Water has long been known to
exhibit many physical properties that distinguish it from other small
molecules of comparable mass. Chemists refer to these as the
"anomalous" properties of water, but they are by no means mysterious;
all are entirely predictable consequences of the way the size and
nuclear charge of the oxygen atom conspire to distort the electronic
charge clouds of the atoms of other elements when these are chemically
bonded to the oxygen.
A
covalent chemical bond consists of a pair of electrons shared between two atoms. In the water molecule H
2O,
the single electron of each H is shared with one of the six outer-shell
electrons of the oxygen, leaving four electrons which are organized
into two non-bonding pairs. Thus the oxygen atom is surrounded by four
electron pairs that would ordinarily tend to arrange themselves as far
from each other as possible in order to minimize repulsions between
these clouds of negative charge. This would ordinarly result in a
tetrahedral geometry in which the angle between electron pairs (and
therefore the H-O-H bond angle) is 109°. However, because the two
non-bonding pairs remain closer to the oxygen atom, these exert a
stronger repulsion against the two covalent bonding pairs, effectively
pushing the two hydrogen atoms closer together. The result is a
distorted tetrahedral arrangement in which the H—O—H angle is 104.5°.
Because molecules are smaller than light waves, they cannot be
observed directly, and must be "visualized" by alternative means. The
two computer-generated images of the H2O molecule shown
below come from calculations that model the electron distribution in
molecules. The outer envelopes show the effective "surface" of the
molecule as defined by the extent of the electron cloud
The H2O
molecule is electrically neutral, but the positive and negative charges
are not distributed uniformly. This is shown clearly by the gradation
in color from green to purple in the image at the above right, and in
the schematic diagram here. The electronic (negative) charge is
concentrated at the oxygen end of the molecule, owing partly to the
nonbonding electrons (solid blue circles), and to oxygen's high nuclear
charge which exerts stronger attractions on the electrons. This charge
displacement constitutes an electric dipole, represented by the arrow at the bottom; you can think of this dipole as the electrical "image" of a water molecule.
As
we all learned in school, opposite charges attract, so the
partially-positive hydrogen atom on one water molecule is
electrostatically attracted to the partially-negative oxygen on a
neighboring molecule. This process is called (somewhat misleadingly)
hydrogen bonding.
Notice that the hydrogen bond (shown by the dashed blue line) is
somewhat longer (117 pm) than the covalent O—H bond (99 pm). This means
that it is considerably weaker; it is so weak, in fact,that a given
hydrogen bond cannot survive for more than a tiny fraction of a second.
How chemists think about water
The nature of liquid water and how the H2O
molecules within it are organized and interact are questions that have
attracted the interest of chemists for many years. There is probably no
liquid that has received more intensive study, and there is now a huge
literature on this subject.
The following facts are well established:
- H2O molecules attract each other through the special type of dipole-dipole interaction known as hydrogen bonding
- a hydrogen-bonded cluster in which four H2
Os are located at the corners of an imaginary tetrahedron is an especially favorable (low-potential energy) configuration, but...
- the molecules undergo rapid thermal motions on a time scale of picoseconds (10
–12 second), so the lifetime of any specific clustered configuration will be fleetingly brief.
A
variety of techniques including infrared absorption, neutron
scattering, and nuclear magnetic resonance have been used to probe the
microscopic structure of water. The information garnered from these
experiments and from theoretical calculations has led to the
development of around twenty "models" that attempt to explain the
structure and behavior of water. More recently, computer simulations of
various kinds have been employed to explore how well these models are
able to predict the observed physical properties of water.
This
work has led to a gradual refinement of our views about the structure
of liquid water, but it has not produced any definitive answer. There
are several reasons for this, but the principal one is that the very
concept of "structure" (and of water "clusters") depends on both the
time frame and volume under consideration. Thus questions of the
following kinds are still open:
- How do you distinguish the members of a "cluster" from adjacent molecules that are not in that cluster?
- Since
individual hydrogen bonds are continually breaking and re-forming on a
picosecond time scale, do water clusters have any meaningful existence
over longer periods of time? In other words, clusters are transient,
whereas "structure" implies a molecular arrangement that is more
enduring. Can we then legitimately use the term "clusters" in
describing the structure of water?
