Hydrogen bonding in water: Introduction
Hydrogen bonding forms in liquid water as the hydrogen atoms of one water molecule are attracted towards the oxygen atom of a neighboring water molecule.
In a water molecule (H2O),
the oxygen nucleus with +8 charges attracts electrons
better than the hydrogen nucleus with its +1 charge. Hence,
the oxygen atom is partially negatively charged and the
hydrogen atom is partially positively charged. The hydrogen
atoms are not only covalently attached to their oxygen
atoms but also attracted towards other nearby oxygen atoms.
This attraction is the basis of the 'hydrogen' bonds.
The water hydrogen bond is a weak bond, never stronger than
about a twentieth of the strength of the O-H covalent bond.
It is strong enough, however, to be maintained during thermal
fluctuations at, and below, ambient temperatures. The attraction
of the O-H bonding electrons towards the oxygen atom leaves
a deficiency on the far side of the hydrogen atom relative
to the oxygen atom. The result is that the attractive force
between the O-H hydrogen and the O-atom of a nearby water
molecule is strongest when the three atoms are in a straight
line (that is, O-H····O)
and when the O-atoms are separated by about 0.28 nm.
Each water molecule can form two hydrogen bonds involving
their hydrogen atoms plus two further hydrogen bonds utilizing
the hydrogen atoms attached to neighboring water molecules.
These four hydrogen bonds optimally arrange themselves
tetrahedrally around each water molecule as found in ordinary
ice. In liquid water, thermal energy bends and stretches
and sometimes breaks these hydrogen bonds. However, the
'average' structure of a water molecule is similar to
this tetrahedral arrangement. The diagram opposite shows
such a typical 'average' cluster of five water molecules.
In ice this tetrahedral clustering is extensive, producing
its crystalline form. In liquid water, the tetrahedral
clustering is only locally found and reduces with increasing
temperature. However, hydrogen bonded chains still connect
liquid water molecules separated by large distances.
There is a balance between the strength of the hydrogen bonds
and the linearity that strong hydrogen bonds impose on the
local structure. The stronger the bonds, the more ordered
and static is the resultant structure. The energetic cost
of the disorder is proportional to the temperature, being
smaller at lower temperatures. This is why the structure of
liquid water is more ordered at low temperatures. This increase
in orderliness in water as the temperature is lowered is far
greater than in other liquids, due to the strength and preferred
direction of the hydrogen bonds, and is the primary reason
for water's rather unusual properties.