Tuesday, April 24, 2007

There's a whole lot of simulating going on

Another biological ‘water wire’ is reported by Guillaume Lamoureux at U. Penn. and colleagues [Biophys. J. doi:10.1529/biophysj.106.102756]. They have simulated that ammonium transporter AmtB of E. coli, which has a hydrophobic channel that is thought to pass ammonia but to exclude water and charged species. The simulations, however, show that water can get inside and form a three-molecule chain. How this affects the permeation of ammonia (and exclusion of ammonium) isn’t yet clear, but it appears that the mechanisms proposed so far may run into problems.

Ulf Ryde at Lund and colleagues have simulated water in the active-site cavities of four human cytochromes: P450, 2A6, 2C8 and 3A4 [Rydberg et al., J. Phys. Chem. B, doi:10.1021/jp070390c]. In contrast to the crystal structures, they find that all the cavities are filled with water. That in 2A6 is small and contains only two waters, but the others have big cavities with around 40-60 water molecules – a volume of 1500-2100 angstroms. In these big cavities, water is rapidly exchanged with the environment through three to six channels. Those in 2A6 remain bound there, although quite mobile.

L-alanine acts rather like a surfactant in a water droplet, according to the density-functional MD simulations of Ivan Degtyarenko at Helsinki University of Technology and colleagues [J. Phys. Chem. B 111, 4227; 2007]. The amino acid moves to the droplet surface with its methyl group exposed, and the primary hydration shell of (on average seven) ordered and rather rigid water molecules forms around the carboxylate and ammonium groups.

More on the hydration of urea comes from Hinonori Kokubo and Montgomery Pettitt [J. Phys. Chem. B doi:10.1021/jp067659x]. They suggest that, to put it crudely, urea passes almost unnoticed in water – certainly, there’s no evidence of it acting as a structure-breaker. Another nail in the coffin for this hoary old idea.

Monday, April 23, 2007

Acid on top

[Victoria Buch and coworkers have just published a paper in PNAS on the acidity of the water surface, which is likely to have biochemical consequences. Here is the pre-edited article that I’ve written on the result for the May issue of Chemistry World.]

Pure water has an acid skin. This striking notion has been confirmed by calculations and tests by an international team of scientists. Victoria Buch of the Hebrew University of Jerusalem in Israel and her coworkers have found that the pH of pure water falls from the perfectly neutral value of 7 within the liquid to just 4.8 or less – about as acidic as beer – where water meets air at the surface [V. Buch et al. PNAS online, doi:10.1073/pnas.0611285104].

The finding could be significant for a number of disciplines. In the atmosphere, many important chemical reactions between trace gases take place at the surface of water droplets in clouds. “pH is an essential factor for many of these reactions”, says Pavel Jungwirth, one of Buch’s collaborators at the Academy of Sciences of the Czech Republic in Prague. “The low pH could also affect the rates of carbon dioxide absorption at the ocean surface”, adds water specialist Richard Saykally of the University of California at Berkeley.

And in molecular biology the effect might be reproduced where water comes into contact with water-repelling (hydrophobic) parts of proteins, changing the acid-base chemistry that goes on at protein surfaces. “The effects of interface acidification on protein and membrane structure could be huge”, says theoretical chemist Gregory Voth at the University of Utah.

“I think that this is a very important paper”, says Saykally, who has previously found indirect evidence of the surface acidification. While the prediction was well-known, he says, “this is the first study to actually predict the pH of the liquid water surface.”

It has become clear in the past several years that the water surface may look very different from the bulk liquid. Saykally and Jungwirth have shown that dissolved ions can become either depleted or concentrated at the surface [e.g. P. B. Petersen & R. J. Saykally, Annu. Rev. Phys. Chem. 57, 333; 2006]. But the accumulation of protonated water molecules (H3O+ or hydronium, the basic component of acidity in water) at the air-water interface happens for unique reasons.

In 2004 Voth and his coworkers used quantum-mechanical computer simulations to show that hydronium prefers to sit at the water surface rather than deep inside the liquid [M. K. Petersen et al. J. Phys. Chem. B 108, 14804; 2004]. Whereas H2O molecules typically form four hydrogen bonds to their neighbours – two via the hydrogen atoms, and two via the electron lone pairs on oxygen – H3O+ can only form three. The three hydrogens can bind to water molecules, but the oxygen atom, where most of the positive charge resides, can’t any longer act as a good ‘acceptor’ for hydrogen bonds.

