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Explanation of the Physical Anomalies of Water (F1-F9)

Water has an unusually high viscosity
Large viscosity and Prandtl number increase as the temperature is lowered
Water's viscosity decreases with pressure below 33 °C
Large diffusion decrease as the temperature is lowered
At low temperatures, the self-diffusion of water increases as the density and pressure increase
The thermal diffusivity rises to a maximum at about 0.8 GPa
Water has an unusually high surface tension
Some salts give a surface tension-concentration minimum; the Jones-Ray effect
Some salts prevent the coalescence of small bubbles
The molar ionic volumes of salts show maxima with respect to temperature

F1    High viscosity (0.89 cP, compare pentane 0.22 cP, at 25 °C)

The viscosity of a liquid is the efficiency with which it flows and is determined by the ease with which molecules can move relative to each other. It depends on the intermolecular forces and molecular momentum exchange. At lower temperatures (< 40 °C) in water, the dominant forces are those holding the molecules together (cohesiveness) [3535]. This cohesiveness is large in water due to its extensive three-dimensional hydrogen bonding, and ( at these low temperatures) the viscosity is proportional to the wavenumber of the O-H stretching of water [3539]. As the hydrogen bonds change with the velocity of flow, the flow-rate affects the viscosity of water [3539].

 

It should be noted that although the viscosity of water is high, it is not so high that it causes too much difficulty being moved around within organisms. The Arrhenius energy of activation for viscous flow is similar to the hydrogen bond energy  (H₂O, 21.5 kJ ˣ mol⁻¹; D₂O, 24.7 kJ ˣ mol⁻¹; T₂O, 26.2 kJ ˣ mol⁻¹, all calculated from [73]; all at 0 °C and all more than doubling at -30 °C). 

 

As with other liquids the kinematic viscosity of water has a minimum; this is of interest but not an anomaly. The minimum occurs at about 700 K at 100 MPa (~10⁻⁷ m² ˣ s⁻¹) [4384]. [  Anomalies page: Back to Top ]

F2    Large viscosity and Prandtl number increase as the temperature is lowered.

The increase in the viscosity with lower temperatures is particularly noticeable within supercooled water (see opposite). The water cluster equilibrium shifts towards the more open structure (for example, ES) as the temperature is lowered. This structure is formed by stronger hydrogen bonding. In turn, this creates larger clusters and reduces the ease of movement (increasing viscosity).

 

There is similar behavior with the dynamic viscosity (η) of D₂O where

η_(D2O)(T) = 1.0544 ˣ η_{(H2O)(T - 6.498 K)}

 

where 1.0544 is the square root of the ratio of the molar masses, and 6.498 K is the thermal offset [3728].

 

The high viscosity at low temperatures is a major cause of the very large increase in the Prandtl number (Pr).

 

 

where ν = kinematic viscosity (m² ˣ s⁻¹), α = thermal diffusivity ( m² ˣ s⁻¹), C_P = specific heat, (J ˣ kg⁻¹ ˣ K⁻¹ = m² ˣ s⁻² ˣ K⁻¹), η = dynamic viscosity ( Pa·s = kg ˣ m⁻¹ ˣ s⁻¹), k = thermal conductivity (W ˣ m⁻¹ ˣ K⁻¹ = kg ˣ m ˣ s⁻³ ˣ K⁻¹).

 

 

 

 

The Prandtl number is a dimensionless number that compares energy convection and conduction. Pr ≪ 1 means thermal conduction dominates over thermal convection (heat diffuses faster than matter moves; as with liquid mercury), whereas Pr ≫ 1 means that convection is more effective in transferring energy away from an area, compared to conduction; as with motor oils.

 

There is a minimum in Pr at 254.7 °C (Pr = 0.8313) [540].

 Anomalies page: Back to Top ]

F5    At low temperatures, the self-diffusion of water increases as the density and pressure increase.

The increase in self-diffusion with density (within the range of about 0.9 g ˣ cm⁻³ up to about 1.1 g ˣ cm⁻³, at low temperatures) is in contrast to normal liquids where increasing density decreases self-diffusion as the molecules restrict each other's movements; i.e., liquid water becomes more fluid when squeezed.

