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February 28

Chemical purification, part 2

Of the many chemicals used for the liquid-liquid extraction of iron from phosphoric acid, which ones have the highest flash points? 2601:646:8E01:7E0B:0:0:0:64DA (talk) 05:07, 28 February 2018 (UTC)[reply]

If you have a list of chemicals, this sort of information can usually be found on the SDS, on the packages themselves, or our wikipedia article for each. But we can't read your mind to know what specific chemicals you mean or even what ref you are reading. DMacks (talk) 17:16, 28 February 2018 (UTC)[reply]

From what I have read on liquid-liquid extractions, the common liquid used is kerosene, with some other chellating substance dissolved in it. Usually this is optimised to be effective and cheap. If you do your own research you may find high flash point non-polar solvents that may work eg perfluorocarbons or chlorocarbons, but they are probably undesirable for other reasons. Graeme Bartlett (talk) 23:15, 28 February 2018 (UTC)[reply]

Reference: https://patents.google.com/patent/US3694153 look for kerosene. Graeme Bartlett (talk) 23:27, 28 February 2018 (UTC)[reply]

Interesting! I had another process in mind, where the acid is first mixed with table salt to complex the iron, and then the iron chloride complex is extracted (I found it in an old book in the university library, I don't remember its name but I can give it to you tomorrow), but this MIGHT be even better! Will it work on 85% acid, or will I have to dilute it first? 2601:646:8E01:7E0B:0:0:0:64DA (talk) 08:05, 1 March 2018 (UTC)[reply]

If you read the patent, it says that the acid become too viscous if it is too concentrated. Graeme Bartlett (talk) 12:43, 1 March 2018 (UTC)[reply]

All right, that's not a big deal -- I'll have to dilute the acid before use anyway, so I can just purify it after dilution. And kerosene is not too flammable for me, so I guess that's what I'll be using as my solvent (although, if diesel fuel would work, I'll use that because it's more readily available). 2601:646:8E01:7E0B:0:0:0:64DA (talk) 06:06, 2 March 2018 (UTC)[reply]

Chemical purification, part 3

Is it possible to remove iron from aluminum hydroxide and copper oxide to levels of 1 ppm or less by selectively dissolving the material to be purified (in, say, sodium hydroxide for the aluminum hydroxide or ammonium carbonate solution for the copper oxide), physically removing the undissolved iron oxide/hydroxide (e.g. by filtration), and then recovering the purified aluminum hydroxide or copper oxide from the solution? Or do other techniques have to be used for this level of purity? 2601:646:8E01:7E0B:0:0:0:64DA (talk) 05:16, 28 February 2018 (UTC)[reply]

For your ammonia solution see https://pubs.acs.org/doi/abs/10.1021/i260041a028 where the answer is no, some iron would probably dissolve in your solution. Graeme Bartlett (talk) 23:42, 28 February 2018 (UTC)[reply]

Thanks for the bad news! And for the mini-Bayer process? 2601:646:8E01:7E0B:0:0:0:64DA (talk) 08:06, 1 March 2018 (UTC)[reply]

I couldn't easily find any sources for that. But if you really oxidise your iron to the red mud stage it is really insoluble. But then it depends on how good your filtration is. Even the dust blowing over your solution could add a part-per-million of iron! But for your copper extraction there appears to be more complex solutions and processes that will get the purity. Graeme Bartlett (talk) 12:21, 1 March 2018 (UTC)[reply]

OK here it says it will contain a few hundredths of one percent of iron, ie a few 100 parts per million in the alumina at the precipitation stage. But this page http://sourcedb.ipe.cas.cn/zw/lwlb/200909/P020090909606185656884.pdf says that you can precipitate more iron by adding methanol. So there will be ways around the problem. Graeme Bartlett (talk) 12:41, 1 March 2018 (UTC)[reply]

Whoa, methanol, I'd rather not go there -- I'll be working near gas appliances, using this stuff would be an accident waiting to happen! 2601:646:8E01:7E0B:0:0:0:64DA (talk) 06:08, 2 March 2018 (UTC)[reply]

... but your kerosene has a much lower autoignition temperature (though I agree that flashpoint of methanol could be a problem). You could always precipitate the iron in a safe environment, then evaporate the methanol safely. Dbfirs 11:48, 2 March 2018 (UTC)[reply]

