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Interfacial water and water-gas interfaces

The behavior of water at interfaces is unusual due to its strongly-hydrogen bonded nature and its ability to hold a charge.

 

V Confined water

V Surface tension

V Aqueous biphasic interfaces

V Nanobubbles
V Self-generation of colligative properties in water
V The surface of water
V The surface of ice (Why is ice slippery?)
V Metal interfaces
V Zeta potential
V Thermodynamics of the liquid-gas surface for water
V Evaporation and condensation

 

'A surface is a geometrical abstraction and not a physical reality'

W. W. Barkas, 1948  

The surface of water

The air-water surface k is a very complex system [2629] and certainly not as simple as often assumed. Even as we struggle over the best description for the bulk structuring of water and water’s interactions with ions and other solutes, the surface of water is even more challenging. Its structure varies with temperature, gas molecules binding, ions binding or being excluded, electrical gradients being set up and its dielectric and chemistry changes with apparent unpredictability. Further, both sides of the interface have µm thick unstirred layers different from the bulk(gas or liquid) phases.
 
How far do the gas and liquid surfaces extend? Certainly, the volume of liquid water considered ‘surface’ should be electrically neutral and this is one definition for the surface thickness for water. The interface must include the outer monolayer, the boundary that gives rise to any zeta potential, any double layer of ions, plus deeper layers where there is surface-influenced anisotropy in either charge or structure. The depth of this surface remains undetermined but it is likely to be somewhat greater than given by most techniques that look only at a restricted range of selected properties. There is much direct evidence for an effective depth of the liquid surface measured in many nanometers [2751]. Clearly, such a definition of surface depth cannot encompass a charged surface. Therefore, any charge seen on the surface may depend upon the surface layer thickness under consideration. It may well vary between methodologies with different probe depths.  Also confusing the issue is that the surface is rough (with ~15% increased surface area [2697] ) rather than planar and is in constant flux rather than thermodynamically static. The water surface is not flat, even in the most still conditions having surface nanoscopic waves (~ Å, [2697] ) existing on water that scatter light with low intensity [2407]. Energetic considerations indicate, however, that these surface capillary waves are unlikely to be more than about a water diameter high due to water’s high surface tension. In addition to this, several million monolayers exchange with the gas phase every second [2408] and cause evaporation-dependent temperature gradients [2714], so giving rise to long-lasting surface discontinuities and non-equilibrium [2409]

 

Typical surface density as described by molecular models. There is often a high density peak just underneath the surface.

The density, dielectric permittivity [738] and dipole moment of interfacial water change from their bulk water values to those of the gas over a distance generally regarded as less than about a nanometer. Thus, water’s solvation and ionization properties change at the interface, with most ions and hydrophilic solutes generally being less evident at the interface but non-polar gasses being more soluble there [1259]. Ions, including hydrogen and hydroxide ions, and other solutes behave differently at the surface to their behavior in the bulk. An important property of the surface concerns how it affects the local ion distribution. Some ions such as thiocyanate (SCN- ), azide (N3 -) and iodide (I-) [2778], prefer the surface whereas others avoid it, as shown by their effects on the surface tension [674] and bubble coalescence [672].

 

Interfacial water molecules at the gas-liquid surface have a strong attraction towards the bulk liquid causing a high surface tension. Liquid water at liquid-solid and liquid-gas interfaces behaves as a separate thermodynamic system from bulk water [2263]. Liquid water at interfaces can be investigated using x-ray reflectivity [2266], vibrational sum-frequency generation spectroscopy (VSFG) [1468] and atomic force microscopy [738]. Gas at air-water and other gas-water interfaces behaves like a flat hydrophobic surface g with the difference that the van der Waals interactions between the liquid and gas surfaces are negligible. The surface will be strongly attracted to probes approaching from the gas side at distances of about a micrometer and jumping into contact when still over 100 nm distant [1294], thus showing the long range nature of the attractive van der Waals forces. Interfacial water absorbs light differently to bulk water with both absorption at 270 nm [1328] and at 670 nm [2377] being described. The structure of the surface is not completely understood j but some information has been determined.

