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The Secret Nature of Hydrogen Bonds

Hydrogen bonds are the chemical bonds that exist between H2O molecules and keep them together. Recent experiments have obtained new insights on the hydrogen bond, by shining x-rays of one color (blue arrow) on an ice crystal and analyzing the color and direction of the x-rays (red arrow) that emerge from the ice.

WHAT IS THE NEWS?

• Confirming a controversal prediction first made by famous chemist Linus Pauling, a new experiment has shown that the bizarre rules of the quantum world cause the weak "hydrogen bonds" in water molecules to get part of their identity from stronger "covalent" bonds within H₂O.
• Because hydrogen bonds play a significant role in determining water's properties, such as its unusual ability to shrink when heated, this experiment is likely to shed light on the numerous mysteries associated with water, which brought about many of the conditions favorable for life on this planet.
• The information from this experiment may help improve the understanding of biological structures that contain hydrogen bonds, such as DNA. It may enable researchers to design better self-assembling materials, which rely heavily on hydrogen bonds. And researchers hope to apply the techniques used in this experiment to learn new things about certain non-hydrogen-bond-containing materials, such as superconductors.

INTRODUCTION

COLLEGE PARK, MD--January 12,1999--A US-France-Canada physics collaboration has unambiguously confirmed for the first time the controversial notion -- first advanced in the 1930s by famous chemist and Nobel Laureate Linus Pauling -- that the weak "hydrogen" bonds in water partially get their identity from stronger "covalent" bonds in the H₂O molecule. As Pauling correctly surmised, this property is a manifestation of the fact that electrons in water obey the bizarre laws of quantum mechanics, the modern theory of matter and energy at the atomic scale. Performed by researchers at Bell Labs-Lucent Technologies in the US, the European Synchrotron Radiation Facility in France, and the National Research Council of Canada, the experiment provides important new details on water's microscopic properties, which surprisingly remain largely unknown and difficult to measure. z be published in the January 18 issue of the journal Physical Review Letters, these new details will not only allow researchers to improve predictions involving water and hydrogen bonds, but may also advance seemingly unrelated areas such as nanotechnology and superconductors.

One of the most important components of life as we know it is the hydrogen bond. It occurs in many biological structures, such as DNA. But perhaps the simplest system in which to learn about the hydrogen bond is water. In liquid water and solid ice, the hydrogen bond is simply the chemical bond that exists between H₂O molecules and keeps them together. Although relatively feeble, hydrogen bonds are so plentiful in water that they play a large role in determining their properties.

Arising from the nature of the hydrogen bond and other factors, such as the disordered arrangement of hydrogen in water, the unusual properties of H₂O have made conditions favorable for life on Earth. For instance, it takes a relatively large amount of heat to raise water temperature one degree. This enables the world's oceans to store enormous amounts of heat, producing a moderating effect on the world's climate, and it makes it more difficult for marine organisms to destabilize the temperature of the ocean environment even as their metabolic processes produce copious amounts of waste heat.

In addition, liquid water expands when cooled below 4 degrees Celsius. This is unlike most liquids, which expand only when heated. This explains how ice can sculpt geological features over eons through the process of erosion. It also makes ice less dense than liquid water, and enables ice to float on top of the liquid. This property allows ponds to freeze on the top and has offered a hospitable underwater location for many life forms to develop on this planet.

In water, there are two types of bonds. Hydrogen bonds are the bonds between water molecules, while the much stronger "sigma" bonds are the bonds within a single water molecule. Sigma bonds are strongly "covalent," meaning that a pair of electrons is shared between atoms. Covalent bonds can only be described by quantum mechanics, the modern theory of matter and energy at the atomic scale. In a covalent bond, each electron does not really belong to a single atom--it belongs to both simultaneously, and helps to fill each atom's outer "valence" shell, a situation which makes the bond very stable.

On the other hand, the much weaker hydrogen bonds that exist between H₂O molecules are principally the electrical attractions between a positively charged hydrogen atom--which readily gives up its electron in water--and a negatively charged oxygen atom--which receives these electrons--in a neighboring molecule. These "electrostatic interactions" can be explained perfectly by classical, pre-20th century physics--specifically by Coulomb's law, named after the French engineer Charles Coulomb, who formulated the law in the 18th century to describe the attraction and repulsion between charged particles separated from each other by a distance.

After the advent of quantum mechanics in the early 20th century, it became clear that this simple picture of the hydrogen bond had to change. In the 1930s, the famous chemist Linus Pauling first suggested that the hydrogen bonds between water molecules would also be affected by the sigma bonds within the water molecules. In a sense, the hydrogen bonds would even partially assume the identity of these bonds!

