Quick links ...................................... Water molecule hydrogen-bonding Phase diagram Water spectrum Dissociation Anomalies Water properties ...................................... Water clusters Protein hydration Hydrocolloids Ions ...................................... Table of Contents Site map

Explanation of the Material Anomalies of Water (M1-M18)

No aqueous solution is ideal
D₂O and T₂O differ significantly from H₂O in their physical properties
Liquid H₂O and D₂O differ significantly in their phase behavior
H₂O and D ₂O ices differ significantly in their quantum behavior
The mean kinetic energy of water's hydrogen atoms increases at low temperature
Solutes have varying effects on properties such as density and viscosity
The solubilities of nonpolar gases in water decrease with temperature to a minimum and then rise
The dielectric constant of water and ice are high
The dielectric constant shows a temperature maximum

The relative permittivity shows a 'kink' in its behavior with the temperature at 60 °C
The imaginary part of the dielectric constant shows a minimum near 20 K
Proton and hydroxide ion mobilities are anomalously fast in an electric field
The electrical conductivity of water increases to a maximum at about 230 °C
The electrical conductivity of water rises considerably with frequency
Acidity constants of weak acids show temperature minima
X-ray diffraction shows an unusually detailed structure
Under high pressure, water molecules move further away from each other with increasing pressure
Water adsorption may cause negative electrical resistance

M1    No aqueous solution is ideal

Ideality depends on the structure of the solvent being unaffected by the solute. Water is not even close to being a homogeneous phase at the molecular level. Local clustering will be affected by the presence of solutes, so changing the nature of the water. Even solutions of HDO in H₂O do not behave ideally. ^j Although most non-aqueous solutions also show deviations from ideality at higher concentrations, the deviations that occur in aqueous solutions are generally much more extensive. However, close to ideality in aqueous solutions can be achieved if the 'free' water is considered when some of the water molecules are bound more strongly to the solutes than they are to other 'free' water molecules [3505].

[  Anomalies page: Back to Top ]

M2    D₂O and T₂O differ significantly from H₂O in their physical properties

Different isotopic forms of compounds usually behave very similarly to each other. However, nuclear quantum effects in the water molecule are significant and differ between the isotopic forms. The 50% lighter mass of hydrogen compared with deuteron cause the properties of light water H₂O to be more influenced by nuclear quantum effects than those of heavy water D₂O. The heavier forms of water (D₂O where D = deuterium, 2.0141 g ˣ mol⁻¹; and T ₂O where T = tritium, 3.0160 g ˣ mol⁻¹) form stronger hydrogen bonds than light water (H₂O where H = protium, 1.0078 g mol⁻¹) and vibrate less. Hence, they are more ordered than ordinary water, as shown by their greater molar volumes, ^o are more tetrahedral and have more hydrogen bonds [1485]. This causes many of their properties (such as the viscosity, self-diffusion coefficient, protein solubility, protein stability [3562], toxicity ^a and biological activity [2265], including the effect on the frequency of circadian oscillations [1395]) to be different from those expected from a simple consideration of their increased mass (for example, the D₂O/H₂O viscosity ratio rises from about 1.16 at 100 °C to around 2.0 in deeply supercooled water [23b]. This difference appears as a shift in the equilibrium position equivalent to a slight increase in temperature [425]; for example, viscosity data has been reconciled if the temperatures are shifted by 6.498 °C and 8.766 °C for D₂O and T₂O, respectively [73]. ^b H₂O is about four-fold stronger as an acid than D₂O at 25 °C, and H₃O⁺ in H₂O is 1.5 times as strong an acid as D₃O⁺ in D₂O. Remarkably, the difference in the specific heat minimum between H₂O and D₂O is over 80 °C. Most of the differences between the behavior of H₂O and D₂O has been explained as due to the nuclear quantum effects ⁱ inherent in the large mass difference between the hydrogen and oxygen atoms [554. However, it has subsequently been found find that liquid D₂O is more structured than liquid H₂O. This work predicts the intra- and intermolecular structures of liquid T₂O, HDO and HTO [4448]. Although the electron densities of the different isotopic forms of liquid water have proved, so far, to be indistinguishable [566], it is expected that the O-D bond length is shorter than that of O-H due to its smaller asymmetric vibration and the smaller Bohr radius of D relative to H. This gives rise to small differences in the size and direction of the dipole moment between HDO and H ₂O [1174], which further confuses any analysis of the structure of water containing mixed hydrogen isotopes.

