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February 11

Lemongrass oil as a chemical solvent?

I just stepped in some spilled lemongrass oil at a supermarket checkout and was surprised to find my shoe dissolving in it. Which got me to wondering about the properties of this and complex mixtures of compounds as solvents in general...

1) I found an analysis [1] that shows that the oil is about 75% geranial and neral, which together are called citral, with a touch of beta-myrcene and 50 other compounds making up most of the remaining 25%. Would the presence of a wide range of compounds make it a significantly better solvent than the pure citral components or a citral mixture?

2) What exactly makes a solvent capable of dissolving a shoe? I remember reading that gamma-butyrolactone, once a common industrial solvent, could do it. Can it be described simply in terms of how hydrophobic the solvent is or is it more empirical than that? (Honestly, I don't even know what shoe "rubber" really is, sorry)

3) Citral has a high boiling point, over 200 C, unlike the organic solvents I readily think of (though looking at the gamma-butyrolactone article I see it describes chemical solvent use, without elaborating). Are high-BP solvents considered for syntheses (say, for a recrystallization) or do drying difficulties just rule it out? I mean, I can picture that if you don't care about product purity per se you might tolerate a non-toxic citral contamination so long as you can wash out impurities from the synthesis reaction itself, or if you followed up with washes with a low-BP solvent that doesn't dissolve the desired product at all... not sure if there's a practical situation where people would want to use it or not. Wnt (talk) 00:36, 11 February 2018 (UTC)[reply]

2) Yes, the less polar the solvent is, the better it will be at dissolving rubber and other (non-polar) polymers (found this out the hard way in o-chem lab with some methylene chloride and my calculator). As for (3), you might want to check out the articles on the Bergius process and on coal liquefaction in general. 2601:646:8E01:7E0B:0:0:0:64DA (talk) 06:48, 11 February 2018 (UTC)[reply]

Colour of tetracyanidonickelate(II)

Potassium tetracyanidonickelate (II) has a very light yellow colour and the tetracyanidonickelate (II) ion is colourless in solution. It can be very easily explained with valance bond theory as d^8 configuration has no vacant d shell for electron rearrangement using the energy of incident photon. But if I apply crystal field theory, I see the d x2-y2 orbital is vacant on-ground state and electron rearrangement is possible there. But the prediction for no colour is based on a completely different phenomenon (to my assumption) that so much rearrangement is possible and it can absorb from all regions of visible spectrum equally to produce no colour. These two explanations are completely different description of same phenomenon. Which one is correct?Sayan Ghosh 04:57, 11 February 2018 (UTC) — Preceding unsigned comment added by Sayan19ghosh99 (talk • contribs)

Not answering your question, but our article on this is tetracyanonickelate. It is interesting that the calcium salt colour depends on polarization. Also the picture of the solution shows it as light yellow, so perhaps the tetracyanonickelate ions have to be oriented in a fixed way (in a crystal) to experience colourlesness from some direction and polarization. When in a solution they would be randomly aligned. These ions are square planar. Contrast this with tetrabromonickelate which has tetrahedral bromine and is intensely coloured. For tetrachloronickelate colour depends on temperature. Graeme Bartlett (talk) 11:59, 11 February 2018 (UTC)[reply]

I think both the valance band theory and crystal field theory agrees with the colour of the tetrabromonickelate in same qualitative manner, i.e. d-d electron transition. I think the answer to my question can be given from the data of absorption spectrum of the tetracyanidonickelate(II) complex in solution.Sayan Ghosh 12:40, 11 February 2018 (UTC) — Preceding unsigned comment added by Sayan19ghosh99 (talk • contribs)

Another comment, if something absorbs from all regions of the visible spectrum, then it will appear black. So your colourless substance absorb nothing from the visible spectrum. Graeme Bartlett (talk) 23:21, 11 February 2018 (UTC)[reply]

Yes, it should, but I think it can not absorb 'all the colours'. It absorbs from 'all regions of visible spectra' which make up white colour (like you mix just red and blue and green and leave other colours). But surely I can not claim this to be true. Sayan Ghosh 03:09, 12 February 2018 (UTC) — Preceding unsigned comment added by Sayan19ghosh99 (talk • contribs)

In your example if you absorb red blue and green, and then leave violet, yellow, and orange you will get some some colour, perhaps grey, it will not be clear. Graeme Bartlett (talk) 08:06, 12 February 2018 (UTC)[reply]