This makes it energetically unfavourable for the oxygen side of hydronium to be in water at all. So, Voth’s team said, hydronium acts somewhat like an amphiphile – a molecule with a water-soluble part and a hydrophobic part, like a soap molecule. The ions gather at the air-water interface with the hydrogens pointing downwards to make hydrogen bonds and the oxygen pointing up out of the liquid [S. S. Iyengar et al. Int. J. Mass Spectr. 241, 197; 2005].

In 2005 Saykally and his colleague Poul Petersen used a spectroscopic technique to adduce indirect evidence for this surface excess of acidic hydronium [P. B. Petersen & R. J. Saykally, J. Phys. Chem. B 109, 7976; 2005]. Buch and colleagues now have fresh evidence, albeit still somewhat indirect: they show that deuterated water D2O can swap hydrogens with H2O about 20 times faster at the surface of ice nanocrystals (which have a liquid-like surface) than deeper inside. This switch depends on the presence of hydronium ions.

The researchers also have new simulation data for the surface acidification, which enables them to estimate the pH increase. But Voth cautions against placing too much weight on this number as yet, because of the simplifications in the model. For example, “they have no possibility of having the hydronium and hydroxide recombine into a water molecule”, he says.

A key question now is whether the experimental evidence could be made more direct – “ice crystals are not that relevant to water.” says Voth. But this will be a challenge: “it is very hard to measure the top surface layer of a volatile liquid like water”, says Jungwirth. Saykally agrees with that. “We are thinking hard about how to do this,” he says, “but it’s not easy.”

Tuesday, April 3, 2007

Smooth folding

A paper in the forthcoming issue of PNAS (104, 6206; doi/10.1073/pnas.0605859104) by Peter Tieleman’s group at the University of Calgary looks at how to reconcile the energy-funnel picture of protein folding with the fact that there are likely to be significant enthalpic barriers to the folding process. They use simulations of the association of two polyalanine and two polyleucine alpha-helices to figure out whether enthalpic barriers exist (they do), and why. It seems they arise in this case from the fact that, as the chains approach, they must become desolvated before there is a compensating enthalpic gain from strong helix-helix interaction. (Interestingly, the researchers see no dewetting transition as the helices approach, of the sort predicted by Lum, Chandler and Weeks (J. Phys. Chem. B 103, 4570-4577; 1999), but only ‘steric dewetting’ when there is simply no longer space to fit in a layer of water. This contrasts with the simulations of hydrophobic plates by Bruce Berne mentioned in my previous post (JACS asap doi:10.1021/ja068305m), where dewetting does feature. It is possible that this might be rather sensitive to the precise geometry, size and hydrophobicity of the two surfaces.) The enthalpic energy barrier is, however, largely compensated by the gain in solvent entropy on desolvation, leading to a free-energy barrier that is very small (for poly-A) or non-existent (for poly-L). Thus, the idea of a relatively smooth free-energy funnel is recovered.

I’ve now had a better look at the Berne paper. It suggests that Hofmeister effects are indeed complicated, and not best explained by a simplistic structure-making/breaking model. In the case of the association of nanoscale hydrophobic surfaces, the effect of ions depends on whether or not they accumulate preferentially at the surfaces. High-charge-density ions induce salting-out (reducing the solubility of hydrophobes) via an entropic effect due to preferential exclusion of ions from the interfaces. Medium-charge-density ions induce salting in because of a different entropic effect, due to strong hydration of the ions and a consequent reduction in solvent entropy when the ions, preferentially adsorbed at the surfaces, are expelled as the surfaces associate (I think I’ve got that right). But low-charge-density ions cause salting in enthalpically, since they bind to the surfaces and lower the surface tension of the plate-water interface, thus lowering the enthalpy of association. As if this isn’t complicated enough, Berne and colleagues say that something quite different applied for electrolytes and small hydrophobic particles (see Zangi & Berne, J. Phys. Chem. B 110, 22736-22741; 2006). Hmm.

Not strictly related to biochemistry, but Rudy Marcus and Yousung Jung have just published a proposed explanation for why there is a rate acceleration of certain organic reactions, particularly those associated with ‘click chemistry’, at the interface of water and an organic phase: see JACS asap doi:10.1021/ja068120f.