Science News
from research organizations

The Secret Nature Of Hydrogen Bonds

January 21, 1999
American Institute Of Physics
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.

COLLEGE PARK, MD--A US-France-Canada physicscollaboration has unambiguously confirmed for the firsttime the controversial notion--first advanced in the1930s by famous chemist and Nobel Laureate LinusPauling--that the weak "hydrogen" bonds in waterpartially get their identity from stronger "covalent"bonds in the H2O molecule. As Pauling correctlysurmised, this property is a manifestation of the factthat electrons in water obey the bizarre laws ofquantum mechanics, the modern theory of matter andenergy at the atomic scale. Performed by researchersat Bell Labs-Lucent Technologies in the US, theEuropean Synchrotron Radiation Facility in France, andthe National Research Council of Canada, the experimentprovides important new details on water's microscopicproperties, which surprisingly remain largely unknownand difficult to measure. To be published in theJanuary 18 issue of the journal Physical ReviewLetters, these new details will not only allowresearchers to improve predictions involving water andhydrogen bonds, but may also advance seeminglyunrelated areas such as nanotechnology andsuperconductors.


One of the mostimportant components of life as we know it is thehydrogen bond. It occurs in many biological structures,such as DNA. But perhaps the simplest system in whichto learn about the hydrogen bond is water. In liquidwater and solid ice, the hydrogen bond is simply thechemical bond that exists between H2O molecules andkeeps them together. Although relatively feeble,hydrogen bonds are so plentiful in water that they playa large role in determining their properties.

Arising from the nature of the hydrogen bond and otherfactors, such as the disordered arrangement of hydrogenin water, the unusual properties of H2O have madeconditions favorable for life on Earth. For instance,it takes a relatively large amount of heat to raisewater temperature one degree. This enables the world'soceans to store enormous amounts of heat, producing amoderating effect on the world's climate, and it makesit more difficult for marine organisms to destabilizethe temperature of the ocean environment even as theirmetabolic processes produce copious amounts of wasteheat.

In addition, liquid water expands when cooled below 4degrees Celsius. This is unlike most liquids, whichexpand only when heated. This explains how ice cansculpt geological features over eons through theprocess of erosion. It also makes ice less dense thanliquid water, and enables ice to float on top of theliquid. This property allows ponds to freeze on the topand has offered a hospitable underwater location formany life forms to develop on this planet.


In water, there are twotypes of bonds. Hydrogen bonds are the bonds betweenwater molecules, while the much stronger "sigma" bondsare the bonds within a single water molecule. Sigmabonds are strongly "covalent," meaning that a pair ofelectrons is shared between atoms. Covalent bonds canonly be described by quantum mechanics, the moderntheory of matter and energy at the atomic scale. In acovalent bond, each electron does not really belong toa single atom--it belongs to both simultaneously, andhelps to fill each atom's outer "valence" shell ofelectrons, a situation which makes the bond verystable.


On theother hand, the much weaker hydrogen bonds that existbetween H2O molecules are principally the electricalattractions between a positively charged hydrogenatom--which readily gives up its electron in water--anda negatively charged oxygen atom--which receives theseelectrons--in a neighboring molecule. These"electrostatic interactions" can be explained perfectlyby classical, pre-20th century physics--specifically byCoulomb's law, named after the French engineer CharlesCoulomb, who formulated the law in the 18th century todescribe the attraction and repulsion between chargedparticles separated from each other by a distance.


After theadvent of quantum mechanics in the early 20th century,it became clear that this simple picture of thehydrogen bond had to change. In the 1930s, the famouschemist Linus Pauling first suggested that the hydrogenbonds between water molecules would also be affected bythe sigma bonds within the water molecules. In a sense,the hydrogen bonds would partially assume the identityof these bonds!

How do hydrogen bonds obtain their double identity? Theanswer lies with the electrons in the hydrogen bonds. Electrons, like all other objects in nature, naturallyseek their lowest-energy state. To do this, theyminimize their total energy, which includes theirenergy of motion (kinetic energy). Lowering anelectron's kinetic energy means reducing its velocity.A reduced velocity also means a reduced momentum. Andwhenever an object reduces its momentum, it must spreadout in space, according to a quantum mechanicalphenomenon known as the Heisenberg UncertaintyPrinciple. In fact, this "delocalization" effectoccurs for electrons in many other situations, not justin hydrogen bonds. Delocalization plays an importantrole in determining the behavior of superconductors andother electrically conducting materials at sufficientlylow temperatures.