 

Even though the amount of deuterium in commonly-found water is low (≈ 16 mM), the properties of such water are different to water containing protium (¹H) only (deuterium-depleted water) [2063,3164]; as example, the proton spin-spin relaxation (T₂) is 0.347 ± 0.024 s for deuterium depleted water (D/H = 3 ppm) compared with 2.000 ± 0.140 s for 'normal' water (D/H = 140 ppm), and the diffusion coefficient is 37% greater in deuterium-depleted water) [3164].

 

Almost pure H₂O and D₂O exist, ^m but HDO can never be more than about 50% pure, existing only in the presence of both H₂O and D₂O. Mixtures of H₂O and D₂O rapidly equilibrate (via proton transfer by self-dissociation and recombination) to form HDO:

 

H₂O + D₂O   2HDO          K_eq = 3.86, 25 °C [609], ΔH_mix = 129.4 J mol⁻¹ [654]

 

which is close to a total randomization of the hydrogen atoms (that is, equal concentrations of HOH, HOD, DOH and DOD that would be K_eq = 4) but is reflected in a slight preference for the partitioning of the deuterium-containing species into the more extensive and stronger hydrogen-bonded clusters. ^l ΔH_mix is the enthalpy of mixing, which may be considered high. Thus a mixture of 20% D₂O and 80% H₂O (by volume and by mol; ≈ 11 mol D₂O per liter solution) gives a mixture of 64% H₂O, 32% HDO, and 4% D₂O (by mol, ≈ 36 M H₂O, ≈ 2 M D₂O and ≈ 18 M HDO) and lower the temperature of the mixture by over 0.5 °C (if no heat flows in from the environment). The K_eq decreases with decreased temperature [126a] and increased hydrogen-bond cooperativity [985]. Even the properties of HDO deviate from those expected from a consideration of the properties of H₂O and D₂O [126b], with the D-atom preferring to be hydrogen-bonded over the H-atom except where the hydrogen-bond is particularly short (as in H₅O₂⁺) [985]. Surprisingly, D₂O does not mix as readily with H₂O as might be pre-supposed [1472], presumably as the preferred D₂O clustering holds the D₂O molecules together. This tendency gives anomalous cooling curves in unstirred D₂O/H₂O mixtures [1472], slow equilibration, and unexpected entropic effects [1855]. Mixing 1:1 H₂O and D₂O gives rise to a small heat of mixing (≈ 1.3 kJ ˣ mol⁻¹).

 

The vibrational spectrum of HDO is fundamentally different from either H₂O or D₂O due to the separation of the two hydroxyl (O-H and O-D) vibrations in HDO but their combined motion in H₂O and D₂O. In HDO, the H atom is more reactive and more easily dissociated than the D atom. As hydrogen-bonding is a property of at least two water molecules, isotopic mixtures contain many differently paired (and more extensive) species, each of which may present different properties to those in natural liquid water. Care must be taken over extrapolating the properties of H₂O/D₂O mixtures (often used in neutron scattering and vibrational spectroscopic studies) to those of normal liquid water (that is, 99.97% H₂O). For example, D₂O is preferentially found at hydrophilic interfaces [1342].

 

Liquid T₂O is corrosive due to self-radiolysis (). The β particles travel only about 6 μm in water and even dilute solutions of HTO produce gaseous hydrogen (including HT) and redox-active products, including highly reactive hydroxyl radicals (·OH and ·OT).

 

Even H₂¹⁸O behaves differently from H₂¹⁶O due to reduced quantum translational motions, reducing the size of the first shell local hydrogen-bonded tetrahedron but leaving the non-bonded water distances almost the same [1035]. Although D₂O has a similar mass (only 0.04% heavier than H₂¹⁸O), its behavior is much more affected by the isotopic substitution due to the altered mass distribution influencing its librations and hence the local environment of both the first and second aqueous shells [1035]. H₂¹⁷O and HDO behave as different molecules to H₂¹⁶O in mixtures with H₂¹⁶O as shown by their colligative properties, the freezing points rising with the molality of H₂¹⁷O or HDO (freezing points rising here as they form mixed solid solutions) [1470].