PubChem gives the main name as "Tetracyano nickel", oddly enough [2], and even more surprisingly, lists "Tetracyanidonickelate(0); Tetracyanonickelate(4-); Tetracyanidonickelate(4-)" as synonyms. I don't know what's up with that exactly, but I assume we're all talking about 4 CN^- ligands associated with what, according to our article and the OP's notation, is one nickel (II) = Ni²⁺ for a net charge of -2. As its configuration is [Ar] 3d⁸ 4s² or[Ar] 3d⁹ 4s¹, but apparently [3] all the electrons move to d orbitals when considered in association with ligands and it is regarded as a "d8 complex" which preferentially forms a square planar molecular geometry. Tetracyanonickelate is diamagnetic and planar whereas tetrachloronickelate is paramagnetic and tetrahedral (but PdCl4 2- is also square planar in part because the pairing energy of two electrons is less in more diffuse orbitals). [4] Now the tetrachloronickelate is thermochromic, looking a lot like tetracyanonickelate at room temp where our article says polymers are present but turning bright blue at higher temperatures where it is tetrahedral. Admittedly this doesn't help me much, except to realize that we would actually have to look up stuff like this about individual spectral lines to have a hope of really knowing what's going on. (Well, me anyway, maybe not the others here) Looking at these papers, I get the distinct suspicion that most if not all of the theory is applied post hoc to rationalize a wide range of behaviors from transition metal complexes, but I don't know how true that is. Wnt (talk) 13:37, 12 February 2018 (UTC)[reply]

What is wrong with hamsters?

Why do they enjoy pointlessly running round that exercise wheel so much? They don't get any reward for it: no food, no sexual partner to display to. No one trains them to do it. Perhaps they are genetically predisposed to keep fit. Maybe we should isolate that gene and inject it into people who are genetically predisposed to lay on the sofa eating doughnuts. SpinningSpark 13:33, 11 February 2018 (UTC)[reply]

"The tiny rodents are born to run. It comes naturally to them. Hamsters truly enjoy running, plain and simple, and their species as a whole are very energetic and lively. According to the ASPCA, hamsters in their natural habitat might run up to 5 miles every night in their quest for sustenance". [5] Couldn't find anything more acedemic, but there are pages of results giving the same answer. Alansplodge (talk) 13:48, 11 February 2018 (UTC)[reply]

Hamster wheel#Use by animals says in paragraph 1: Hypotheses to explain such high levels of running in wheels include.... The rest of the paragraph gives various hypotheses. Loraof (talk) 13:52, 11 February 2018 (UTC)[reply]

The most astonishing thing there is that wild mice will use wheels installed in the field. Clearly not doing it merely for exercise then. And not in any way connected with captivity. It must be pure enjoyment. As I said in the title, there is something wrong with them :) SpinningSpark 14:04, 11 February 2018 (UTC)[reply]

The wild mouse research is here: Wheel running in the wild Johanna H. Meijer, Yuri Robbers. Alansplodge (talk) 14:48, 11 February 2018 (UTC)[reply]

Often there is nothing wrong with hamsters [6], [7]. 86.169.57.217 (talk) 12:48, 17 February 2018 (UTC)[reply]

Is high prevalence of hypothyroidism caused by lack of food in our evolutionary history?

It has been estimated that about 10% of all women have thyroid hormone deficiency. The question is then how this situation could have arisen as a result of natural selection. Could it be that during famine conditions women would volunteer to eat less and those with (mild) thyroid conditions would then be more likely to survive? Count Iblis (talk) 22:04, 11 February 2018 (UTC)[reply]

It could be that as a typical Pakled you have mistaken the ruminations of a holodeck character with the basis for factual referenced reliably-sourced non-speculation, allowed by the policies of Wikipedia and the guidelines of this page. Indeed, I calculate the odds of that probabilitity at 97.36888... Do you have any other random thoughts, Jim, or may I return to my serach for dilithium before the local star goes nova? μηδείς (talk) 00:53, 12 February 2018 (UTC)[reply]

Thyroid Disorders at Midlife: An Evolutionary Perspective (2017) by Lynnette Leidy Sievert suggests that it may be an evolutionary response to "low environmental iodine or by increased iodine needs associated with pregnancy" rather than starvation. I'm not sure that I really understand what she's saying, but there's a list of citations, some of which are viewable online. Happy reading. Alansplodge (talk) 19:05, 12 February 2018 (UTC)[reply]

Thanks! Count Iblis (talk) 18:44, 14 February 2018 (UTC)[reply]