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.”

1 comment:

martin chaplin said...

Using a simulation where hydroxide and hydronium ions cannot recombine may lead to serious artifacts. Note that other workers (M. Boström, W. Kunz and B. W. Ninham, Hofmeister effects in surface tension of aqueous electrolyte solution, Langmuir 21 (2005) 2619-2623) suggest the opposite effect, i.e. that hydroxide ions prefer the surface.