 

At 50 °C, the vapor pressure of H₂O is 10% greater than D₂O, so the concentration of deuterium in the liquid phase of water increases on evaporation. However, at about 220 °C both vapor pressures are equal and between 220 °C, and its critical point, D₂O has a higher vapor pressure than H₂O. Both light (H₂O) and heavy water (D₂O) are used as moderators in nuclear reactors; reducing the speed of fast neutrons and turning them into thermal neutrons capable of sustaining a nuclear chain reaction. Moderation is more rapid using light water but more efficient using heavy water, as light water scatters and absorbs significantly greater than heavy water. Although for many years it was thought that light (H₂O) and heavy water (D₂O) apparently could not be distinguished by taste, it has been found that heavy water elicits a sweet taste, mediated by the human Taste Receptor type 1 member 2/Taste receptor type 1 member 3 (TAS1R2/TAS1R3, sweetness receptor) taste receptor [3989]. Large amounts of heavy water may be toxic due to reaction kinetic effects [3567].

[  Anomalies page: Back to Top ]

M3    Liquid H₂O and D ₂O differ significantly in their phase behavior.

The phase behavior of liquid H₂O and D₂O differ, with the triple point of D₂O being 3.79 °C and 48.439 Pa higher than that of H₂O, their vapor pressure curves crossing at 221 °C, and the critical point of D₂O being 3.249 °C and 393 kPa lower [1007]. This isotope effect originates in the reduced zero-point vibration ^p of D₂O that reduces its van der Waals volume (by about 1%) and its associated repulsive effect within the hydrogen bonds at lower temperatures, increasing the D₂O-D₂O hydrogen-bond strength. ^c At higher temperatures, the transition to the excited state is more easily accomplished in D₂O (≈ 2450 cm⁻¹, relative to H₂O ≈ 3280 cm⁻¹). Due to the asymmetry of the vibration, this increases D₂O's effective van der Waals volume and reverses the relative repulsive effect, reducing the D₂O-D₂O hydrogen bond strength at higher temperatures. ^d

 

The miscibility properties of liquid H₂O and D₂O also differ, with 3-methylpyridine being always miscible in liquid H₂O but immiscible in liquid D₂O between 38.5 °C and 117 °C at normal pressure [1976].

[  Anomalies page: Back to Top ]

M4    H₂O and D ₂O ices differ significantly in their quantum behavior (nuclear isotope effect).

Replacing hydrogen (¹H) by deuterium (²H) results in a greater volume (≈ 0.1%) of hexagonal ice (the volume isotope anomaly; the nuclear isotope effect). Thus, H₂O has a larger molar density (50.899 mol ˣ kg⁻¹) than D₂O (50.816 mol ˣ kg⁻¹) at their melting points. This is in contrast to the expected behavior (e.g., solid ²⁰Ne is 0.6% more voluminous than ²²Ne) due to the expected greater nuclear motion of lighter isotopes [1819]. Furthermore, the anomalous effect in water increases with temperature, even though a typical isotope shift should decrease with temperature [1819]. This anomalous behavior is due to the significantly greater covalent O-H asymmetric vibration and the O-H distance being longer than the O-D distance. Also, the hydrogen-bond donor-acceptor (oxygen–oxygen; H-O-H····OH₂) distance changes upon isotopic substitution (the Ubbelöhde effect also called the secondary isotope effect) [3528].

 

A similar anomaly occurs with ice VIII under pressure (see right) but not with the ionic ice X [2936]. Ices Ic, II, III, V, IX, XII, XIII, and XIV are also likely to have this volume isotope anomaly [2936, 3152]. Unexpectedly, while the ice II to ice Ic transition is endothermic for H₂O (+40 J ˣ mol⁻¹), it is exothermic for D₂O (−140 J ˣ mol⁻¹) [4132].

 

As the K_eq decreases with decreased temperature [126a] and increased hydrogen-bond cooperativity [985] (see above), at temperatures close to 0 K, this may mean that H₂O and D₂O may form separate phases and are no longer in equilibrium [985].

[  Anomalies page: Back to Top ]

M5    The mean kinetic energy of water's hydrogen atoms increases at low temperature (disputed)

The mean kinetic energy of liquid water's hydrogen atoms has been reported to reach a maximum at about 7 °C, then reducing on cooling to a minimum value at about 0 °C [1596]. However, it should be noted that a simulation does not show this anomaly [1760]. Also, new experimental results, benchmarked with those obtained for ice [2768], do not show this anomaly. At the current time, this anomaly should be considered suspect.

 

[  Anomalies page: Back to Top ]

M6    Solutes have varying effects on properties such as density and viscosity

Solutes will interfere with the cluster equilibrium by favoring either open or collapsed structures. Any effect will cause the physical properties of the solution, such as density or viscosity, to change. Solutes have a lower than expected impact on both the cryoscopic constant (that is, an effect of solute on freezing point depression, 1.86 K kg mol⁻¹, compare CCl₄ 30 K kg mol⁻¹) and ebullioscopic constant due to water's low molar mass and high heats of fusion and evaporation respectively.

