gatewood
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Posts posted by gatewood
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holy ppl... im merely asking, what might you think happens to matter, as it compresses, beyond neutron degeneracy. E.g. what state of matter could it be said it is in... if at any? (like the hypothetical quark stars).
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2 hours ago, swansont said:
It's not clear that gravitational singularities exist, and we don't have the physics to describe such things.
"fundamental form of energy" doesn't make sense.
ohhh... please, cut the chase, we all know gravitational singularities are basically a placeholder for where our understanding of physics breaks down.
What I mean to say/ask: (fun question) what would you think we'll see if we could compress , say, 10 solar masses to the size of an atom (not an infinitesimally small volume)? Suppose there where no such thing as a schwartschild radius, but we could still compress stellar amounts of matter down to atomic/subatomic scales and also still observe it.
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2 hours ago, Phi for All said:
I don't know why so many people get this idea that we'll become or be superseded by an artificial superintelligence by this century.
We'll probably go on for thousands of years as the same, limited biological species in a civilization "a la star wars/trek".
Facepalm*
N/A
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I don't know why so many people get this idea that we'll go on for thousands of years as the same, limited biological species in a civilization "a la star wars/trek".
We'll probably become or be superseded by an artificial superintelligence by this century.
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Just a fun question: what state of energy/matter, could it be argued, that a gravitational singularity is in?
I would say that, it broke down to the most fundamental form of energy. Could it be said that, it is an extremely exotic form of atom?
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On 5/9/2021 at 6:19 AM, exchemist said:
I certainly can't - and I doubt anyone else can since it seems, from the link I gave you earlier in the thread, to depend on what borate product you are using. The DOT product in the link is said to give a more or less ideal pH on its own. At least, that is how I interpret this passage:
"When used in pools at typical dilution, the new formulation has an ideal pH of ≈7.6. Above pH 8, the chlorine becomes much less effective as a sanitizer, but below pH 7.4 the equilibrium trends toward hypochlorous acid and chlorine is lost more rapidly."
So, as I say, I think the best bet is to check with the supplier's recommendations for whatever borate product you are using - or have in mind to use.
Let me see if I understand. The measurement is simply, how acidic you want the pool water to be? That'll be easy, simply by calculating or experimenting, to get the precise ratio for the desired pH.
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On 6/13/2021 at 4:12 AM, chenbeier said:
The urea is more or less decomposed by the acid, there is nothing to purify.
It might depend, try to see the precise amount of aqua regia (nitric+hydrochloric acid) you spilled onto how many beads. If the ratio favors the beads, you probably have a mixture of urea and whatever product(s) resulted from the reaction.
Depending on how pure you need your urea and if the mass ratio heavily favors the urea, you could probably still do without separating it. If you need to / insist on separating the urea, it miiight be possible to separate using regular/fractional distillation. Investigate and do some balancing equations, to see what products you got from the reaction, to know boiling and volatilization points (see if you get any azeotropes).
Finally, do be careful, however, of the possible presence of urea nitrate, as it is explosive (though, if its there, you probably have too little of the compound, but still):
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16 hours ago, exchemist said:
Like all these things it gets more complex when you delve into it. Electron affinity is the energy released by an electronegative atom when it gains an extra electron and becomes an anion. For example, all the halogens release energy on gaining an electron, meaning the anion has lower energy than the neutral atom. The same is true for oxygen when it gains one electron. However when it gains a second, that is energetically unfavourable, due to the repulsion from the net -ve charge of the anion towards a second electron. (I had forgotten this, and only remembered after looking it up.)
However electron affinity is only a measure of the energy change when a free atom or ion gains an extra electron. In the case of metal oxides, the oxygen atom is not free. It is sitting in a crystal lattice, in the present case (CaO) surrounded by 6 nearest neighbour Ca2+ ions. That makes its environment much more energetically attractive for oxygen to pick up a second electron and form O2-. Hence it is common to find metal oxides with O2- anions even though, if the oxygen atom were free, you would have to "force" it to accept a second electron. (It's significant that if you put these oxides in contact with water you never get hydrated O2- ions. When they come out of the crystal lattice they pinch an H+ ion from water to make OH- (hydroxide) - plus another OH- from what is left of the water molecule:
O2- +H2O -> 2OH- .
As for the question about close proximity lowering potential, that's just replaying what you said, in effect, about magnets. You bring opposite magnetic poles together, or opposite electric changes together, and you lower the magnetic or electrostatic potential energy. That is reflected in the fact that you have to do work to pull them apart again. And, as they come together, magnets can gain kinetic energy, just as you said, at the expense of the magnetic potential energy. Similarly, ions with opposite charges approaching one another gain kinetic energy at the expense of electrostatic potential energy - which, in the context of molecular scale processes, means the heat energy given off in an exothermic reaction.
The alkaline earth (Group II) metal oxides react with CO2 in the air, yes. I'm less sure about the alkali (Group I) metals. The carbonate anion has a charge of 2- and this means it needs 2 M+ atoms to go with it, so I'm not sure how the kinetics and thermodynamics of that work out.
- First paragraph
Ok, then its just like stacking a bunch of N magnets onto a single S magnet, the atom loses the electron affinity, simply because its electric field becomes more balanced with the new electron. E.g. the Ca++ has higher energy than a Ca+ because its electric field is even more unbalanced, and an O- has more trouble accepting an electron than a neutral O, because... well, it is even more electronegative. Simple.
- Second paragraph (first half)
But wouldn't the Ca++ actually compete with the O-- for the electrons to reduce themselves?
- Second paragraph (second half)
You taught me something rather interesting, thank you very much
- Third paragraph
Sure, I realized too late, how easy that question was 🤪 . Thanks for taking the time to clarify anyhow
- Fourth paragraph
Yes, it does pick up 2 alkaline metals. E.g. 2Na+:
17 hours ago, studiot said:Look around you.
Water, alcohol, oil, resins, methane, nitrogen, oxygen, carbondioxide, benzene; These are common liquids or gases.
And they are all covalently bonded, albeit some have straightforward covalency some have dative covalency (ie are polar).Compare this with sodium chloride, iron oxide, copper sulphate; These are all common solids.
And they are all ionically bonded.This situation represents the very large majority of cases.
This is no accident, there are good reasons for this.That is to observe that ionic compounds tend to from solids whilst covalent compounds often appear as fluids.
Yes one of these reasons is molecular weight.
But there are plenty of examples of ionic solids with a lower molecular weight than covalent liquids, eg sodium chloride is 58 whilst benzene is 78, both from my list.So it is instructive to consider what is different.
The difference is that in a fluid the molecules have a degree of autonomy not present in a solid.
They can move about as a molecule.
And most important you can for instance identify one particular carbon atom with two particular oxygen atoms forming the 'molecule'.
However you cannot identify a particular sodium atom (ion) with one particular chlorine atom (ion) in the solid.
In fact electric forces link one (each) sodium+ to 6 chloride-
The coordination number is said to be 6.
So the intensity of the charge difference is distributed that way.https://courses.lumenlearning.com/cheminter/chapter/ionic-crystal-structures/
This is the key difference of importance to your question.
The chemical implications of these can be very complex indeed as both John Cuthber and exchemist are trying to tell you.Well, I'm aware of the melting and boiling point difference between covalent bonds, which have low melting and boiling points, because the energy that binds the molecules is low, and ionic ones, which bind their molecules far more strongly, that's why, huge temperatures are needed to even begin to have them behave as fluids.
And about the molecular weight, I was kinda aware of that, but it'll be a whole different topic to talk about more complex molecules such as aromatic compounds (I mean, cellulose is a huge polymer and it'll decompose way before even melting).
Finally, I'm definitely gonna study that last part, thanks for sharing it
18 hours ago, John Cuthber said:So, your point was that the cliffs are mainly made from a highly soluble thing that doesn't dissolve.
The dover white cliffs are an entirely different creature, they're made mostly of chalk (calcite minerals), product of ancient coccolithophores and other microorganisms, their shells are made of calcium carbonate (most sea shells are made of it), a metal carbonate that is NOT soluble (the very part of hardwood ashes that won't dissolve with water).
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Where the difference lies exactly?
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14 hours ago, John Cuthber said:
The White Cliffs of Dover indicate that the solubility isn't just down to the Ca++ ion.
The Ca++ ions in water are surrounded by a bunch of water molecules which are more or less firmly attached.
Well, that's precisely the point isn't it? A hydration sphere surrounding a cation (firmly attached due to the high electropositivity of the cation).
11 hours ago, exchemist said:Yes you are right, the answer I gave before was a simplistic one, for ionic compounds generally. The case of these oxides is a bit more involved, since in fact the second electron affinity of oxygen is +ve, endothermic.
The stability of compounds like CaO relies on a high lattice energy, i.e. the reduction in electrostatic potential that comes from close approach of a large number of oppositely charged ions, in a crystal lattice. In the case of CaO, the O2- anion and the Ca2+ cation are of very similar size, allowing a very efficient packing arrangement which minimises inter-ionic distances, releasing more energy, which compensates for the energy required to get a second extra electron onto the O atom.
So my first answer was a bit misleading. Sorry about that.
Its ok, I wasn't all too clear myself either.
Hmmm... interesting. Forgive my ignorance, but you gave me a lot of questions buddy:
1. Could you elaborate a bit more on your first paragraph, if you would? (that last part)
2. Why does the close proximity of the ions produce a reduced electrostatic potential?
3. Why exactly does the second electron (to complete the valence shell) requires an extra kick to orbit oxygen?
4. Aren't the alkaline and alkali oxides, actually rather reactive? If exposed to the atmosphere, they will form metal carbonates... if I understand it correctly.
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Thanks for the response.
That's interesting, since carbonic acid is, more or less, an aqueous solution of CO2, wouldn't it be possible to volatilize it, by aerating the water? Or by degassing it?
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Fair enough.
I have another question: what happens to a metal carbonate when it is hydrated? Does it split into its metal cation and carbonic acid? (like how sodium chloride splits into its respective ions?).
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Yes, you explained it neatly and pretty well, though I'm aware of this, given it's ionic bonding 101.
The core of my question, lies more in the fundamental physics that are going on, during the bonding of 2 ions. The way I think I understand ionic bonds, is akin to how, 2 magnets have high potential energy when they're far apart, which gets turned into kinetic energy when "falling" into each other, then all such energy is released and the system is now at a lower energy state, once the magnets come in contact and are at rest.
So given such model, it is not actually the transferring of electrons, which releases energy in itself, but the actual process of coming in contact. I mean, given how electropositive it is, wouldn't a +2 calcium cation quickly bind with basically anything with an inkling of electronegativity or negative polarity? (which is what I think, makes it so soluble in water, just like how sodium cations and chlorine anions break apart to bond with the negative and positive polarity ends of the H2O molecule, their electric fields have found more positive and negative fields on which to be balanced). The cation remains highly reactive till it finds an electronegative atom/molecule, on which to bond and balance the electric field.
Though, to be fair, the O2 molecule is probably a bad candidate, since it is non-polar (it would require some energy to break the covalent bond).
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Would it be possible to have, say, calcium cations in aqueous solution, react with O2 to form calcium oxide, if the concentration of the dissolved gas was high enough?
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On 6/9/2021 at 5:10 AM, John Cuthber said:
It is utterly impractical to consider getting K2O by heating the carbonate. It will not work.
If you want KOH there are ways that do work.
The same chemistry they discuss here
https://pubs.acs.org/doi/pdf/10.1021/ie51398a020
is probably the easiest way.That paper is very interesting (though the term "lime", I think should be replaced with "quicklime"). Most appreciated input kind sir.
Just wanted to say that, the thing I have about using a thermal process, is that it would be rather easy to set up a kiln with charcoal basically anywhere, with no need for reactants or electricity, and the process would be rather straightforward.
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On 6/10/2021 at 4:59 AM, John Cuthber said:
Do you understand the distinction between " we can not imagine" something and " we recognise the physical impossibility (or impracticability) of something?
Your idea seems to be a less practical version of this.
https://en.wikipedia.org/wiki/ElectrodialysisWell that's exactly the kind of feedback I'm looking for, most appreciated (that's why I'm here, to learn from the experts).