 

The necessarily under-coordinated water molecules at the surfaces of both ice and water form similar ice-like, low-density phases that are hydrophobic, stiffer, superfluidic and thermally more stable than the bulk water [2004]. Hydrogen bonding in the surface is stronger than in the bulk [1261] (and this has an effect on the osmotic pressure) but some hydrogen bonding is lost, giving a more reactive environment [594] and greater ice nucleation [914]. The increased strength of surface water hydrogen bonds is partially due to the reduced competition from neighboring water molecules [2030] but has little effect on their vibrational lifetime [1262]. This stronger bonding is due to lower anti-cooperativity and compensation for the increased chemical potential on the loss of some bonding. This surface hydrogen-bonding gives rise to long-range specific ion effects on the aerial surfaces where tiny amounts of dodecahedral cluster-stabilizing ions (such as ClO4-) affect the water clustering around distant similar ions (such as I-) [2139]. The diffusion within the surface is increased for some molecules (in the surface) but decreased for others and depends on the number of hydrogen bonds and size of the water clusters [1263]. The O···O distance, between surface water molecules, within 2-3 nm from the surface, expands by 5.9 % at 25 °C. Analysis of simple thermodynamics c shows the surface probably has considerable structuring, having identical density to that of bulk water at just under 4 °C. In addition, the surface water structuring varies less with temperature than the bulk. Refractive index study of the water-air surface reveals it to be about 1.7 nm thick at 22 °C and more dense than the bulk liquid (that is, it behaves like water at a lower temperature) [1482]. Recent vibrational spectroscopy shows this surface to be relatively homogeneous [1468a,c] although this work is questioned [1468b]. About a quarter of the water molecules each have a 'dangling' O-H group [415, 1613] pointing at a slight angle out of the water [594, 1261, 2541] whilst slightly more have 'dangling' acceptor electron positions [2334] similar to water-hydrophobe surfaces, creating a slight negative charge on the surface. a 2D-IR VSFG analysis shows fast intra- and intermolecular energy transfer processes at aqueous interfaces with stretch excitations moving both along the surface plane and down into the bulk [2412].

 

Sum-frequency vibrational spectroscopy (SFVS)shows that water's libration frequency at the air/water interface is 834 cm-1 , compared with the bulk value of ~670 cm-1 while the OH stretch, time-averaged structure and ps structural dynamics all show no change [2785]. This indicates that although there is little change in water's hydrogen bonding length at the interface, these hydrogen bonds are stiffer in terms of their rotation, as happens at other hydrophobic surfaces. SFVS also shows the doubly donor-acceptor and singly donor-acceptor water molecules are the main hydrogen bonded species in bulk water with singly donor-acceptor water molecules the primary species at the air/water interface [2894].

 

The density, dielectric (permittivity, [738b]) and dipole moment of interfacial water change from their bulk water values to that of the gas over a distance of less than about a nanometer. Thus water’s solvation and dissociation properties also change at the interface, with ions and hydrophilic solutes generally being less soluble but non-polar gases more soluble [1259]. Ions, including hydrogen and hydroxide ions, and other solutes will behave differently at the surface to their behavior in the bulk. Perhaps the most important property of the surface, after the surface tension, concerns how it affects the local ion distribution. Some ions prefer the surface as shown by their effect on the surface tension (Jones-Ray effect) [674] and bubble coalescence. Both OH- [1025, 1266] and H3O+ [1308] can sit at gas/water interfaces, although clearly not at the same time due to their rapid recombination to form H2O within this lower dielectric interface. As OH- ions are preferred over H3O+ ions (above about pH 3-4), f this generally reinforces the interface's negative charge compared with the bulk. Strong acids like HCl and particularly HNO3 re-associate at the interface [2190], so allowing their evaporation. However weak acids, such as formic and acetic acids, have strong surface preference, but dissociate more rapidly when at the surface, so reducing evaporation [2787].

 

Density of interfacial water and concentration of ions at a liquid water-gas interface; charges are not meant to be exact representations

As proven by Michael Faraday in 1843 [2471], the ‘natural’ state of such interfaces appears to be negative [1266, 1433, 1477, 1591, 1905, 1951], e as at hydrophobic surfaces [1347, 1607]. The interactions of ions [1351], and particularly H+ and OH- ions, with the interface has been reviewed [1641]. Chaotropic ions with low surface charge density and/or high polarizability (such as Cl-, Br-, I-, HO2- and O2·-, but see ion effects in foams) will favor the gas-liquid interfaces [928a] as they only interact weakly with water (and so can easily lose bound water) but are influenced by the highly polarized surface. Charge transfer causes the surface to reflect the charge on the ions close to the surface [2147], usually anions. Due to charge transfer, the water molecules at the surface may become negatively charged, even when the anion is over 15 Å away from the surface [2147]. An additional effect is charge transfer where the outside water molecules contain more hydrogen-bond acceptors whereas the water molecules just to the inside of the slip-plane contain an excess of hydrogen-bond donors [2811].

 

Aqueous radicals also prefer to reside at such interfaces [939], as do some molecular species that prefer to hydrogen bond on the outside of clathrate-like structures; superoxide h for example [1260]. The presence of radicals at the surface is further shown by their release when microbubbles collapse [2068]. Excess electrons have been found to be stable at the surface of ice for several minutes [1836].

 

A combination of kosmotropic/kosmotropic (e.g. LiCl) or chaotropic/ chaotropic (e.g. CsI) ions leads to the formation of ion-pairs at the liquid-air interface whereas there appears to be no, or only small amounts of, such contact ion-pairs in the bulk water [2416]. Small cations (kosmotropes, but see ion effects in foams) are found away from the interface towards the bulk where their requirement for efficient hydration may be satisfied and as they cannot easily be stripped of the bound water by the interface. Also, there is a very large electrostatic solvation free energy cost that prevents adsorption of low polarizability ions at hydrophobic interfaces such as oils or air [2117]. Such cations only approach the interface in response to the surface negative charge. An exception to this is the oxonium ion (H3O+), which has a much stronger preference for the surface than other small cations [1500]. In acid solutions, oxonium ions point away from the surface as they only poorly accept hydrogen bonds (but strongly donate three), with their oxygen atom pointing at the surface [1308]. This encourages these ions to sit in the surface layer [1308] in the absence of competing anions such as OH- (see interfacial ions), and can lead to the charging of hydrophobic surfaces in acid solution [1584]. Mostly however at neutral pHs, there is a lower concentration of hydrogen ions than anions at the surface.

 

Iguaco Falls, Brazil

 

The zeta potential b of the surface of water is considerable and changes markedly with solute concentration (-65 mV for deionized water [1264], -46 mV for 0.1 mM NaCl, -18.8 mV for 0.1 M NaCl [1265] ). This give a surface charge density varying from about an electron per 1000 nm2 for pure water to about an electron per 10 nm2 for 0.1 M NaCl (using the equations from [1267]). The charge at the surface of deionized water with air is similar to that found on small oil droplets in water [711c].The aerosol mists formed at waterfalls (see left) are found to be negatively charged [2049].

generation of an 'image' charge

 

Controlling the approach of ions to the air-water interface is the 'image charge' force, where ions within a dielectric medium are repelled from a dielectric interface towards the bulk liquid by an “image” charge of the same size and polarity (see right).i This repulsion pushes any surface hydroxide ions away from the interface but only by an Ångström or so. If a charge lies on the outside of the interface then its image charge is attractive within the liquid phase as shown at bottom.

Water coats all hydrophilic surfaces open to ambient atmospheres that are not dried and the first absorbent layer is generally held strongly. Importantly this layer will effect the properties of the surface including their electrical properties and can cause negative resistance [2223]. The viscosity of water at hydrophilic surfaces may be orders of magnitude greater than that in the bulk but may be reduced considerably by light (670 nm [2378]). Disrupting the 'unstirred layer' next to a surface has pronounced effects on the surface charge and surface chemistry which can last for several minutes after the stirring has ceased [2138]. Therefore under flow conditions, all surfaces should be considered as dynamic including gas-liquid interfaces where there is continuous evaporation and condensation.

 

Surface binding of gasses to expanded water clusters

Atomic force microscopy at air/water interfaces has indicated that the surface polarization causes the presence of nano-sized clusters of water d within about 250 nm of the interface [738]. The reduced density and stronger hydrogen bonds within the surface will both contribute to the stabilization of water clusters; particularly that of ES over CS full and partial clustering. Small gas molecules bind to these surface clusters due to multiple van der Waals dispersion interactions, and good fit, between the gas molecules and the clusters without the possibly negative influence caused by the necessary closure of the clusters within the bulk. This offers an explanation for the greater solubility of the hydrophobic gases at the interface as they can occupy clathrate-type water dodecahedra. They are also thought important in clathrate hydrate formation at liquid gas interfaces [2563].

 

This is also supported by the exceptionally small difference in surface density from the bulk density as shown by the abnormally large pressure coefficient of the surface tension

(Change in surface tension with pressure=change in volume on change in surface area) [1280] at ambient temperature.

 

Interestingly, the air-water interface may give rise to chiral selectivity and recognition [1285]. The surface may therefore act, as the isotropic bulk cannot, as a mechanism for the choice of chirality early in the formation of life's molecules; for example, the D-series of carbohydrates and the L-series of amino acids.

 

Recently, it has been discovered that the charge on the interface affects the freezing point of supercooled water [1737]. On a surface with no electric field, water droplets were found to freeze at around -12.5 °C. On a positively charged surface, however, the freezing point is raised to -7 °C, while if the surface is negatively charged the droplet does not freeze until the temperature reaches -18 °C. Whether this strange behavior is due to the reversal of the charge of the natural negatively-charged surface destroying the water clustering has yet to be determined.

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The surface of ice (Why is ice slippery?)

There is a recent review of the surface of ice [3006]. Hexagonal ice is a very soft material (1.5 on Mohs scale) being scratched by most crystalline solids except soapstone, Mg3Si4O10(OH)2. The surface is liquid below 0 °C due to the pressure melting caused by the one-sided attractive interaction with the underlying solid ice [1686]. There is good evidence that the surface layer is indistinguishable from normal liquid water with a thickness that depends upon the temperature (that is, the further below the freezing point, the thinner the layer; it being many nanometers thick at -1 °C but has disappeared by about -38 °C). There appears to be a weakening of the hydrogen-bonded structure of the outermost water (corrugated) layers on the basal face at 257 K. This has been interpreted in terms of a stepwise change from one to two molten layers [2842].

 

Recently it has been proposed that the quasi-liquid layer is made up of two phase types that exhibit different morphologies (droplets and thin layers) [2581]. The surfaces can be examined by laser confocal microscopy combined with differential interference contrast microscopy (LCM-DIM), which can directly visualize the 0.37-nm-thick elementary steps on ice crystal surface. The non-equilibrium growth of these layers seems to form by the deposition of (critically) supersaturated water vapor onto the ice surface rather than by the surface melting of the ice or the sublimation of ice caused by undersaturated water vapor [2581].

 

As the surface of ice, near its melting point and exposed to the air, is partially melted, it is able to take up ions and organic molecules with ease. Dissolved materials lower the freezing point and so increase the thickness of the molten layer. Also, reactions take place within this liquid-like phase. However, the pre-melted surface of ice has unique solvation properties, different from those of liquid water, in that ions dissolve but small molecular weight hydrophilic organic molecules, such as glyoxal (O=CH-CH=O), do not [2969]. At low temperatures (~150 K), the melted surface layer has completely disappeared and the surface of ice is highly ordered. It has two surface-specific vibrational modes at higher frequencies (∼3530 cm-1 and ∼3700 cm-1) [3001]. The vibration at ∼3700 cm-1 corresponds to the free surface hydroxyl group whereas the ∼3530 cm-1 vibration corresponds to the vibration of surface water molecules involving the free surface lone pair.

 

Although ice is often perceived as 'slippery' (for example, ice skating), it is also very 'sticky' (for example, the difficulty in removing ice from from car windscreens, the compaction of snow to form 'snowballs' and the ease with which two ice cubes stick together). The underlying slipperiness of ice can be explained by the tetrahedral open structuring [1859] of the liquid water surface that aids the formation of a slipping plane on confinement, whereas its 'stickiness' is due to the refreezing of liquid water confined between ice surfaces. Capillary condensation of liquid water between a tungsten tip and a hydrophobic graphite surface using a friction force microscope has been proposed to form a sticky 'ice' at room temperature [1033]. l These different perceptions depend on the speed of the relative movement between the surfaces and the presence and properties of the ultra-thin layer of quasi-liquid water/amorphous ice on the crystalline surfaces [937] that may be 10 nm thick or greater. This surface layer is easily melted further by frictional heating with the low thermal conductivity of ice reducing the loss of heat. Also, there is a deformation of the ice, on skating, due to the pressure and the ease that the ice may deform (ploughing). The trails behind the skates are due to both the melting and the ploughing. At low skating speeds, the ploughing dominates whilst at high speeds the friction in the water layer dominates [2944].

Metal interfaces

The oxygen atoms of water bind to metal atom surfaces whereas the hydrogen atoms hydrogen bond surface water [2734].The 100 surface of Pt is four coordinated allowing
fully hydrogen-bonding water Surface oxide/hydroxide, if present due to the temperature and reactivity of the metal surface, also hydrogen bond to water. The hydrophilic/hydrophobic balance of a metal surface depends on the metal atom packing arrangement and spacing and the interaction time and temperature and pressure of the water contact. The structures are often uncovered experimentally by use of scanning tunneling microscopy (STM) [2739].

 

The binding energy of the water oxygen atoms to the metal atoms Au < Ag < Cu < Pd < Pt < Ru < Rh vary between 12 and 40 kJ ˣ mol-1 [2735], mostly slightly stronger than waters' hydrogen bonding. The structure of the bound H2O has little change upon adsorption, where they preferentially lie flat on the surface due to metal atoms' interactions of their 1b1 p-like lone pair orbitals (perpendicular to the plane of the water molecule) to metal 4d-orbitals [2740].

The 111 surface of Pt is six coordinated not allowing
fully hydrogen-bonding water in first layer

 

As an example, the 100 face of a platinum crystal is four coordinated allowing a hydrogen-bonded water sheet to form (see right top [2733]). As this water layer is fully hydrogen-bonded it forms a surprisingly substantial hydrophobic barrier to further water adhesion. Any further water added forms spherical droplets on the surface and do not spread. The 111 platinum crystal face (see right bottom, bonding to the most stable flat-on-surface water molecules about 34 kJ ˣ mol-1, bonding to upright water molecules about 9 kJ ˣ mol-1 ,Pt···O distances 2.36 Å, [2735]), however, is locally six coordinated and frustrates water's hydrogen bonding in the first layer so leaving further bonding sites on the effective hydrophilic surface layer and allowing further water to spread [2733]. The interfacial water molecules can diffuse within layers and exchange with secondary hydrating water through hydrogen bond rearrangements and quantum tunneling. Due to the different binding of the 111 face compared with the 100 face, the relaxation times of water molecules adsorbed on the 111 face are much faster [2736].

 

Low-temperature scanning tunneling microscopy and density functional theory have shown the possibility of many water nano-clusters forming on metallic Cu surfaces. These hydrogen-bonded ring structures resemble the resonance structures of polycyclic aromatic hydrocarbons and allow a refinement of the so-called “2D ice rules”, which have proved useful in understanding water−ice structures at solid surfaces [2915].