...except there is no other place for me to do the purifications -- the only other available area is off-limits to any work with hazardous chemicals because there's food being prepared there. And autoignition temperature is mainly a problem if I actually spill the kerosene on the gas appliances (which most emphatically WILL NOT happen -- my work area is a fair distance away), whereas the flash point and/or fire point is a much bigger hazard -- the way the whole thing is set up, the main fire hazard is if the methanol vapors get sucked into the gas appliances and ignite! 2601:646:8E01:7E0B:59E2:B6:A6B8:FAC7 (talk) 09:41, 3 March 2018 (UTC)[reply]

You have to consider why you are doing this. Different grades of chemicals are available for different applications. For food use, buy food grade, where a little bit of iron will not matter. For micronutrient growth experiments you will really need known low levels of important nutrients. There is also spectrophotometry grade materials, that will have very low colour levels in certain wavelength ranges. For serious chemical experiments you will probably have to purify it yourself, and not get advice from random people on the internet. For industrial application employ a qualified chemical engineer. These separations using kerosene also need much harder to get ligands to dissolve in it, which will cost you plenty, and probably more than then small amount of pure material that you want to make. Electrochemical methods may be more suitable to purify copper. And dissolving pure aluminium or copper may be a cheaper way to get your pure compounds. Graeme Bartlett (talk) 23:19, 3 March 2018 (UTC)[reply]

This is a specialty material synthesis (namely, synthetic turquoise) -- and in my case, most impurities don't matter but iron is critical because even tiny amounts can discolor the product. 2601:646:8E01:7E0B:59E2:B6:A6B8:FAC7 (talk) 02:45, 4 March 2018 (UTC)[reply]

Calculating the activation energy of a redox reaction

The reaction in question involves a galvanic cell setup. If I know the electrode potentials of the reactants, can I multiply that by the total charge flow (given by the area under a current-time graph) to get the activation energy? I think this because potential*charge = potential energy which I suspect is the energy needed for the reaction to proceed. The Average Wikipedian (talk) 15:10, 28 February 2018 (UTC)[reply]

This set of lecture slides describes how to calculate the activation energy of an electrochemical reaction quantitatively, and seems to directly answer your query. --Jayron 32 15:14, 28 February 2018 (UTC)[reply]

Why is estradiol cypionate so much more potent than estradiol valerate?

AFAIK, estradiol cypionate is not active in itself and neither is the valerate ester: they're both prodrugs for estradiol. However, 5 mg of estradiol cypionate apparently is just as effective as 20 mg of estradiol valerate and estradiol cypionate is not manufactured at higher dosages. What is the reason behind this increased efficacy? Yanping Nora Soong (talk) 21:10, 28 February 2018 (UTC)[reply]

According to cypionic acid:

The lipophilicity of the cypionate group allows the prodrug to be sequestered in fat depots after intramuscular injection.[2] The ester group is slowly hydrolyzed by metabolic enzymes, releasing steady doses of the active ingredient.

So maybe it's not so much potency as time-release? --21:15, 28 February 2018 (UTC)

To elaborate a bit on the answer; this is a pharmacokinetics issue rather than a pharmacodynamics issue. The way that the cypionate is metabolized affects how the drug is released into the body; the drug itself has the same effects, but the body's action on the two molecules is different. --Jayron 32 00:23, 1 March 2018 (UTC)[reply]

This is the sort of thing one takes a full course in organic chemistry, plus a full course in psychopharmacology for. I am not sure the ref desk is certified to teach such courses, and I have had both, but at an Ivy-League university, and I had to repeat Organic because a mere one out of ten of my syntheses blew up, virtually guaranteeing a D grade, which would count as an F toward the Bio major. Enrole in a good STEM school, then get back with the stuck point. μηδείς (talk) 03:02, 1 March 2018 (UTC)[reply]

Actually our questioner sounds to be the same kind of student. And so probably has access to better experts on campus. Graeme Bartlett (talk) 12:15, 1 March 2018 (UTC)[reply]

Honestly, I think the graph in estradiol cypionate sums it up very well, and an expert would bring little to the table -- at least on this narrow issue of dosage -- other than by finding data of that type. Medical dosage recommendations are often quite arbitrary, really - consider the history of the birth control pill. [1] To call dosage empirical is too generous - empirical means that people trying it have found that X works; but medical testing means more like "somebody tried it once and found that X works and it will cost more than you can possibly spend to have the bureaucrats revisit the issue." Wnt (talk) 03:09, 2 March 2018 (UTC)[reply]

Yes, it is a good thing that the two above users have learned to imitute me exalctly. μηδείς (talk) 04:53, 2 March 2018 (UTC)[reply]