Water has unusually high melting
point Water has unusually high boiling
point Water has unusually high critical
point Solid water exists
in a wider variety of stable (and metastable) crystal and
amorphous structures than other materials The thermal conductivity
of ice reduces with increasing pressure The structure of liquid water changes at high pressure Supercooled water has two phases and a second critical point at about -91°C Liquid water is easily supercooled but glassified with difficulty Liquid water exists at very low temperatures and freezes on heating Liquid water may be easily superheated Hot water may freeze faster than cold
water; the Mpemba effect Warm water vibrates longer than cold
water
Partial phase diagram of water (H2O) showing the melting (M.Pt.), boiling (B.Pt.) and triple (T.Pt) points.a
The melting point of water is over 100 K higher than expected by extrapolation of the melting points of other Group 6A hydrides, here above shown compared with Group 4A hydrides. It is also much higher than O2 (54 K) or H2 (4 K). See also below for further comparisons.
In ice (Ih), all water molecules participate in four hydrogen bonds (two as donor and two as acceptor) and are held relatively static. In liquid water, some of the weaker hydrogen bonds must be broken to allow the molecules to move around. The large energy required for breaking these bonds must be supplied during the melting process and only a relatively minor amount of energy is reclaimed from the change in volume (PΔV = -0.166 J mol-1). The free energy change (ΔG=ΔH-TΔS, where ΔH=ΔU+PΔV) must be zero at the melting point. As temperature is increased, the amount of hydrogen bonding in liquid water decreases and its entropy increases. Melting will only occur when there is sufficient entropy change to provide the energy required for the bond breaking. The low entropy (high organization) of liquid water causes this melting point to be high.
Although ice is very difficult to superheat above its (equilibrium) melting point, tiny
amounts of ice (Ih) have been superheated
to 290 K (without melting) for very short periods (>250
ps) [954a] with the limit of superheating (>1 ns) established at about 330 K [954b].
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]
The boiling point of water is over 150 K higher than expected by extrapolation of the boiling points of other Group 6A hydrides, here shown compared with Group 4A hydrides. It is also much higher than O2 (90 K) or H2 (20 K). See also below for further comparisons.
There is considerable hydrogen bonding in liquid water resulting in high cohesion (water's cohesive energy density is 2.6 times that of methanol), which prevents water molecules from being easily released from the water's surface. Consequentially, the vapor pressure is reduced. As boiling cannot occur until this vapor pressure equals the external pressure, a higher temperature is required.
The pressure/temperature range of liquidity for water is much larger than for most other materials (for example, under ambient pressure the liquid range of water is 100°C
whereas for both H2S and H2Se it is
about 25°C. [ Anomalies page : Back to Top ]
The critical point of water is over 250 K higher than expected by extrapolation of the critical points of other Group 6A hydrides, here shown compared with Group 4A hydrides. For example, the critical point (647 K, 22.06 MPa 322 kg m-3) is far higher than ethanol (514 K, 6.14 MPa 276 kg m-3), which also hydrogen bonds (but in chains not 3-dimensional) and is much larger and more massive.
The critical point can only be reached when the interactions between the water molecules fall below a certain threshold level. Due to the strength and extent of the hydrogen bonding, much energy is needed to cause this reduction in molecular interaction and this requires higher temperatures. Even close to the critical point, a considerable number of hydrogen bonds remain, albeit bent, elongated and no longer tetrahedrally arranged [92].
The critical points (C.Pt.), boiling points (B.Pt.)
and melting points (M.Pt.) of the molecules isoelectronic
with water shows water to have higher values.
Ammonia
and hydrogen fluoride also have somewhat raised values
as they form molecular clustering, albeit with three donor
H-atoms and one lone pair acceptor group or one donor
H-atom and three lone pair acceptor groups, respectively;
giving a maximum of two hydrogen bonds per molecule, on average. Although
solid HF forms stronger hydrogen bonds, these form linear zigzag chains with no
rings or polygons and hence its three-dimensional structure
is weaker. The hydrogen bonds in solid NH3 can form three-dimensional arrangements but are distorted and weakened.
Water has two donor H-atoms and two lone pair acceptor
groups with close to tetrahedral angles giving the possibility
of four hydrogen bonds per molecule with little distortion.
