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Liquid water  

Liquid water

Liquid Water

Liquid water has several key properties that define its unusual behavior.

 

< Water molecule

< Water hydrogen-bonding

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< Anomalies of water

< Supercooled water

 


'Water is the driver of Nature'

Leonardo da Vinci   

 

A single water molecule

 

A single water molecule

Liquid water is the most extraordinary substance, challenging its apparently simple molecular formula. On this page, I give some of its defining properties.

 

The water molecule is bent (see left), and has two hydrogen atoms and one oxygen atom. The hydrogen atoms are slightly positively charged whereas the oxygen atom has a negative charge.

 

The molecules can interact with each other in three ways, hydrogen-bonding, van der Waals interactions, and electrostatic interactions. Most importantly, they can form hydrogen bonds with each other. These are moderate forces that connect the oxygen atoms of neighboring water molecules by means of single hydrogen atoms donated by one of the water molecules. A water molecule can, therefore, form extensive chains (see below right).

 

A chain of water molecules; a water wire

A chain of water molecules

 

Such chains can form complex three-dimensional networks. In liquid water, the network is a 3-D
percolating infinite cluster. Each linkage in such networks only lasts a short time (~ ns), such that the network is in a continual structural flux. However, the connections between water molecules within the network transmit structural information throughout the liquid.

 

Hydrogen-bonding between neighboring water molecules, together with the high density of molecules present due to their small size, produces a great cohesive effect within liquid water that is responsible for water's liquid nature at ambient temperatures. In addition to this complex connected system, van der Waals interactions act locally around each water molecule to hold and orient unbound water molecules, and electrostatic effects act at all distances (if decaying with distance) around the molecules to orient all of the water molecules. The combination of these three interactions gives rise to a two-state structuring of high-density and low-density networks.

 

Mobile protons and electrons

Mobile protons and electrons

In the liquid state, the three atoms do not stay together. The hydrogen atoms are constantly exchanging between water molecules, due to protonation/deprotonation processes and transfer along the water wires. Both acids and bases catalyze this exchange and even when at its slowest (at pH 7), the average time for the atoms in an H2O molecule to stay together is only about a millisecond. As this brief period is, however, much longer than the timescales encountered during investigations into water's hydrogen-bonding or hydration properties, water is often treated erroneously as a permanent structure.

 

Unlike most materials, water dissociates,

2 H2O(aq) equilibrium arrows  H3O+(aq) + OH-(aq)

 

Water can support acid-base equilibria over an extensive range. As a solvent, water can both accept protons from an acid and donate protons to a base, so revealing its amphoteric character; a property of great importance to life as we know it.

 

A typical tetrahedral group of five water molecules. The

central H2O donates hydrogen bonds to the top and front

and accepts hydrogen bonds from the side waters.

A typical local group of five water molecules
In liquid water, all water molecules have at least one hydrogen bond to neighboring water molecules with effectively no free water molecules found under ambient conditions. Water consists primarily of a mixture of clusters of water molecules with different degrees of hydrogen-bonding existing in rapid and complex equilibria. Liquid water contains a mix of short, straight strong hydrogen bonds and long, weak, bent hydrogen bonds with many intermediate forms between these extremes. Recent advances support the hypothesis that this is equivalent to a simple equilibrium between extreme examples of these clusters; low-density clusters consisting of mostly tetrahedrally hydrogen-bonded water molecules (see left) and high-density clusters with many long, weak and bent hydrogen bonds plus many adjacent, non-hydrogen bonded, van der Waals water-water interactions. The unusual variations with temperature and pressure of many physical properties of liquid water (the anomalies of water) are most easily explained using this two-state model which is also supported by much experimental data.


Liquid water contains by far the densest hydrogen-bonding of any material with almost as many hydrogen bonds as there are covalent bonds. The hydrogen bonds can rapidly rearrange in response to changing conditions and environments (for example, solutes). The hydrogen-bonding patterns are apparently random in water but influence each other. For any water molecule chosen at random, there is an equal probability (50%) that the four hydrogen bonds (that is, the two hydrogen bond donors and the two hydrogen-bond acceptors) locate at any of the four sites around the oxygen. Water molecules surrounded by four hydrogen bonds tend to clump together, forming clusters, for both statistical and energetic reasons. Hydrogen-bonded chains (that is, O-H····O-H····O) are cooperative, with the breakage of the first bond being the hardest, but then the next one is weakened, and so on. Such cooperativity is a fundamental property of liquid water where hydrogen bonds are up to 250% stronger than the single hydrogen bond in the dimer, H2O····H-OH. A strong base at the end of a chain may strengthen the bonding further. The cooperative nature of the hydrogen bond means that acting as an acceptor enhances the water molecule serving as a donor.

 

However, there is an anticooperative aspect in so far as acting as a donor weakens the water molecule's capability to act as another donor, for example, O····H-O-H····O. It is clear therefore that a water molecule with two hydrogen bonds where it serves as both donor and acceptor is somewhat stabilized relative to one where it has just two donors or two acceptors. This is the reason why it is suspected that the first two hydrogen bonds (donor and acceptor) give rise to the strongest hydrogen bonds.

 

Notable amongst the physical properties of water are the opposite influences of hot and cold water, with behavioral changes being more accentuated at low temperatures where the properties of supercooled water often diverge notably from those of warmer water. As cold liquid water is heated individual molecules shrink, bulk water shrinks and becomes less easy to compress, its refractive index increases, the speed of sound within it increases, gasses become less soluble, it takes less energy to heat, and it conducts heat better. In contrast, as hot liquid water is heated it expands, it becomes easier to compress, its refractive index reduces, the speed of sound within it decreases, gasses become more soluble,  it takes more energy to heat and it is a poorer conductor of heat. With increasing pressure, individual water molecules expand, cold water molecules move faster but hot water molecules move slower. Hot water freezes faster than cold water and ice melts when compressed except at high pressures (> 1 GPa) when liquid water freezes when compressed.

 

The dielectric constant (permittivity) of liquid water is unusually high due to its high concentration of dipoles enhanced by the correlations of their dipole orientations. This endows water with excellent solvent properties for salts. Water is also an excellent solvent for hydrophilic solutes, such as low molecular weight alcohols. Hydrogen-bonding is responsible for this property as they help the solubilization of hydrophilic molecules with oxygen and nitrogen atoms. The solubilities change significantly with temperature, and these changes correlate with the hydrogen-bonding donation and acceptance changes in the water.

 

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This page was established in 2018 and last updated by Martin Chaplin on 10 November, 2018


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