[  Anomalies page: Back to Top ]

M7    The solubilities of nonpolar gases in water decrease with increasing temperature to a minimum and then rise. ^e

nonpolar gases are poorly soluble in water. Most gaseous solutes dissolve more in most solvents as the temperature is raised. However, nonpolar gases are much more soluble in water at low temperatures than would be expected from their solubility behavior at high temperatures.

 

The solubilities of the noble gases are shown opposite [IAPWS, 1166] and are tabulated below. Their hydration may be considered as the sum of two processes: (A) the endothermic opening of a clathrate pocket in the water and (B) the exothermic placement of a molecule in that pocket, due to the multiple van der Waals dispersion interactions (for example, krypton dissolved in water is surrounded by a clathrate cage with 20 Kr···OH₂ such interactions [1357]). In water at low temperatures, the energy required by process (A) is minimal as such pockets may be easily formed within the water clustering (by CS ES).^f Using the noble gases to investigate the solvation of nonpolar gases is helpful as they are spherically symmetrical and have low polarizability. In contrast, the shape and polarizability may confuse the hydration of other gases. The solubility of the noble gases increases considerably as the temperature is lowered. Their enthalpy and entropy of hydration become more negative as their fit into the water dodecahedral clathrate improves.

 

Solubility properties of the noble gases

Property		He	Ne	Ar	Kr	Xe	Rn
Atomic number		2	10	18	36	54	86
Atomic radius, Å [1167]		1.08	1.21	1.64	1.78	1.96	2.11
Partial molar volume in H2O [3763], cm3 ˣ mol−1		24.6	24.6	31.5	37.4	45.0	49.4
ΔG° of solution in H2O at 25 °C, kJ ˣ mol−1 [1296]		29.41	29.03	26.25	24.80	23.42	19.29 *
ΔH° of solution in H2O at 25 °C, kJ ˣ mol−1 [1296]		-0.59	-3.80	-11.98	-15.29	-18.99	-21.37 *
ΔS° of solution in H2O at 25 °C, J ˣ mol−1 ˣ K−1 [1296]		-100.6	-110.1	-128.2	-134.5	-142.2	-144.0 *
ΔCP of solution in H2O at 25 °C, J ˣ mol−1 ˣ K−1 [678]		105.2	150.8	178.3	204.4	197.5	292.8
Solubility, mM, 5 °C, 101,325 Pa [1166]	H2O	0.41	0.53	2.11	4.20	8.21	18.83
	D2O	0.49	0.61	2.38	4.61	8.91	20.41
Solubility minima, °C [IAPWS, 678]	H2O	30	50	90	108	110	91
	D2O	53	53	98	108	116

 

Oxygen (O₂) and nitrogen (N₂) molecules behave similarly (solubility minima of N₂ at 74 °C and O₂ at 94 °C, IAPWS), although their solubilities are low (O₂, 1.92 mM in H₂O, 2.14 mM in D₂O; N₂, 0.94 in H₂O, 1.05 mM in D₂O; all at 5 °C, 101,325 Pa [1168]). The higher solubility of O₂ over N₂, despite its lesser clathrate forming ability [1168] has been proposed due to its formation of weak hydrogen bonds to water [1168].^g

 

Therefore, the solubilization process is exothermic (that is, has negative ΔH) and (as predicted by Le Chatelier's principle) solubility decreases with temperature rise. At high temperatures (often requiring high pressure), the natural clustering is much reduced, causing greater energy to be required for opening the pocket in the water. The solubilization process becomes endothermic, and (as predicted by Le Chatelier's principle) solubility goes through a minimum before increasing with temperature rise (being fully miscible under supercritical conditions).

 

The more attractive the solute-water van der Waals dispersion interactions (due both to atomic number dependence and goodness of fit within the clathrate pocket), the greater the inherent exothermic nature of the process and therefore, the higher the temperature minimum (see Table above) and the greater the temperature range of negative temperature solubility coefficient. Similarly, Henry's constants ( = partial pressure/mole fraction; represents volatility, see opposite) exhibit increasing maxima with increasing size (the maxima are the same as the solubility minima above).