Implicit in this quantum mechanical picture is that allobjects--even the most solid particles--can act likerippling waves under the right circumstances. Thesecircumstances exist in the water molecule, and theelectron waves on the sigma and hydrogen bonding sitesoverlap somewhat. Therefore, these electrons becomesomewhat indistinguishable and the hydrogen bondscannot be completely be described as electrostaticbonds. Instead, they take on some of the properties ofthe highly covalent sigma bonds--and vice versa. However, the extent to which hydrogen bonds were beingaffected by the sigma bonds has remained controversialand has never been directly tested by experiment--untilnow.


Workingat the European Synchrotron Radiation Facility (ESRF)in Grenoble, France, a US-France-Canada research teamdesigned an experiment that would settle this issueonce and for all. Taking advantage of the ultra-intensex-rays that could be produced at the facility, theystudied the "Compton scattering" that occurred when thex-ray photons ricocheted from ordinary ice.


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

Compton scattering is a very powerful technique,because it is one of the few experimental tools thatcan obtain direct information on the low-energy stateof an electron in an atom or molecule. By measuring theenergy lost by a photon and its direction as itscatters from a solid, one can determine the momentumit transfers to the electrons in a molecule--and learnabout the original momentum state of the electronitself. From this information, one can reconstruct theelectron's "ground-state wavefunction"--the completequantum-mechanical description of an electron in ahydrogen bond in its lowest-energy state.


The effect that theexperimenters were looking for--the overlapping of theelectron waves in the sigma and hydrogen bondingsites--was a very subtle one to detect. Rather thanstudy liquid water, in which the H2O molecules andtheir hydrogen bonds are pointing in all differentdirections at any given instant, the researchersdecided to study solid ice, in which the hydrogen bondsare pointing in only four different directions becausethe H2O molecules are frozen in a regularly repeatingpattern.

Still, the effect was expected to be fairly small--onlya tenth of all the electrons in ice are associated withthe hydrogen bond or sigma bond. The rest are electronswhich do not form bonds. What also complicates mattersis that Compton scattering records information on thecontributions from all the electrons in ice, not justthe ones in which the researchers were interested.

However, the experimenters had a couple of advantages.First, the ESRF is a latest-generation facility thatcan produce very intense beams of x-rayphotons--allowing the experimenters to obtain enoughCompton-scattering events to perform a meaningfulstatistical analysis that would allow them to uncoverthe effect in the data. Second, the researchers shinedthe x-rays from several different angles. Measuring thedifferences in the scattering intensity from thesedifferent angles allowed them to subtract outuninteresting contributions from nonparticipatingelectrons.


Takingthe differences in scattering intensity into account,and plotting the intensity of the scattered x raysagainst their momentum, the team observed wavelikefringes corresponding to interference between theelectrons on neighboring sigma and hydrogen bondingsites.

The presence of these fringes demonstrates thatelectrons in the hydrogen bond are quantum mechanicallyshared--covalent--just as Linus Pauling had predicted. The experiment was so sensitive that the team even sawcontributions from more distant bonding sites. Fromtheoretical analysis and experiment the team estimatesthat the hydrogen bond gets about 10% of its behaviorfrom a covalent sigma bond.


For many years, manyscientists dismissed the possibility that hydrogenbonds in water had significant covalent properties This fact can no longer be dismissed. The experimentprovides highly coveted details on water's microscopicproperties. Not only will it allow researchers in manyareas to improve theories of water and the manybiological structures such as DNA which possesshydrogen bonds. Improved information on the h-bond mayalso help us to assume better control of our materialworld. For example, it may allow nanotechnologists todesign more advanced self-assembling materials, many ofwhich rely heavily on hydrogen bonds to put themselvestogether properly. Meanwhile, researchers are hopingto apply their experimental technique to study numeroushydrogen-bond-free materials, such as superconductorsand switchable metal-insulator devices, in which onecan control the amount of quantum overlap betweenelectrons 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 ReviewLetters. For a copy of the article please contact BenStein at 301-209-3091 or

Story Source:

The above post is reprinted from materials provided by American Institute Of Physics. Note: Materials may be edited for content and length.

Cite This Page:

American Institute Of Physics. "The Secret Nature Of Hydrogen Bonds." ScienceDaily. ScienceDaily, 21 January 1999. <>.
American Institute Of Physics. (1999, January 21). The Secret Nature Of Hydrogen Bonds. ScienceDaily. Retrieved October 6, 2015 from
American Institute Of Physics. "The Secret Nature Of Hydrogen Bonds." ScienceDaily. (accessed October 6, 2015).

Share This Page: