 3 p. ; 28 cm.
Subject (LCSH): Photography --Standards.
Photographic chemicals --Standards.
Other Name: National Association of Photographic Manufacturers, Inc.
Other Title: Specification for photographic grade sodium tetraborate pentahydrate and sodium tetraborate decahydrate.
Notes: Caption title.
"Revision of ANSI PH4.230-1976."
"ANSI PH4.230-1982."
"Approved February 26, 1982."
Other Identifying Number: ANSI PH4.230-1982
Come to think of it, people in Virginia use it on Smithfield hams. The ham must be soaked in water for quite some time, the water removed and replaced by fresh, and the process repeated several times to get the salt to the point where a person can safely eat the ham. And you pay a lot for the privilege of doing it.
And it is also a chemical fact that the solution I am keeping is not the solution you are talking about.
I missed that you were tossing the first leaching of the tech grade borax on the first reading myself. I did not find it clearly stated it was discarded. Anyway...
I still have issues with using saturated solutions as a form of calibration. It's just so tricky using them as thye are so dependant on temperature. It's just so much simpler to weigh something out and dissolve it in solvent. I can't think of any analytical procedures where a saturated solution of anything is used as a "calibrated" solution. They are usually used as a reagent solution that is being added to a solution in order to supply an excess of a reagent, never to supply a precise amount. If you are careful and use good temperature controls and make sure you don't get your solution too cool and crash out your reagent, you could use a saturated solution as you describe.
But it really looks like a pain... Just buy a grade of high quality reagent and work from there. None of these mechanations to turn a sow's ear into a silk purse. (There, I was able to tie it back in with your ham comment!)
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