 

The density increase may be due to increasing temperature, below 4 °C, at atmospheric pressure or due to increasing pressure at low temperatures. Liquids typically show reduced self-diffusion when they are squeezed. However, water at 0 °C increases its diffusivity by 8% under a pressure of about 200 MPa [226], and the diffusivity of supercooled water at -30 °C increases by 60% with a similar pressure increase. The temperature limit for this anomalous behavior is ≈ 42 °C ±5 °C in agreement with the limit of the compressibility anomaly [1970]. Further increase in pressure reduces the diffusivity in common with the behavior of other liquids. The movement of water becomes restricted at low temperatures as the more open (lower density) structure produced on cooling (see above) is formed by stronger and more complete hydrogen bonding, which reduces the self-diffusion.

 

The dashed lines (above and left) indicate the diffusivity maxima. The two-state mode has been used to explain the maxima. The strongly hydrogen-bonded and structured low-density water is converted to high-density water under pressure due to bending and breakage of the hydrogen bonds [2890].

 

The strength of the hydrogen bonding is a controlling influence in this anomalous region, where the hydrogen bond angles and the inter-molecular distances are strongly coupled, and this order decreases on compression [169] due to the collapse of ES structures to CS structures. Simulation studies have shown that self-diffusion goes through a minimum as the density of water is reduced below about 0.9 g cm⁻³ followed by an increase with further density reduction, as might be expected from most liquids [402], due to the disruption of the network at low-density as the now-stretched hydrogen bonds are broken [626]. The maximum in the self-diffusion is brought about as at even higher pressures (> 200 MPa) there is an increased packing density due to the gradual phase transition to interpenetrating hydrogen-bonded networks.

 

For similar reasons involving the collapse of the hydrogen bonding, the ice surface diffusion coefficient at high pressure (10 MPa) is more than twice that observed at atmospheric pressure [1708].

 

For the same reasons, the molecular rotational movement of water (reciprocal rotational relaxation time; Debye relaxation time, τ_D ) also varies in direct proportion to the changes in self-diffusion (translational motion) [2890]. Thus, the rotation and translation of water are coupled [1839]. [  Anomalies page: Back to Top ]

F6    The thermal diffusivity rises to a maximum at about 0.8 GPa.

The thermal diffusivity (),^b which arises from vibrations in the water network [713], shows less anomalous temperature and pressure behavior than might be expected. This is due to dependence on the anomalously behaving, counteractive thermal conductivity, density and specific heat capacity. There is, however, a steep increase in thermal diffusivity at low temperatures and a maximum in the low-temperature thermal diffusivity - pressure behavior at about 0.8 GPa (see left, 25 °C) [614].

 

There will likely be a minimum in the thermal diffusivity-temperature behavior at about -30±15 °C at atmospheric pressure in line with the specific heat (C_P) and thermal conductivity. A modeling approach using TIP5P gives the minimum at ≈ 250 K [1352]. [  Anomalies page: Back to Top ]

F7    High surface tension (72.75 mJ ˣ m⁻², compare CCl₄ 26.6 mJ ˣ m⁻² at 20 °C)

Surface tension was first described by Galileo in 1612 ('Discorso intorno alle Cose che Stanno in su l'Acqua'). It is the reason that liquid surfaces tend to shrink to give the minimum surface area possible, and denser-than-water material appears to float on water, e.g., water striders, a paper clip can be made to float. Macroscopically, it can be understood as a force acting along the interface or the excess free energy per unit surface area of an interface between two bulk phases. The units of surface tension is that of energy per unit area (J ˣ m⁻²; and equivalent to an attractive force per unit length N ˣ m⁻¹) with an order of magnitude equivalent to the intramolecular bond energy divided by the molecular cross-sectional area.