 

The poor solubility of nonpolar gases in water, despite their negative enthalpy change on dissolution, is due to the positive free energy changes (+ve ΔG) attributed to the anomalously large negative entropy changes (-ve TΔS). These entropy changes are caused by the structural enhancement of the water (ES) clusters. This conclusion is reinforced by the enhanced heat capacity (ΔC_P) of these solutions (+ve C_p, characteristic of a decrease in the degrees of freedom of the water solvent). This structural enhancement may include fixing the cluster centers, preventing the randomizing flickering between clusters otherwise evident, and ordering the inner dodecahedral water shells surrounding the solute molecules. There is also a decrease in volume (-ve ΔV) showing a reduction in the unoccupied space within the water solvent and also indicative of the gases occupying the pre-existing, if collapsed, clathrate sites.

 

Very large hydrophobic molecules and structures, such as membranes and surfactants (and oils and large hydrocarbons) are not soluble in water, which allows for the organelles that we need to make up our cells.

 

Counterintuitively despite forming stronger hydrogen bonds, D₂O is a better solvent than H₂O for nonpolar gases, it is a more static molecule and more easily forms the ES water clustering. Therefore D₂O can accommodate the guest molecules more easily without breaking its hydrogen bonds [874]. The addition of positively hydrating salts (for example, LiCl) that destroy the water low-density ES clustering reduces the solubility ('salt out') of the nonpolar gases. In contrast, hydrophobic hydrating salts (for example, tetramethylammonium chloride) that increase water low-density ES clustering stability also increase nonpolar gas solubility ('salt in'). Small nonpolar organic molecules also behave similarly to nonpolar gases, but their increased size alters the clathrate structuring. Thus benzene has a solubility minimum, at a lower temperature than expected from above, at about 20 °C [210].

 

Interestingly, the change in solubility of nonpolar gases with respect to their diameters has a maximum (and their free energy of hydration has a minimum) when diameters are about the same as that of the dodecahedral cavities (that is, ≈ 4.5 Å) in the icosahedral network [196]. The solubility behavior of larger hydrophobic molecules is discussed briefly elsewhere. It should also be noted that the solvent properties of superheated liquid water also change with temperature as water's dielectric permittivity reduces towards that of common organic solvents as the temperature rises towards its critical point. Even though the amount of air (that is, N₂ + O₂ + Ar) dissolved in water is very low, it is sufficient to lower the density of water by almost 5 ppm (that is, 0.0005%) at 0 °C [870].

 

It should be clear from the above discussion that the solubility of nonpolar gases (at equilibrium), in water at its point is not zero, an error propagated by some textbooks. However, the non-equilibrium production of steam would lead to the temporary loss of dissolved gas from the liquid phase, and this would increase the pressure in the closed system required for solubility determinations.

 

The solubility of gases (and other solutes such as salts) in ice is very low and this explains the usefulness of freeze-thaw operations under reduced pressure for degassing water.

[  Anomalies page: Back to Top ]

M8   The dielectric constant of water and ice are high (78.4 at 25 °C)

The high relative permittivity (dielectric constant) of water and its small molecular size combine to endow water with the strong ability to dissolve salts. Polar molecules, where the centers of positive and negative charges are separated, possess a dipole moment. This means that in an applied electric field, polar molecules tend to align themselves with the field. Although water is a polar molecule, its hydrogen-bonded network tends to oppose this alignment. The degree to which a substance does this is called its dielectric constant (permittivity). Because water possesses a hydrogen-bonded network that transmits polarity shifts extensively through rapid and linked collective changes in the orientation of its hydrogen bonds, it has a high relative permittivity (dielectric constant). This allows it to act as a solvent for ionic compounds, where the attractive electric field between the oppositely charged ions is reduced by about 80-fold, allowing the thermal motion to separate the ions amongst the solution. An ab initio molecular dynamics study showed that the increase of the average dipole moment (due to hydrogen-bonding) of water molecules and the local alignment of these molecular dipoles both contribute almost equally to this high relative permittivity (dielectric constant) [1672]. On cooling, as the water network strengthens and water's dipole moment increases, the dielectric of liquid water climbs to 87.9 (0 °C), increasing on conversion to ice to 96 (0 °C) then rising further to about 183 as the ice is cooled to 140 K. The increase in relative permittivity on freezing is anomalous as the relative permittivities of most polar liquids fall abruptly on freezing (e.g., nitrobenzene 37 3) to close to the square of the refractive index due to molecular rotations ceasing. The freezing of water is different due to the relative permittivity being mainly due to proton hopping rather than molecular rotation.

 

The relative permittivity (dielectric constant) drops on heating, and liquid water becomes far less polar, down to a value of about six at the critical point. The relative permittivity (dielectric constant) similarly reduces if the hydrogen bonding is broken by other means such as strong electric fields but not with pressure. The change in the dielectric with temperature gives rise to considerable and anomalous changes in its solubilization and partition properties with temperature. These are particularly noticeable in superheated water [610], where the dielectric is low, and in supercooled water high dielectric. Also, the dielectric increases (107.7 at -35 °C, with a sudden jump to >10,000 at -40 °C in nano-sized pores [1774] ) even as the density decreases. Pressure increases the relative permittivity (dielectric constant) (101.42 at 0 °C and 500 MPa), due to its effect on the density.

 

Perhaps the high relative permittivity (dielectric constant) of water should not be considered anomalous as other small polar molecules (with higher dipole moments) form liquids also having high dielectric constants (see below). The ratio dielectric constant/(dipole moment)² is often also reckoned, by others, to be anomalously high in liquid water (but note that the gas-phase, rather than liquid, dipole moments are used for comparing these substances). Although high, clearly molecules with zero dipole moment (e.g., CCl₄) have infinite such values. Also anomalous is the 60% increase in the static dipole moment of water molecules in liquid water compared to the gaseous state. This is due to hydrogen bonding causing greater intermolecular charge transfers during the changed distance between the molecular charge centers. Interestingly, the dynamical dipole moment ⁿ is not similarly enhanced and may even be reduced due to the electrostatic dragging effect of surrounding electrons [1993].

 

 

Shown opposite are the dipole moments (blue triangles below) and dielectric constant/(dipole moment)² ratios (red diamonds above) relative to the dielectric constants for a range of solvents. The data 1-17 correspond to 1, diethyl ether; 2, chloroform; 3, methylene dichloride; 4, methyl ethyl ketone; 5, acetone; 6, ethanol; 7, methanol; 8, acetonitrile; 9, ethylene glycol; 10, dimethyl sulfoxide (DMSO); 11, hydrazine; 12, formic acid; 13, water; 14, sulfuric acid; 15, formamide; 16, hydrogen cyanide; and 17, N-methyl formamide respectively.

 

By measuring electrostatic forces, water shows an anomalously low dielectric constant (≈ 2) next to boron nitride (BN) surfaces [3319]. The depth of this (out-of-plane dielectric constant ≈ 2) layer is about 0.75 nm (3 water layers); it remains anomalously small (out-of-plane dielectric constant < 20) up to 20 nm and does not return to its bulk value (dielectric constant ≈ 80) for ≈ 500 nm. The rationale for this behavior is that the rotational freedom of water dipoles decreases near surfaces.

 

[  Anomalies page: Back to Top ]

M9   The relative permittivity shows a temperature maximum.

 

Anomalous dielectric behavior of water is found over a range of microwave frequencies between about 2 and 100 GHz . The real (ε') and the imaginary (ε'') part of the complex dielectric constant increase then decrease with increasing temperature. Examples at two close frequencies for liquid (including supercooled) water are shown opposite [588]. This may be understood by noting the shifts with the temperature of the maximum frequency of microwave absorption and the dielectric permittivity. Analysis of the complex permittivity gives a discontinuity at about 30 °C [1045].

 

 

 

 

 

 

 

 

There is a frequency (2.96 ± 0.17 kHz.) where the relative permittivity (strictly speaking, the dielectric constant is the relative permittivity of a material for a frequency of zero) is independent of temperature (the isopermittive point). Below this frequency, the relative permittivity increases with temperature; above, it decreases. This has been explained by considering water as a system of two species: ions and dipoles [1698].

[  Anomalies page: Back to Top ]

M10   The relative permittivity shows a 'kink' in its behavior with increasing temperature at 60 °C

The relative permittivity (dielectric constant) increases considerably at low temperatures. If carefully drawn, there appears to be a kink in the behavior at about 60 °C. This anomaly was associated with a subtle change in the water's average effective dipole moment from 2.17 D to 1.87 D [2741, 2755].

 

This behavior may be associated with solute solvatochromic shifts at about 43 °C in the specific hydrogen-bonded donor acidity, the hydrogen-bonded acceptor basicity, and the general nonspecific polarizability and dipolarity, as determined using solvatochromic probes [2892].^q The temperature for this anomaly has a pH dependency from ~35 °C at pH 3.5 to ~45 °C at pH 7 to 9 [3828].