How do hydrogen bonds obtain their double identity? The answer lies with the electrons in the hydrogen bonds. Electrons, like all other objects in nature, naturally seek their lowest-energy state. To do this, they minimize their total energy, which includes their energy of motion (kinetic energy). Lowering an electron's kinetic energy means reducing its velocity. A reduced velocity also means a reduced momentum. And whenever an object reduces its momentum, it must spread out in space, according to a quantum mechanical phenomenon known as the Heisenberg Uncertainty Principle. In fact, this "delocalization" effect occurs for electrons in many other situations, not just in hydrogen bonds. Delocalization plays an important role in determining the behavior of superconductors and other electrically conducting materials at sufficiently low temperatures.

Implicit in this quantum mechanical picture is that all objects--even the most solid particles--can act like rippling waves under the right circumstances. These circumstances exist in the water molecule, and the electron waves on the sigma and hydrogen bonding sites overlap somewhat. Therefore, these electrons become somewhat indistinguishable and the hydrogen bonds cannot be completely be described as electrostatic bonds. Instead, they take on some of the properties of the highly covalent sigma bonds--and vice versa. However, the extent to which hydrogen bonds were being affected by the sigma bonds has remained controversial and has never been directly tested by experiment--until now.

Working at the European Synchrotron Radiation Facility (ESRF) in Grenoble, France, a US-France-Canada research team designed an experiment that would settle this issue once and for all. Taking advantage of the ultra-intense x-rays that could be produced at the facility, they studied the "Compton scattering" that occurred when the x-ray photons ricocheted from ordinary ice.

Crystal structure of ordinary ice. Red balls give the position of oxygen and white balls give the position of hydrogen. (Figure courtesy Bell Labs/Lucent Technologies.)

Named after physicist Arthur Holly Compton, who won the Nobel Prize in 1927 for its discovery, Compton scattering occurs when a photon impinges upon a material containing electrons. The photon transfers some of its kinetic energy to the electrons, and emerges from the material with a different direction and lower energy . By studying the properties of many Compton-scattered photons, one can learn a great deal about the properties of the electrons in a material.

Compton scattering is a very powerful technique, because it is one of the few experimental tools that can obtain direct information on the low-energy state of an electron in an atom or molecule. By measuring the energy lost by a photon and its direction as it scatters from a solid, one can determine the momentum it transfers to the electrons in a molecule--and learn about the original momentum state of the electron itself. From this information, one can reconstruct the electron's ground-state wavefunction--the complete quantum-mechanical description of an electron in a hydrogen bond in its lowest-energy state.

The effect that the experimenters were looking for--the overlapping of the electron waves in the sigma and hydrogen bonding sites--was a very subtle one to detect. Rather than study liquid water, in which the H₂O molecules and their hydrogen bonds are pointing in all different directions at any given instant, the researchers decided to study solid ice, in which the hydrogen bonds are pointing in only four different directions because the H₂O molecules are frozen in a regularly repeating pattern.

Still, the effect was expected to be fairly small--only a tenth of all the electrons in ice are associated with the hydrogen bond or sigma bond. The rest are electrons which do not form bonds. What also complicates matters is that Compton scattering records information on the contributions from all the electrons in ice, not just the ones in which the researchers were interested.

However, the experimenters had a couple of advantages. First, the ESRF is a latest-generation facility that can produce very intense beams of x-ray photons--allowing the experimenters to obtain enough Compton-scattering events to perform a meaningful statistical analysis that would allow them to uncover the effect in the data. Second, the researchers shined the x-rays from several different angles. Measuring the differences in the scattering intensity from these different angles allowed them to subtract out uninteresting contributions from nonparticipating electrons.

Taking the differences in scattering intensity into account, and plotting the intensity of the scattered x rays against their momentum, the team recorded wavelike fringes corresponding to interference between the electrons on neighboring sigma and hydrogen bonding sites.

The presence of these fringes demonstrates that electrons in the hydrogen bond are quantum mechanically shared--covalent--just as Linus Pauling had predicted. The experiment was so sensitive that the team even saw contributions from more distant bonding sites. From theoretical analysis and experiment the team estimates that the hydrogen bond gets about 10% of its behavior from a covalent sigma bond.

For many years, many scientists dismissed the possibility that hydrogen bonds in water had significant covalent properties This fact can no longer be dismissed. The experiment provides highly coveted details on water's microscopic properties. Not only will it allow researchers in many areas to improve theories of water and the many biological structures such as DNA which possess hydrogen bonds. Improved information on the h-bond may also help us to assume better control of our material world. For example, it may allow nanotechnologists to design more advanced self-assembling materials, many of which rely heavily on hydrogen bonds to put themselves together properly. Meanwhile, researchers are hoping to apply their experimental technique to study numerous hydrogen-bond-free materials, such as superconductors and switchable metal-insulator devices, in which one can control the amount of quantum overlap between electrons in neighboring atomic sites.

This research is reported by E.D. Isaacs, A. Shukla, P.M. Platzman, D.R. Hamann, B. Barbiellini, and C.A. Tulk in the 18 January 1999 issue of Physical Review Letters.

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