[ Anomalies page : Back to Top ]
The ability for water to form extensive networks of hydrogen bonds
increases the number of solid phases possible.
The open structure of hexagonal ice (19.65
cm3 mol-1), which contains only about 7.5 cm3 mol-1 of water molecules, gives plenty of scope for different arrangements
of the water molecules as the structure is compressed. For comparison,
hydrogen sulfide has only four distinct solid phases [119]. [ Anomalies page : Back to Top ]
Hexagonal ice shows anomalous reduction in thermal conductivity
with increasing pressure (as do cubic ice and low-density amorphous ice but not high-density amorphous ice ), which
behavior is different from most crystals where thermal conductivity increases with increasing density. Low-density amorphous ice is the only glass to show this peculiar behavior. This anomaly is due to the pressure-induced bending of the hydrogen bonding decreasing the transverse sound velocity
[617]. [ Anomalies page : Back to Top ]
In a similar manner to the formation of the high density crystalline (ice-five and ice
seven) and amorphous (HDA) ice phases, it is likely that liquid water undergoes a significant
change in structure at high pressure (about 200 MPa for liquid water).
The pressure-viscosity, self-diffusion, compressibility and structural properties of water change above about 200 MPa. Other changes also occur around 200 MPa, such as the loss of the density maximum and the discontinuity in fast sound in liquid water. The explanation for all these effects is that there appears
to be an increase in interpenetration of hydrogen bonded networks
at about 200 MPa (at 290 K); interpenetration of hydrogen
bonded clusters being preferred over more extreme bending
or breaking of the hydrogen bonds. This structuring for liquid water at high pressures is consistent to that found by neutron scattering [1001] and indicates that liquid water structuring at high pressure has similarity to that of its high pressure ice phases [1254]. [ Anomalies page : Back to Top ]
As water is supercooled it converts mainly into its expanded form
(for example, ES)
at ambient pressures, which at low enough temperatures (< -38°C)
may result in it forming metastable low-density amorphous ice (LDA;
although normally it will form hexagonal ice at this temperature). If the pressure on LDA is increased above
about 200 MPa then LDA undergoes a 30% collapse forming metastable
high-density amorphous ice (HDA) but notably in a continuous process
without breaking the hydrogen bonds [394].
This phase change cannot continue to higher
temperatures (so creating a second critical
point, [45]) as neither
of these phases is stable in the presence of liquid water although
they may convert into their metastable supercooled liquid forms.
The presence of these low- and higher-density forms of liquid (supercooled) water leads to the breakdown of the Stokes-Einstein relationship in supercooled water [1040] occurring far above the glass-transition temperature, in contrast to many supercooled liquids where this behavior is found only at temperatures just above this transition [1040b]. [ Anomalies page : Back to Top ]
Water freezing is not simply the reverse of ice melting [1110]. Melting is a single step process that occurs at the melting point as ice is heated whereas freezing of liquid water on cooling involves ice crystal nucleation and crystal growth that generally is initiated a few degrees below the melting point even for pure water. Liquid water below its melting point is supercooled water. It may be expected that the directional hydrogen bonding capacity of water would reduce its tendency to supercool as it would encourage the regular structuring in cold liquid that may lead to a crystalline state. Liquid water, however, is easily supercooled down to about -25°C and with more difficulty down to about -38°C with further supercooling possible, in tiny droplets (~5 μm diameter), down to about -41°C under normal atmospheric pressure. Water, supercooled down to -37.5°C, is sustained in storm clouds and the condensed clouds formed by aircraft at high altitude. Rather strangely, at the limit of this supercooling (also known as the homogeneous freezing point) the water activity is always 0.305 lower than that of water melting at the same temperature [457]. Where salts or hydrophilic solutes are present, the homogeneous freezing point reduces about twice as much as the melting point [663].
Liquid water may be maximally supercooled to about -92°C and 210 MPa. It should be noted that bulk water never forms a glass as the glass transition temperature (Tg, = ~136 K) for water is far lower, relative to its melting point (Tm, 273 K), than expected; Tm/Tg ~ 2 rather than Tm/Tg ~ 1.3-1.5 as for more typical liquids. Thus supercooled bulk water (i.e. not affected by surfaces or solutes) always crystallizes before its temperature can be sufficiently lowered, whatever the cooling rate [558]. Water glass may only be produced by extremely rapid cooling (105 K s-1) of tiny volumes of water (<~100 μm diameter).
As water is cooled, the cluster equilibrium shifts towards the more open structure (e.g. ES ) with higher viscosity. In order for crystallization to occur at least 3 - 4 unit cells worth of water molecules have to come together in the correct orientation.b The formation of whole or part icosahedral clusters interferes with this process whilst not allowing cluster crystallization due to their five fold symmetry. Lowering the temperature further, which should encourage crystallization, is partially counteracted by the increase in icosahedral clustering. The presence of ES clusters is, in principle, in agreement with computer simulation studies requiring the presence of metastable states [216]. Methods that break the hydrogen bonding in these clusters, such as ultrasonics [296], cause the supercooled water to immediately freeze.
There is a recent comprehensive review of the properties of supercooled
water [569].
[ Anomalies page : Back to Top ]
Deeply supercooled liquid water can be produced from glassy amorphous ice between
-123°C and - 149°C [74]
and may coexist with cubic ice up to -63°C [137].
This behavior is particularly anomalous as the liquid
(deeply supercooled water) is a 'strong'
liquid (compared with supercooled water that is a 'fragile'
liquid [493]) that
changes to crystalline solid (cubic ice)
on increasing the temperature whilst keeping the pressure
constant. Such a 'strong'
to 'fragile'
(Arrhenius to non-Arrhenius) change in a liquid is not normal. Deeply supercooled water exists in the liquid state where it appears to be too cold to diffuse sufficiently quickly to crystallize noticeably. A possible explanation of this low-temperature-range liquid water
may be the formation of strands of icosahedral
structures. This model can also explain the high viscosity and strong (that is, low specific heat) liquid behavior of this
extremely supercooled water [215]. The unusual behavior of this liquid (that is, deeply supercooled water), by solidifying on heating, has been found with other liquids (for example, methyl cellulose and some cyclodextrin solutions [1026]).
[ Anomalies page : Back to Top ]
Liquid water can be easily superheated above its boiling point away from its surface with the atmosphere [1128, 1184]. This may be particularly important when heating foods and drinks in a microwave oven where explosive production of steam from the superheated water may cause severe injuries. Superheating is also causes the boiling point of water to vary, in much the same way as its freezing point, and of irregular boiling, that is, 'bumping' [1184]. Liquid water may be superheated to about +240°C to +280°C in capillaries or small droplets within high-boiling immiscible solvents. Superheating is also apparent at low temperatures but at negative pressures (i.e. stretched water). Water may be superheated by reducing the pressure to below -100 MPa at 20°C [1128]. Superheating is facilitated by dissolved gas that may increase its hydrogen-bonded order [821] but prevented by the presence of gas bubbles or nanobubbles (that is, cavities) that act as initiation sites for vaporization.
Water vapor (gas) may easily be cooled below its condensation temperature (dew point) for its partial pressure (i.e. its boiling point ) in the absence of dust, or other, particles or surfaces that help the nucleation process [1184].
An interesting,
if unrelated effect (the Leidenfrost effect), is that water
droplets remain far longer on a hotplate just above 200°C
than if the hotplate was just above 100°C. (see [960]
for an amusing scientific answer to how water boils). [ Anomalies page : Back to Top ]
The ability of hot water to freeze faster than cold seems counter-intuitive as it would seem that hot water must first become cold water and therefore the time required for this will always delay its freezing relative to cold water. However experiments show that hot water (for example, 90 °C) does often (but by no means always) appear to freeze faster than the same amount of cold water (for example, 18°C) under otherwise identical conditions [158]. This has been recognized even as far back as Aristotle in the 4th century BC but was brought to the attention of the scientific community by the perseverance of Erasto Mpemba [1388] a Tanzanian schoolboy,c who refused to reject his own evidence, or bow to disbelieving mockery, that he could freeze ice cream faster if he warmed it first. For a recent review of the Mpemba effect, see [959].