- The possible locations of neighboring molecules around a given H2O
are limited by energetic and geometric considerations, thus giving rise
to a certain amount of "structure" within any small volume element. It
is not clear, however, to what extent these structures interact as the
size of the volume element is enlarged. And as mentioned above, to what
extent are these structures maintained for periods longer than a few
picoseconds?
The
view first developed in the 1950's that water is a collection of
"flickering clusters" of varying sizes (right) has gradually been
abandoned as being unable to account for many of the observed
properties of the liquid. The current thinking, influenced greatly by
molecular modeling simulations beginning in the 1980s, is that on a
very short time scale (less than a picosecond), water is more like a
"gel" consisting of a single, huge hydrogen-bonded cluster. On a 10-12-10-9
sec time scale, rotations and other thermal motions cause individual
hydrogen bonds to break and re-form in new configurations, inducing
ever-changing local discontinuities whose extent and influence depends
on the temperature and pressure. It is quite likely that over very
small volumes, localized (H2O)n polymeric
clusters may have a fleeting existence, and many theoretical
calculations have been made showing that some combinations are more
stable than others. While this might prolong their lifetimes, it does
not appear that they remain intact long enough to detect as directly
observable entities in ordinary bulk water at normal pressures.
Think of liquid water of as a seething mass of H2O molecules in which hydrogen-bonded clusters are continually forming, breaking apart, and re-forming.
Theoretical models suggest that the average cluster may encompass as many as 90 H2O
molecules at 0°C, so that very cold water can be thought of as a
collection of ever-changing ice-like structures. At 70° C, the average
cluster size is probably no greater than about 25.
Prof.
Martin Chaplin of the London South Bank University has reviewed much of
the existing literature on water clustering, and has recently proposed
an icosohedral clustering model in which twenty 14-molecule tetrahedral units form an icosohedron containing a total of 280 H
2O units. This model is consistent with X-ray diffraction data and is able to explain all of the unusual properties of water.
It
must be emphasized that no stable clustered unit or arrangement has
ever been isolated or identified in pure bulk liquid water. A 2006 report
suggests that a simple tetrahedral arrangement is the only long-range
structure that persists at time scales of a picosecond or beyond.
Water
clusters are of considerable interest as models for the study of water
and water surfaces, and many articles on them are published every year.
Some notable work reported in 2004 extended our view of water to the
femtosecond time scale. The principal finding was that 80 percent of
the water molecules are bound in chain-like fashion to only two other
molecules at room temperature, thus supporting the prevailing view of a
dynamically-changing, disordered water structure.
Liquid and solid water
Ice, like all solids, has a well-defined structure; each water molecule is surrounded by four neighboring H
2Os. two of these are hydrogen-bonded to the oxygen atom on the central H2O molecule, and each of the two hydrogen atoms is similarly bonded to another neighboring H2O.
The
hydrogen bonds are represented by the dashed lines in this
2-dimensional schematic diagram. In reality, the four bonds from each O
atom point toward the four corners of a tetrahedron centered on the O
atom. This basic assembly repeats itself in three dimensions to build
the ice crystal.
When ice melts to form liquid water,
the uniform three-dimensional tetrahedral organization of the solid
breaks down as thermal motions disrupt, distort, and occasionally break
hydrogen bonds. The methods used to determine the positions of
molecules in a solid do not work with liquids, so there is no
unambiguous way of determining the detailed structure of water. The
illustration here is probably typical of the arrangement of neighbors
around any particular H2O molecule, but very little is known
about the extent to which an arrangement like this gets propagated to
more distant molecules.
Below are three-dimensional
views of a typical local structure of liquid water (right) and of ice
(left). Notice the greater openness of the ice structure which is
necessary to ensure the strongest degree of hydrogen bonding in a
uniform, extended crystal lattice. The more crowded and jumbled
arrangement in liquid water can be sustained only by the greater amount
thermal energy available above the freezing point.
ice
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water |
The anomalous properties of water
