Talk:High-energy phosphate

Is the delta G physiologic or standard conditions?

Misleading idea that breaking bonds releases energy
I am thinking of breaking this article up into sections so that the final paragraph on the misleading use of the term 'high-energy' is emphasized more. As this paragraph states, bond breaking almost always requires input of energy. The energy released is due to the net total hydrolysis reaction, not the 'bond-breaking' process. I would like to do this because I cannot count the number of students I have had whom were taught in biology that "breaking bonds releases energy." Does anyone have a problem with this approach? Sirsparksalot (talk) 14:42, 12 November 2012 (UTC)
 * Not me!! I've had major trouble even getting this point into the ATP article to the extent that it's there now (see the TALK section). Bond breaking ALWAYS requires energy, or it wouldn't be a bond. The opposite idea is one of the most pervasive falsehoods in science, akin to the idea that mass can disappear and be converted to energy (that's matter not mass). I blame Lippman's 1941 squiggle bond representation, that looks like a coiled spring. A nice bit of showmanship for his ideas, but a source of endless misconception as a result, for 70+ years. S  B Harris 19:30, 12 November 2012 (UTC)
 * I would check out the talk page but I can't seem to find the archives. In any event, I can certainly understand your frustration and am sorry that I wasn't here at the time to help your cause! As for my original comment, I was not thinking of providing an overhaul of the article, just creating a new subsection with the heading "Miconceptions about bond-breaking", or something to that effect, and using the paragraph presently at the end of the article. If you think it is best left as is then I won't change it. Maybe it's something we can keep in mind for the long term? On another note, I was hesitant to say that bond breaking always requires energy because there are some (albeit very few) well-documented examples of metastable bonds that will release energy when broken. However, such examples are exceedingly rare and only occur in very special circumstances. Suffice it to say that they are virtually non-existent in naturally-occuring phenomena, including phosphate bonds. However, the scientist in me refrains from using the word "always." Sirsparksalot (talk) 23:22, 25 November 2012 (UTC)
 * There are no archives; all this is in the regular ATP TALK page. Just start here and read downward. Horrifying. Do feel free to make changes and back me up. My problem was in being one guy in sea of people who had read sloppy texts. I have had a few other experiences like it on WP: at the articles on weight and heat for example. But if there are two of us, one can take on the Injuns while the other sleeps! Didn't even realize there were metastable chem bonds. One bond? I see no reason it can't happen in theory, but these things happen in timescales usually of 10^-8 sec, so metastable would mean "only" 10^-3 sec? In metastable nuclear isomers you can get kinetic stabilty from high spin forbiddenness (ala Ta-180m). BUt there's not much to hang you up in a single chemical bond. Electrons in a single bond soon do what they want. It takes big collections of atoms (preferrably solid) to stop them. But in those cases the lower energy state is a very complicated distance away from the higher energy state, and single bonds aren't involved. Do you have a best example? S  B Harris 01:21, 26 November 2012 (UTC)
 * Fun times! I now see why you are so hesitant to make edits. There is certainly a lot of confusion and mis-information out there. I think that the root of the problem is that chemistry is very rarely taught "one molecule at a time," but instead as "one reaction at a time." In principle, this isn't a flawed approach since single-molecule reactions are exceedingly rare in nature. However, people seem to forget that anything that happens to one compound is inherently coupled to another molecule. That being said, it did take me a lot of time to figure this out. In retrospect, if I didn't have a background as a gas-phase physical chemist, I'm not sure if I would have been exposed to the type of information to to figure it out at all.


 * As for your question about metastable chemical bonds, the are quite rare and are generally isolated to multiply-charged ions in the gas-phase. Some have lifetimes of 10^-3 sec (as you predicted) and I believe that some have been stored for as long as seconds. I can try to dig up some references if you would like. Because of the excess charge on the ion, there is long range repulsion involved in forming the bond. From long distances, one charged fragment is repelled from another like-charged fragment. At short distances, this repulsion is masked by what can be thought of as "normal" covalent bonding, which arises from the interaction of fragment electrons with the nuclear framework of the other fragment. While I know only very little about nuclear physics, I will go out on a limb and say that the bond metastability is similar in nature to the metastable nature of nuclei, which arises from the long-range Coulomb repulsion of nucleons combined with the short-range strong/weak nuclear attraction. For chemical bonds,the timescale is dictaed by entropic effects and tunnelling. Typical molecules have some degree of residual energy in them and unimolecular decay will be dictated by the amount of time necessary for energy to find its way through phase space into the right nuclear coordinate. Typically, molecular fragments are too large for tunelling to occur on a reasonable time scale, unless the molecule has sufficiently-high energy to be close to the top of the barrier. All this being said, once you add solvent to the mix, it is able to polarize and reduce the effects of the long-range electrostatic repulsion between the moleucle framents. This makes the molecule go back to a regime of normal, stable bonding (much like we would think of an H2 potential energy surface).


