View source

== Hydrogen bonds in water ==
[[Image:Liquid_water_hydrogen_bond.png|thumb|right|300px|Snapshot from a simulation of liquid [[water]]. The dashed blue lines from the molecule in the center of the picture represent hydrogen bonds.]]
The most ubiquitous, and perhaps simplest, example of a hydrogen bond is
found between [[water]] molecules.  In a discrete water molecule, water has two hydrogen atoms and one oxygen atom.  Two molecules of [[water]] can form a hydrogen bond between them; the simplest case, when only two molecules are present, is called the [[water dimer]] and is often used as a model system. When more molecules are present, as is the case in liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogens on two other water molecules. This can repeat so that every water molecule is H-bonded with up to four other molecules, as shown in the figure (two through its two lone pairs, and two through its two hydrogen atoms.) 

[[Liquid]] water's high [[Boiling Point|boiling point]] is due to the high number of hydrogen bonds each molecule can have relative to its low [[molecular mass]], not to mention the great strength of these hydrogen bonds. Realistically the water molecule has a very high boiling point, melting point and viscosity compared to other similar substances not conjoined by hydrogen bonds. The reasoning for these attributes is the inability to, or the difficulty in, breaking these bonds. Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is up to four. For example, hydrogen fluoride—which has three lone pairs on the F atom but only one H atom—can have a total of only two bonds ([[ammonia]] has the opposite problem: three hydrogen atoms but only one lone pair).

:H-F<sup>...</sup>H-F<sup>...</sup>H-F

The exact number of hydrogen bonds in which a molecule in liquid water participates fluctuates with time and depends on the temperature. From [[water model|TIP4P]] liquid water simulations at 25 °C, it was estimated that each water molecule  participates in an average of 3.59 hydrogen bonds. At 100 °C, this number decreases to 3.24 due to the increased molecular motion and decreased density, while at 0 °C, the average number of hydrogen bonds increases to 3.69.<ref>{{cite journal|author= W. L. Jorgensen and J. D. Madura| title=Temperature and size dependence for Monte Carlo simulations of TIP4P water| journal=[[Mol. Phys.]]|date=1985| volume=56 |issue=6 |pages=1381 |doi=10.1080/00268978500103111}}</ref> A more recent study found a much smaller number of hydrogen bonds: 2.357 at 25 °C.<ref>{{cite journal|author=Jan Zielkiewicz|title= Structural properties of water: Comparison of the SPC, SPCE, TIP4P, and TIP5P models of water| journal=[[J. Chem. Phys.]]|volume= 123|pages=104501|date=2005| doi=10.1063/1.2018637}}.</ref> The differences may be due to the use of a different method for defining and counting the hydrogen bonds.

Where the bond strengths are more equivalent, one might instead find the atoms of two interacting water molecules partitioned into two [[polyatomic ion]]s of opposite charge, specifically [[hydroxide]] (OH<sup>&minus;</sup>) and [[hydronium]] (H<sub>3</sub>O<sup>+</sup>). (Hydronium ions are also known as 'hydroxonium' ions.)

:H-O<sup>−</sup>  H<sub>3</sub>O<sup>+</sup>

Indeed, in pure water under conditions of [[standard temperature and pressure]], this latter formulation is applicable only rarely; on average about one in every 5.5 × 10<sup>8</sup> molecules gives up a proton to another water molecule, in accordance with the value of the [[dissociation constant]] for water under such conditions. It is a crucial part of the uniqueness of water.