 

There is also a change in the dielectric with pressure (pressure coefficient) that also shows an unusual behavior [1912] with a minimum at about 95 °C thought due to the two-state structuring of liquid water.

 

[  Anomalies page: Back to Top ]

M16    X-ray diffraction shows an unusually detailed structure

The unusually detailed X-ray diffraction is shown elsewhere and is simply explained by the presence of ordered clustering within the liquid phase. There also appears to be ordered structuring amongst the amorphous ices [3501].

[  Anomalies page: Back to Top ]

M17   Under high pressure water molecules move further away from each other with increasing pressure, a density-distance paradox. 

When liquid water is put under pressure (below about 200 MPa), the water molecules approach their neighbors more closely, as might be expected from the increase in density. However, if the pressure is increased from about 200 to 400 MPa, the average distance between neighboring water molecules increases [51]. This also occurs in supercritical water [3600]. At higher pressures, the distances reduce again (but less so) with increasing pressure. Similar and corroborative behavior is seen with the O-H stretch vibration frequency (v₁), which increases with pressure between 200 to 400 MPa [533] while reducing with pressure at higher or lower pressures. The O-H stretch data has been confirmed at 23 °C but not found at 52 °C, indicating that it requires larger clusters to be recognized [824]

 

 

 

 

The explanation for all these effects is that there appears to be an increase in the 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 (see cartoon left where interpenetration increases the density without reducing the intermolecular distances. Such behavior can also be seen to cause other anomalies (1, 2, 3, 4, 5 ). When liquid water is put under pressure (below about 200 MPa) the water molecules approach their neighbors more closely, as expected from the increase in density. The interpenetration of the water networks gives rise to high-density water similar to the interpenetrating high-density ices.

 

 

 

 

 

 

 

 

The 200 MPa pressure required for the formation of high-density water has also been confirmed by infrared spectroscopy [2099]. Similar boundaries have been found using the isothermal compressibility data (see left) [2194].

 

A similar effect is found in supercooled water, where the distance between the water molecules decreases [1489] as the density decreases as the supercooling temperature is lowered. For example, the volume of one kg of water at 4 °C is 1.000 L with the water molecules 0.282 nm apart, but at -6 °C, the volume is 1.001 L with the water molecules only 0.279 nm apart, and in ice at 0 °C, the volume of one kg is 1.091 L with the water molecules only 0.276 nm apart. This decrease in density (increase in volume) is primarily due to the reduction in nearest neighbors. Also similar is what happens in the high-density ices where, for example, ice-seven (with two inter-penetrating cubic ice lattices) under a pressure of over 2200 MPa (density 1.65 g cm⁻³) has an average O····O nearest neighbor distance about 3.5% greater than that in cubic ice (density 0.92 g cm⁻³ at 0.1 MPa). Thus the density of ice-seven is somewhat less than twice the density of cubic ice (i.e., 2 ˣ 0.92/(1.035)³ = 1.65 g ˣ cm⁻³). Also, neighboring water molecules in hexagonal ice lie closer than those in liquid water, but the ice density is less due to ice molecules having only four neighbors.

 

A similar anomaly is that the first coordination shell positions of the amorphous ices move further apart as pressure increases and the amorphous ice densities increase [3501].

 

	LDA	HDA	VHDA
Formation Pressure	0.1 MPa	70 MPa	1.1 GPa
rO-O , Å	2.750	2.780	2.803
Density, g ˣ cm−3	0.9435	1.1347	1.2616

 

The bond lengths in the tetrahedral glass GeO₂ also behave similarly as they increase under increasing pressure from 5 GPa to ≈ 30 GPa when the coordination of the Ge atoms increases [1930].

[  Anomalies page: Back to Top ]

M18   Water adsorption may cause negative resistance

 

Water coats all hydrophilic surfaces open to ambient atmospheres that are not dried, and the first absorbent layer is generally held strongly. Significantly this layer will affect the properties of the surface, including their electrical properties, and can cause negative resistance [2223]. This effect is shown right for Sr₃Co₂Fe₂₄O₄₁ under moist air conditions; a) normal ohmic conduction; b) breakdown tunneling; c) negative differential resistance, where the current decreases as the potential increases; d) water absorption [2223]. The negative resistance is expected due to the electrolytic breakdown of the adsorbed water.