 * Unfortunately, all of this lends some confusion for the case of phosphate bonds. I will have to go back, dig-up, and review some references to be sure, but I'm almost 100% certain that phsphate, because of its highly-negatively-charged state, won't form a chemical bond in the gas phase. From what I can gather, a lot of people use this general idea of phosphate's Coulombic instability to justify the "breaking a bond releases energy" argument. What people fail to realize is that this is only true in the GAS PHASE. In fact, if they really thought about it, this wouldn't even help their case since, for phosphate in the gas-phase, there is no bond to be broken! The only reason that phosphate as we know it exists is is because of solvent screening effects!


 * Anyway, all that aside, phosphate bond breaking does not occur in a vacuum but in aqueous solution. This means that, in addition to all the quantum effects present in all of the bonds of all involved species, the differential solvation of all reactants and products MUST be included. In these cases, we have to treat it just like any other stable bond (more or less), e.g. H2. In other words, breaking a bond, even phosphate, requires energy. If it wasn't coupled to a hydrolysis reaction, it would be an endothermic process and life, so far as we know it, wouldn't exist.


 * Okay, I'll stop preaching now. Hopefully that answered your question about metastable bonds. I'll see if I can help out on the other talk page when I get some more time to find a good entry into the thread. If you think any of the above conversation helps your cause, let me know and I can paste some of it into other talk pages. Sirsparksalot (talk) 17:12, 8 December 2012 (UTC)

I think I'll paste this entire section into the other talk page, complete with header. It never hurts for doctors and biochemists to learn more basic chem. Here's a great example of the foolishness that happens when they get it second or third hand. I do not except myself from such errors. S B Harris 20:46, 8 December 2012 (UTC)

ATP in vacuum and the origin of life (organic evolution)
Thanks for the concise energetic picture. I had pictured ATP molecules flying around as charged molecular ions in molecular mass spectrographs and hadn't ever considered that the things might not be stable at all outside water.

ATP is not even very stable IN water, except at extraordinary concentrations, and with some Mg2+ to cut the charge-charge interactions in the polyphosphate.

This brings up the two worst problems in organic evolution (evolution of life to the one-cell stage). From one self-replicating prokaryote cell to us seems straightforward, but getting to one cell is not. The basic reason for this is that all the polymers of life, both proteins and nucleic acids, are not stable as polymers in water. They tend to break down into monomers. The idea that some "primordial soup" would make polymers like this from any reasonable concentration of monomers in a water-based primordial soup (some type of ocean or pool), just won't fly. It's just not going to happen no matter how long you have. So the Stanley Miller-Leslie Orgel experiments that make monomer-thingies from gases and energy are interesting, but they don't get past the next step, which is a worse one, of how to string these monomers up in to polymers of life.

Life does it, of course, by coupling all the polymerizations in proteins, DNA and RNA to ATP-hydrolysis. And then it makes ATP by the damnest irreproducibly-complex Rube Goldberg contraption involving a membrane (the second big problem), AND a proton gradient across the membrane, AND an ATP synthase which is a complex 8 unit protein even in prokaryotes. And it needs this membrane to keep the H+ gradient up-- no primordial soup can work this. So where is the chicken and where is the egg? With no ATP you can't make a protein like an ATP synthase, and without some high-energy compound like ATP you can't make polymers of any kind, protein or RNA. So you need to START with ATP, or something like it, and a lot of it. The synthesis of it has to be coupled to some energy-producing reaction, or else it is an input from space.

It's a problem so bad that some Russian suggests that a giant meteorite made of polyphosphate slammed into a little pond somewhere and provided all the polyphosphate pro-anhydric-polymeric potential for the putting-together of all the polymers of the first protolife. Sounds wonky to me, but it's still the best idea I've heard so far on this. Everybody else ignores what it is a totally fundamental problem. Okay, how ELSE are you going to make nucleotides? Put in nucleotides and it all goes off into RNA-world with no problem. But you have to INPUT these.

Panspermia helps a little, since it gives us some more billions of years for all this to happen someplace else. But we still run up against thet time limit of the Big Bang and the need to make at least some second-generation stars. Okay 5 billion more years, but we need a mechnaism.

How about another type of life entirely, which we don't see? I've been wondering if we might not even need some kind of semi-creationist scenario ("semi," as this doesn't require a God-with-no-origin). Semi-creation would posit that irreducibly complex RNA-protein life was created by some other type of non-RNA, non-protein life that itself used some kind of polymer system that isn't so nutty and unstable in water as P~P~P, so that it DID make it from scratch, 10 billion years ago. Perhaps with 5 billion extra years, they then they had time enough to go from non-life to one-cell to intelligence, to hyper-intelligence (though not omniscience/omnipotence), and then they made the first simple RNA-protein-membrane-ATP-synthase life, using ATP as the basic entropy-energy currency, and sent it out as seeds, ala panspermia. Leaving us all here, 5 billion years after THAT, scratching our water-filled proteinacious heads with muscles that use ATP to power the action, and looking very confused. S B Harris 01:36, 9 December 2012 (UTC)