A number of explanations have been put forward. One that has gained some support is that there is sufficient evaporation from the hot water that this causes faster cooling plus a reduction in mass, so faster freezing [1390]. There does not seem to be sufficient mass lost in experiments to support such an explanation as the sole cause, however, and the effect is apparent even when the vessel has a lid. Related to this explanation, part of the cause may be due to the larger and more persistent convective currents created when hot water cools; however the effect is again evident even if the vessel has an insulated base and lid. The most likely scenario (described in [158], but disputed [1415].) is that the degree of supercooling is greater, under some circumstances, in initially-cold water than initially-hot water. The initially-hot water appears to freeze at a higher temperature (less supercooling) but less of the apparently frozen ice is solid and a considerable amount is trapped liquid water. Initially-cold water freezes at a lower temperature to a more completely solid ice with less included liquid water; the lower temperature causing intensive nucleation and a faster crystal growth rate. If the freezing temperature is kept about -6°C then the initially-hot water is most likely to (apparently) freeze first. If freezing is continued, initially-cold water always completely freezes before initially-hot water.
Why initially-cold water supercools more is explained in terms of the gas concentration and the clustering of water. Icosahedral clusters do not readily allow the necessary arrangement of water molecules to enable hexagonal ice crystal initiation; such clustering is the cause of the facile supercooling of water. Water that is initially-cold will have the maximum (equilibrium) concentration of such icosahedral clustering. Initially-hot water has lost much of its ordered clustering and, if the cooling time is sufficiently short, this will not be fully re-attained before freezing. Experiments on low-density water around macromolecules have shown that such clustering processes may take some time [4]. Also of relevance here is that the formation of clathrate ices, which have structures closely related to the icosahedral clusters, behaves in an opposite manner. Thus, the supercooling from hot water is far greater than that from cold water [1391].
It is also possible that dissolved gases may encourage supercooling
by (1) increasing the degree of structuring, by hydrophobic hydration,
in the previously-cold water relative to the gas-reduced previously-hot
water (the critical effect of low concentrations of dissolved gas
on water structure is reported in [294];
re-equilibration taking several days) and (2) increasing the pressure
as gas comes out of solution when the water starts to crystallize,
so lowering the melting point and reducing the tendency to freeze
(see guest book). Also, the
presence of tiny gas bubbles (cavities produced on heating) may increase the
rate of nucleation, so reducing supercooling [428]. Recently another possibility has been described depending on changes in dissolved material with temperature (such as the reduction in bicarbonate in heated 'hard' water), but this has not yet been experimentally tested [1014]. The rationale for the Mpemba effect in this case concerns differences in the solute concentration at the ice-liquid interface causing a localized lowering of the melting point [1014].
[ Anomalies page : Back to Top ]
It is expected that the lifetime of an excited molecular vibration should decrease as the temperature increases as the energy and likelihood of interactions with other molecules also both increase. For example, the lifetime of the excited liquid HCl stretch vibration decreases from 2.1 ns at 173 K to 1.0 ns at 248 K.
In liquid water, the excited OH-stretch vibration has a lifetime
of 0.26 ps at 298 K and this lifetime increases to 0.32 ps at 358 K [592].
The reason for this is due to the effects of the hydrogen-bonded
network. The OH-stretch vibration normally relaxes by transferring
energy to an overtone of the H-O-H bending vibration. However, as
the temperature increases the hydrogen bonds of water get weaker,
which leads to an increase of the frequency of the stretch vibration
and a decrease of the frequency of the bending vibration. As a result,
the overtone of the bending mode shifts out of resonance with the
stretching mode, thereby making the energy transfer less likely.
[ Anomalies page : Back to Top ]
a The surface temperature on Mars lies below the triple point of water and its atmospheric pressure is close to this value, such that no liquid water may be found there. [Back]
b Theoretical considerations concerning the ice nucleation site size gives estimates of 45,000 water molecules at -5°C down to 70 water molecules at -40°C [265]. Molecular dynamics studies show that these do not need to form a crystalline structure for crystallization to occur [347]. [Back]
c Erasto Mpemba is now enjoying his retirement from being Principal Game Officer in Tanzania. [Back]
Density anomalies (D1-D20) explanations
Material anomalies (M1-M12) explanations
Thermodynamic anomalies (T1-T11) explanations
Physical anomalies (F1-F9) explanations
Home | Site Index | The anomalies of water | Water: Introduction | The icosahedral water clusters | Top
This page was last updated by Martin Chaplin on 27 June, 2008