 

[  Anomalies page: Back to Top ]

---

Footnotes

^a D₂O is toxic to many organisms at high levels (20%-100% D₂O, where it affects many processes, including mitosis and membrane function). However, it is not generally considered harmful at much lower levels where it is used in human physiological research. It would take many days of drinking 2-3 liters a day of deuterated water for a person to be poisoned enough to die. By use of a desiccation-rehydration process, it has been shown that the anhydrobiotic nematode Panagrolaimus superbus is able to tolerate and proliferate in 99% D₂O due to its possession of a wide water channel pore [4234].

 

There is some evidence to show that artificially reducing its natural abundance in water (0.03% w/w) may have positive effects on the health of organisms [424]. [Back]

 

^b This method for reconciling the data works poorly at low temperatures [1049]. [Back]

 

^c The reduced zero-point energy when switching D-atoms for H-atoms from free to hydrogen-bonded positions within water clusters has been shown due to the energetic consequences of lowering the bend and torsional bond energies, which are greater than the raising of the stretching bond energy [986]. [Back]

 

^d Also contributing to this effect are the relative isotopic differences between the zero-point energies of the liquid and gaseous phases. Librational vibrations (due to hydrogen-bonding) release energy when the phase changes from liquid to gaseous (where they are absent), with H₂O librations (being greater) releasing more energy and so increasing the volatility of H₂O relative to D₂O at lower temperatures. Opposite effects are apparent at higher temperatures where there is less hydrogen bonding. However, energy still needs to be supplied to provide for the increased zero-point vibrational energy of the stretch vibrations (the gaseous stretch vibrations being more energetic than those for the liquid phase) [992]. [Back]

 

^e In many gaseous-solute solvent systems (for example, N₂ in CCl₄), the solubility increases with temperature increase. Although solubility decreases with temperature increase are encountered with some other solute-solvent combinations (for example, methane in n-heptane), the behavior is a more general property of water and deserves comment. [Back]

 

^f There is evidence [157, 269, 274] that the first (clathrate) shell possesses stronger hydrogen-bonding which weakens the hydrogen-bonding out to the next shell. [Back]

 

^g The formation of O=O···H-OH hydrogen bonds may be seen as the first stage in the natural low-level formation of oxygen redox products (for example, H₂O₂) in water. As the ratio of O₂/N₂ solubilities has a maximum at 290 K, there is an indication that partial clathrate cages may be responsible for the polarization that encourages the hydrogen-bond formation. [Back]

 

ⁱ Nuclear quantum effects concern the different energies of the vibrational states. The deuterium atom bonds (being about twice as heavy as the protium atom) vibrate with less amplitude and frequency. Nuclear quantum effects are seen particularly in differences in their zero-point energy, the vibrational energy remaining at close to absolute zero. [Back]

 

^j The term 'ideal' indicates adherence to a particular equation; it does not indicate any more fundamental truth. [Back]

 

^l Within LDA clusters, the equilibrium lies more to the left, with HOD molecules being somewhat excluded [1779]. [Back]

 

^m D₂O is highly hygroscopic, absorbing water vapor from the atmosphere and so decreasing in D₂O and increasing in HDO if handled in open vessels. [Back]

 

ⁿ The dynamical dipole is the behavior of the dipole (μ) upon molecular rotation, ∂μ/∂Ω_ξ (ξ = x, y, z), where Ω_ξ is the molecular rotation coordinate around the ξ axis [1993]. [Back]

 

^o The molar volumes of most liquids and solids exhibit a volume decrease upon substitution with a heavy isotope [2621]. [Back]

 

^p   Zero-point energy is energy that remains even if the material could be cooled to 0 K (absolute zero, 0 K = −273.15 °C). As the hydrogen atoms of the water molecule are light, their motion is relatively great, and the molecules retain significant vibrational energy even at absolute zero. This zero-point energy allows easier transfer of protons between molecules and, hence, water's anomalously high conductivity. [Back]

 

^q The solutes involved are for the specific hydrogen-bonded donor acidity (o,o-di-tert-butylstilbazolium betaine/ o-tert-butyl-stilbazolium betaine), the hydrogen-bonded acceptor basicity (5-nitroindoline/ 1-methyl -5-nitroindoline) and the general nonspecific polarizability (tetra-tert-butylnonaene ttbP9) and dipolarity (N,N-dimethyl-amino-7-nitrofluorene/fluoro-7-nitrofluorene). [Back]

 

 

Home | Site Index | The anomalies of water | Water: Introduction | The icosahedral water clusters | LSBU | Top

 

This page was established in 2006 and last updated by Martin Chaplin on 30 May, 2022

This work is licensed under a Creative Commons Attribution
-Noncommercial-No Derivative Works 2.0 UK: England & Wales License