If you (I mean, not yourself, per se) only have: "well go try distillation or reverse osmosis" to say, when I'm proposing some brainstorm idea, then it seems, I have to recur to stronger criticism, to get the actually good answers (which I have, which I'm thankful for).
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On 6/9/2021 at 9:42 PM, Sensei said:
If somebody would like to remove water from solution he/she can:
- freeze it (ice forms on the top so you should remove it when it forms.. also solubility of salt depends on temperature therefore I told you to gather and learn properties of compounds that you have in solution and analysis of data),
https://en.m.wikipedia.org/wiki/Fractional_freezing
"Freeze distillation is a misnomer, because it is not distillation but rather a process of enriching a solution by partially freezing it and removing frozen material that is poorer in the dissolved material than is the liquid portion left behind."
- vacuum distillation.
https://en.m.wikipedia.org/wiki/Vacuum_distillation
"Vacuum distillation is often used in large industrial plants as an efficient way to remove salt from ocean water, in order to produce fresh water. This is known as desalination."
(you just have different salt than seasalt)
Having vacuum pump in the home lab is essential thing (apart glassware and stirrer with temperature control).
- evaporate excess of water.
https://en.m.wikipedia.org/wiki/Evaporation
You have to start listening and reading the all materials which are given to you..
If you would ask "how would you efficiently extract alkaline carbonates from solution?" people would give you their ideas. But you asked "how to filter alkaline carbonates from solution?" and upset because people answered "it won't work".. Ask the right questions in the first place..
The question is the correct one... that's exactly why I asked for "filtering" and not "separation", because it is generally MUCH cheaper to get a simple material to catch and remove the contaminant.
I know it sounds hard without using things like ion exchange or reverse osmosis (e.g. things like arsenic ions can be captured using iron oxide), that's why I thought you might have some clue.
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I get just how unimaginative you people are.
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2 hours ago, Sensei said:
OP should ask for methods of separation of chemical compounds. "filtering" is not surprising term for typical layman.
https://en.m.wikipedia.org/wiki/Separation_process
1) read Wikipedia articles of compounds in solution.
2) write down their properties.
3) pick up the one which is the most suitable, utilitizing differences in properties.
I know about separation techniques, only distillation and precipitation techniques I've found suitable, that's why I ask if anyone here has had experience with what I'm asking. Thanks for the "help".
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I'm looking for ways to filter (besides distillation) soluble alkaline carbonates, such as sodium and potassium carbonates from water. Do things like activated charcoal remove it?
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Now, for a bit of social and urban engineering trivia, here's Kowloon, the walled city of Hong Kong:
https://www.businessinsider.com/kowloon-walled-city-photos-2015-2?r=MX&IR=T
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Well, distilled water is not deionized water (some azeotropes carry on) and it is far more energy intensive. Also my hypothetical electrode technique wouldn't require regeneration (which can get expensive) and no resins.
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2 hours ago, SergUpstart said:
And those used in water ionizers. (The water ionizer is essentially an electrolyzer)
ok, can I ask you if you could elaborate on that... please? (about the whole cotton membrane)
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Well, yes you're right there, it would be hard to get there (to deionized water, which is what I want). How about this:
1. Thoroughly filter water with sand, HEPA and activated carbon filters.
2. Boil it a bit to remove dissolved gases.
This way, we reduce contaminants to a bare minimum, only ions remains.
3. Pass them through electronegative and electropositive graphite electrodes (this will trap ions)
If any reaction did take place, and i end up with things like sodium or potassium chlorate, then those can then be filtered with another round of activated charcoal.
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State of "matter" of a singularity
in Astronomy and Cosmology
Posted · Edited by gatewood
Sure, but the question doesn't need it.
Say we paused a black hole just after its schwarzschild radius gobbled all the core of the star that formed it, and, hypothetically, we could take a peek inside. What would we see? What would matter compressed down, further than neutron (or quark) degeneracy, would be like?
I mean, the core of my question would be: all that fell inside a black hole... still exists in some form? And if so, what you think it that form is? A Bunch of elementary particles? Energy resulting from annihilated particles? A bunch of photons and neutrinos?