 

Surface tension at a gas-liquid interface is produced by the attraction between the molecules being directed away from the surface as surface molecules are more attracted to the molecules within the liquid than they are to molecules of the gas at the surface (see right),

 

Surface free energy,

 

In contrast, molecules of water in the bulk are equally attracted in all directions. If the surface area of the liquid is increased, more molecules are present at the surface, and work must be done for this to occur. In order to achieve the greatest possible interaction energy, surface tension causes the maximum number of surface molecules to enter the bulk of the liquid and, hence, the surface area is minimized. The surface area is known but not the volume of liquid involved.

 

Water has an abnormally high surface tension ^c and surface enthalpy ^d with an abnormally tightly packed surface and very high intramolecular bond energy with a small molecular cross-sectional area. ^e Using a dynamic system, the complex surface tension of water

 

σ* = 0.073 + i(0.017 ± 0.002) N ˣ m⁻¹

 

at room temperature has been determined [2155]. Water molecules at the liquid-gas surface have lost potential hydrogen bonds directed at the gas phase and are pulled towards the underlying bulk liquid water by the remaining stronger hydrogen bonds [214]. Energy is required to increase the surface area (removing a molecule from a more-fully hydrogen-bonded interior bulk water to the lesser hydrogen bonded surface), minimizing and holding the surface under tension. The forces between the water molecules are several and relatively large on a per-mass basis, compared to those between most other molecules. Also, as the water molecules are tiny, the surface tension is large. Lowering the temperature greatly increases the bulk hydrogen bonding, causing increased surface tension. The high surface tension controls drop formation in clouds and rain.

 

From a simple theory [4117], the surface tension should be a function of the temperature difference from the critical point (∝ρ^2/3). This is observed approximately for most liquids but not for water [4118]. There is no apparent anomaly in the surface tension/temperature behavior [IAPWS] above the deeply supercooled regime. There were inflection points reported at about +4 °C [865] and 256 °C [2142] (262 °C [427]). The high-temperature anomaly has been found to coincide with the maximum in the surface enthalpy and the hydrogen bonding percolation threshold of water molecules at the liquid-vapor interface [3971]. The low-temperature inflection point was explained by use of a two-state mixture model involving low-density, and higher density water clusters [866]. However, this anomaly has not been found using the capillary elevation method [2142] or a counter-pressure capillary rise method [2632]. There is, however a significant positive deviation from the extrapolated surface tension behavior below 235 K consistent with the tail of an exponential growth in surface tension as temperature decreases [2737]. This is thought due to support the coexistence of two liquid forms in pure water of macroscopic size at these low temperatures. In 2020, a significant deviation from the extrapolated [IAPWS] formulation for the surface tension of ordinary water was detected below 253 K (-20 °C) in agreement with an anomaly in the course of surface tension in the deeply supercooled region [3949].

 

However, the surface enthalpy/ temperature behavior is anomalous, with a clear minimum at the temperature of maximum density.^a This is a consequence of the minimum in the surface entropy/ temperature behavior. Using molecular simulations there is a clear local minimum in the surface excess entropy (ΔS/Area) at a temperature ~240-250 K, indicating a second inflection point [3547]. This is due to the SPC/E water surface being denser than expected at these low temperatures. An interesting, if usually ignored, phenomenon is the linear reduction of surface tension with increasing relative humidity; ≈ 0.1% drop per 1% increase in humidity at 5 °C [1854].

 

Surface tension changes differently from bulk water properties due to surface enrichment with water clusters. The effect of microwave irradiation on the surface tension of water also shows unique properties, with the water surface tension remaining well below its original value for an extended period (minutes after the return of the temperature) [2208].

 

The more significant than expected drop in surface tension with temperature increase (0.155 mJ ˣ m⁻² ˣ K⁻¹ at 25 °C) is one of the highest known and similar to that of the liquid metals. It has been quantitatively explained using spherically symmetrical water clustering [376]. The thermodynamic change in surface tension with pressure is very high at 25 °C [1280]. ^e

 

It is interesting to note that surfactants lower the surface tension because they prefer to sit within the surface layer, attracting the surface water molecules in competition to the bulk water hydrogen-bonding and reducing the net forces away from the surface (that is, the surface tension). Many organic molecules, both hydrophilic (ethanol) and hydrophobic (neopentane), prefer the surface of the water to its bulk [1889].

 

The high surface tension of water endows it with some rather unexpected properties. Thus, water drops may rise up an inclined plate against gravity if subjected to symmetrical vibrations of about 100 Hz [1311]. This is due to the unequal changes in contact angle at the top and bottom surfaces, creating upwards forces greater than that due to gravity and the nonlinear friction effects. High surface tension is responsible for many biological effects, including the ease with which water produces lingering clouds of small droplets (aerosols) that can spread pathogens widely from sneezes and high flush toilets [2259]. The water surface can support insects on its surface (see above left) or hanging from its surface (see above right).

 

From the data on charged droplets [2660], surface charge gives rise to a significant reduction in surface tension. Also, if a small drop of water (typically 1 mm diameter) is coated in a fine (typically 20 μm diameter) hydrophobic dust, then the drop can roll and bounce without leakage [225], and the aqueous spheres can even float on water. ^j Capillarity holds the dust at the air-liquid interface, with the elasticity being due to the high surface tension. Similar material is known as 'dry water', behaving as a dry powder but releasing ≈ 95%-98% liquid water on a mechanical action such as rubbing on the skin in cosmetics [1660]. This 'dry water' powder can efficiently take up and hold large amounts of CO₂ as its clathrate (≈ 25% CO₂ by weight) [1929]. [  Anomalies page: Back to Top ]

F8    Some salts give a surface tension-concentration minimum; the Jones-Ray effect.

 

There is a shallow minimum in the surface tension of many ions at very low ionic concentrations (known as the Jones-Ray effect [674]; first dismissed erroneously as an artifact by Irving Langmuir [1518]). It has been definitively demonstrated that the Jones-Ray effect is not caused by surface-active impurities [3464]. At low concentration (< 1 mM), the surface tension of KCl solutions drops (≈ -0.01% change) with increasing concentration. A decreasing surface tension usually corresponds to a surface enhancement, whereas an increased surface tension corresponds to a surface deficit of the material. However, in this case, the drop in surface tension has been attributed to a bulk effect rather than a surface effect [2550]. At such low concentrations, the decrease in Δγ does not stem from the surface segregation of the bare ions but rather from the relative stability of the weakly oriented water surrounding the bulk ions. An increase in the orientational order of the water hydrogen-bond network entails an entropic penalty that is greater in the bulk than at the surface, leading to a net favorable interaction with the surface and, hence, to a decrease in Δγ.

 

The affinity of chaotropic ions for the expanded and weakly hydrogen-bonded surface water structure (aided by the excess of 'lone pair' electrons directed towards the bulk [594]) may help explain the sum frequency generation vibrational spectroscopy (SFG) data. This data suggests that the ions interact with the outermost interfacial water molecules, weakening their average water dipole moment, and decreasing their Gibbs free energy and thereby reducing the surface tension [2250].

 

The supposed preference of H₃O⁺ for the surface in some acid solutions (presumed due to its likely surface active nature, as its O atom is not hydrogen-bonded) is indicated by the slight drop in surface tension with high HCl, HNO₃, and HClO₄ acid concentrations (see right, [2396]). However, the anion is critical as the same effect is not shown by H₂SO₄; thus the H₃O⁺ ions in H₂SO₄ solutions show no preference for the surface. Also, NH₄OH (but not NH₄Cl or NaOH) shows a much more marked reduction in surface tension with concentration than these acids, which by a similar, if possibly an equally erroneous argument, might indicate a greater preference of hydroxide ion for the surface. ^h The surface-active nature of these acids and bases is more easily and more consistently explained by the formation of uncharged species (for example, HCl, NH₃) at the surface, coincident with their volatility (HCl forms 10% ion-pairs [2654]). Since writing this, the effect has been confirmed for HCl and HNO₃ [2190].

 

The drop in surface tension with surface-active acids is easily explainable due to greater entropy at the surface and hence lower surface energy as given by the Gibbs adsorption equation:

 

-dγ = Σ_i (Γ_idμ_i)

 

where dγ is surface tension change, Γ (capital Gamma) is the surface excess of component i, and μ is the chemical potential of component i.

 

The slight rise in surface tension with salts is less easy to understand but is thought due to the relative depletion of salt within the surface, which means that when such ions do absorb near the surface a depletion layer closer to the surface must also be created. Also, higher concentrations of such salts must disproportionately increase the attractive forces on the surface water molecules, consequently adding to the increase in the surface tension. This has been further explained by the salts being at a lower concentration nearer the outside due (1) to their need for hydration on all sides, and hence residence away from the outside surface, and (2) their 'image charges' repelling them from the surface [2397]. Although the outside surface of the water is apparently no different from 'pure' water, it must be more organized (lower entropy) due to the salt-organized hydration water just under this surface and thus causing the rise in surface energy (surface tension).

 

Kosmotropic cations and anions prefer to be fully hydrated in the bulk liquid water and increase the surface tension by the latter mechanism at all concentrations [1885]. This partitioning is noticeable in NaCl solutions, such as seawater; the weakly chaotropic Cl⁻ occupying surface sites, whereas the weakly kosmotropic Na⁺ only resides further from the surface [928]. The polarizability of large chaotropic anions (such as I⁻) is accentuated due to the asymmetric solvent distribution at the surface and increases the strength of chaotrope-solvent interactions when at the surface [989]. Similarly to chaotropic ions, hydroxyl radicals also prefer to reside at air-water interfaces [939], the radicals are donating one hydrogen bond but accepting less than two [943]. For higher molar concentrations, the ionic surface tension increments, k_i = dγ/dc_i (mN m⁻¹ M⁻¹) have been tabulated [1981].

 

Interestingly the surface tension of pure (no CO₂) water shows little change with pH over the range of pH 1 - 13 on adjusting with just HCl or NaOH, except for a small local minimum around pH 4, which has been attributed as a probable manifestation of the Jones-Ray effect [2073]. The constancy of the surface tension indicates that OH⁻ ions saturate the surface above pH 4 and by H₃O⁺ between pH 0 and 3 such that bulk concentration changes have no effect.

[  Anomalies page: Back to Top ]

F9    Some salts prevent the coalescence of small bubbles.

Higher concentrations (often about 0.17 M) of many, but not all, salts prevent the coalescence of small gas bubbles (reviewed [672]) in contrast to the expectation from the raised surface tension and reduced surface charge double layer effects (DLVO theory). Higher critical concentrations are required for smaller bubble sizes [599]. This is the reason for the foam found on saltwater seas but not on freshwater lakes. The salts do not directly follow the Hofmeister effects (that are primarily described in terms of the individual cations or anions) with both the anion and cation having importance together with one preferentially closer to the interface than the other; for example, excess hydrogen ions [1205] tend to negate the effect of halides [622]. The explanation for this unexpected phenomenon is that bubble coalescence entails a reduction in the net gas-liquid surface. The reduction in this surface is preferred when it gives rise to an increase in the (closer) interactions between the oppositely charged ions.

 

It has been proposed that anions and cations may be divided into two groups α and β with α cations (Na⁺, K⁺, Mg²⁺, NH₄⁺) and α anions (OH⁻, F⁻, Cl⁻, Br⁻, SO₄²⁻) avoiding the surface and β anions (ClO₄⁻, CH₃CO₂⁻, SCN⁻) and β cations (H⁺, (CH₃)₄N⁺) attracted to the interface; αα and ββ anion-cation-pairs then cause inhibition of bubble coalescence whereas αβ and βα pairs do not [1657]. ⁱ Bubble coalescence is inhibited when a bulk solvated or a surface-active ion-pair is present in solution (αα or ββ, respectively), creating an effectively uncharged interface [1657].

 

This property of saltwater enables a method for desalinating seawater [2186]. A high surface area of many tiny bubbles can be produced in a bubble column evaporator (using a sintered glass gas entry) using seawater, but not in freshwater as then the bubbles coalesce. This allows the rapid production of water-saturated air that can then be condensed.

 

Related to the coalescence behavior is the electrical conductivity of water that changes on degassing [4347]. The difference in electrical conductivity after degassing the solution is small but reproducible, with the salts behaving similarly at concentration below 0.17 M. However, the salts behave quite differently and specifically when their concentration is 0.5 M. The αα pairs show reproducible lowering in the measured conductivity when degassed. However, The αβ pairs show a smaller increment in the measured conductivity when degassed.

 

These groupings do not behave as bulk-phase ionic kosmotropes and chaotropes, which indicates the different properties of bulk water to that at the gas-liquid surface. The ions likely reside in the interfacial region, between the exterior surface layer and interior bulk water molecules, where the hydrogen bonding is naturally most disrupted [605]. A similar phenomenon is the bubble (cavity) attachment to microscopic salt particles used in microflotation, where chaotropic anions encourage bubble formation [758].

 

Interestingly, the salt concentration in our bodies corresponds to the minimum required for optimal prevention of bubble coalescence [622]. As small bubbles are much less harmful than large bubbles, this fact is beneficial. [  Anomalies page: Back to Top ]

F10    The molar ionic volumes of salts show maxima with respect to temperature.

The molar ionic volumes of salts, at infinite dilution, depend on both the positive intrinsic volume of the ions and the negative volume change of the water due to the ion's electric field pulling on their neighboring water molecules. Their behavior with respect to temperature is thus mainly derived from how they can disrupt the structuring in water (i.e., contract the clustering). Shown opposite is the behavior of two ionic kosmotropes, SO₄²⁻ (intrinsic volume 52.94 cm³ ˣ mol⁻¹) and Mg²⁺ (intrinsic volume 1.62 cm³ ˣ mol⁻¹) and two ionic chaotropes, ClO₄⁻ (intrinsic volume 60.15 cm³ ˣ mol⁻¹) and Cs⁺ (intrinsic volume 21.38 cm³ ˣ mol⁻¹) [1599, 1912].

 

They all can hold water increasingly strongly (relative to the water holding water itself) at higher temperatures (where water-water hydrogen-bonding is more disrupted), and as the temperature is lowered at low temperatures (where the salt interacts with the clathrate clustering).

 

The effect also causes a shrinkage in the volume that changes with temperature (see left, [2750).

 

Additionally, kosmotropes and chaotropes behave differently, with the chaotropes' molar volumes changing less with temperature and reaching their maxima at higher temperatures [1599, 1912].

 

[  Anomalies page: Back to Top ]

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Footnotes

^a If the equation for 'slip boundary' solutes, where the solute diffusion does not involve the fixed shell of solvent molecules assumed in the above equation, is used then the hydrodynamic water radius is close to correct at 1.64 Å at 25 °C. [Back]

 

^b At temperatures between 100 °C and 400 °C, the thermal diffusivity scales as the square root of the absolute temperature (Diffusivity/√T  density [614]). [Back]

 

^c A freshly exposed surface of water would be expected to have much higher surface energy (≈ 0.180 J m⁻² [1255] ), with the surface tension reducing as hydroxide ions build-up at the water-air interface [1905]. [Back]

 

^d Surface enthalpy (H_surf, also known as the total surface energy) may be calculated from the binding energy lost per unit surface area (= molecules per surface area ˣ binding energy lost per molecule. If the surface is only half occupied with water molecules that have lost about a third of their hydrogen bonds, the surface enthalpy should be = 0.5 ˣ (10¹⁹ molecule ˣ m⁻²) ˣ (1/6.022 ˣ 10²³ mol molecule⁻¹) ˣ 1/3 ˣ (45 kJ ˣ mol⁻¹) = ≈ 0.125 J ˣ m⁻² (compare with the actual value of 0.118 J ˣ m⁻² at 25 °C). As surface tension (γ) can be described as,

 

γ = ΔG_surf = ΔH_surf -TΔS_surf

 

Surface enthalpy, H_surf ≡ (dH/dA)_P and surface entropy S_surf ≡ (dS/dA)_P (where A is the area) can be derived by measuring the temperature dependency of the surface tension of water (γ) since S_surf = -(dγ/dT)_P and H_surf = γ - T ˣ (dγ/dT)_P [2302]. [Back]

 

^e The influence of pressure on the surface tension of water, as with other liquids, is not straightforward. There are two clear effects. Firstly, the thermodynamic relationship relating surface tension to pressure has been shown to equal the change in volume associated with forming a more extensive surface, [1283]. The expression may be taken as a measure of the difference in density of the liquid in the bulk compared with that at its surface and is therefore generally positive (that is, the surface tension should increase with pressure about +0.7 mJ ˣ m⁻² ˣ MPa⁻¹ for water at 25 °C). The pressure coefficient of the surface tension ( = surface enthalpy/internal+external pressure, = 0.696 nm at 25 °C) is generally much higher than for other liquids; for example, methanol (0.159 nm), diethyl ether (0.176 nm), benzene (0.178 nm) and even mercury (0.398 nm) [1280]. This high value for water indicates that the density at the surface of the water is more similar to the bulk liquid than occurs in most other liquids (see the thermodynamic derivatization). Anomalously amongst liquids, the densities of surface and bulk water are equal at 3.984 °C (at atmospheric pressure, as calculated from the equations given in [1280]). Below this temperature, the bulk liquid is less dense than the surface liquid.

 

However, the thermodynamic relationship does not hold for real liquid-gas systems where the application of pressure will cause water vapor to condense and gas molecules to adsorb on to the liquid-gas interface. The adsorption of gas molecules to the surface of liquid water lowers the surface tension by a greater extent than the thermodynamic effect outlined above (except perhaps for helium). Thus, the surface tension of water, in contact with other molecules in the gas phase, drops with an increase in pressure due to the surface activity of surface-absorbed gas molecules [1282]. The extent of this lowering depends upon the gas involved and is much greater for hydrophilic gases, such as CO₂ (-7.7 mJ ˣ m⁻² ˣ MPa⁻¹), than nonpolar gases such as N₂ and O₂ (-0.8 mJ ˣ m⁻² ˣ MPa⁻¹). [Back]

 

^f A similar effect occurs in kosmotropic salt solutions such as MgCl₂. As the concentration of salt increases, the glass transition temperature increases, as does the disparity between the rotational and translational diffusivities. With the rotational diffusivity almost independent of the viscosity, this involves a breakdown of the Stokes-Einstein relationship [1451]. This is explained, like the explanation for the effect in deeply supercooled pure water, as sticky clusters (here, hydrated salt ions) continually jam into each other (thus high viscosity) but with intervening space where water molecules can rotate unimpeded. [Back]

 

^g The surface enthalpy/temperature curve was calculated from a combination of sixth power fits four ranges of surface tension data, given in [865] and [IAPWS]; and assuming the effect of changes in pressure on the change in surface tension with temperature was insignificant. Due to noise in the data and the lack of data below 250 K, the form of the curve at very low temperatures is error-prone. [Back]

 

^h It has been proposed that the lesser hydration energy of OH⁻, relative to H₃O⁺, results in OH⁻ preferring the surface over the H₃O⁺ [1025], which also has some, but less, preference for the surface [1205,1308], and biases a pure aqueous interface to give it a negative potential [1205c, 1308]. This phenomenon, even if correct, cannot be the whole story as ions with even lower hydration energy do not seem to readily replace hydroxide ions at the interface [1505]. [Back]

 

ⁱ Originally, it was proposed that β anions (ClO₄⁻, CH₃CO₂⁻, SCN⁻) avoided the surface and α anions (OH⁻, Cl⁻, SO₄²⁻) were attracted to the interface [1305] and there is still some confusion over this theory [2650]. [Back]

 

^j Pure water droplets also bounce off the surface of the water and may float on the water surface [3285]. [Back]

 

 

 

 

 

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