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Surface potential convention

Footnotes

a The sign convention for the surface potential is generally taken as measured from the air into the solution; therefore a positive potential (χ, chi) means the bulk is positively charged and the surface is negatively charged and a negative potential (see right) means the bulk is negatively charged (Φ, phi. Galvani inner potential) )and the surface is positively charged (φ, psi, "Volta" outer potential); Χ = Φ - φ. Most water models give negative surface potentials. Not all workers follow this convention. [Back]

 

b

Zeta potential

 

 

 

 

Nanobubble showing the position corresponding to the zeta potentialZeta potential (ζ) is the apparent effective charge on a moving charged particle. It is determined from the electrophoretic mobility. Although often proposed to be the charge at the 'slipping plane', it probably lies well within the unstirred surface layer (except when maximally moved by an electric field) and several water diameters away from the actual charged surface. It is likely to be close to, if numerically smaller, than the actual charge at the surface (surface potential) of the nanobubble (particle) with that charge counteracted by a smaller amount of counterion held close by. Although its exact physical meaning is unclear, it does give an indication of the electrostatic interaction between particles.

 

Zeta potential is easily determined from the movement of such particles in an electric field and depends on the relative permittivity (dielectric constant) and viscosity. It is calculated using the Smoluchowski equation.

the Smoluchowski equation

where νE is the electrophoretic mobility, ε0 and εr are the relative permittivity and the electrical permittivity of a vacuum respectively, ζ is the zeta potential, μ is the solution viscosity, r is the particle radius, η0 is the bulk ionic concentration, z is the valence of the ion, e is the charge of an electron, kB is the Boltzmann constant and T is the absolute temperature. There is a problem distinguishing the zeta potential of similarly sized sub micron bubbles and particles in solution; thus zeta potentials of such solutions would be mixed.

 

The zeta potential changes with pH and there will be a pH at which it is zero, the isoelectric point. Colloids are least stable at this isoelectric point and , generally, most stable at higher positive or negative zeta potentials found at more extreme pH. [Back]

 

c

Thermodynamics of the liquid-gas surface for water

 

 

 

 

At the liquid-gas surface the following thermodynamic relationship holds:

dG = -SdT + VdP + γdA + Σi μdni

 

where G, U, S, T, V, P, γ, A, μ and ni are the thermodynamic quantities, Gibbs (free) energy, internal energy, entropy, temperature, volume, pressure, surface tension, surface area, chemical potential and number of moles of substance i, respectively.

From the properties of the differential: (dG/dA)TPn=surface tension (gamma),       (dV/dA)TPn=(dgamma/dP)TAn      and       (dS/dA)TPn=-(dgamma/dT)PAn.

 

Also there is the Maxwell relation (dS/dV)T=(dP/dT)V (see for example [1287]).

Although (dV/dA)TPn has units of length, this may be misleading. It is better to consider its reciprocal (dA/dV)TPn as a measure of a difference in density between the surface density and bulk density.

As H = G + TS,       (dH/dA)TPn=(dG/dA)TPn + T(dS/dA)TPn       then       (dH/dA)TPn=gamma - T(dgamma/dT)PAn (surface enthalpy).

 

From inspection of the surface tension changes with temperature, it is clear that this term (the surface enthalpy) is always positive.

As H = U + PV,       (dH/dA)TPn=(dU/dA)TPn + P(dV/dA)TPn


Also,       (dU/dA)TPn =(dU/dV)TPn  x (dV/dA)TPn       where       (dU/dV)TPn=Pi=internal pressure.

 

Therefore,       (dH/dA)TPn=(Pi + P)x(dV/dA)TPn

 

The internal (cohesive) pressure (Πi) [1279] is the work required to increase the volume at constant temperature, external pressure and solute concentrations, having the same units as pressure.

 

As dU = TdS - PdV,       (dU/dV)T =T(dS/dV)T  - P,       therefore       (dU/dV)T =T(dP/dT)V  - P       (see above),

 

and as (dP/dT)VAn  =- (dP/dV)TAn x (dV/dT)PAn = expansion coeficient/isothermal compressibility, the internal pressure may be calculated from (Πi + P) = Tα/κT , where α is the coefficient of thermal expansion and κT is the coefficient of isothermal compressibility [1279].

 

 

It follows from

(dH/dA)TPn=(Pi + P)x(dV/dA)TPn and (dH/dA)TPn=gamma - T(dgamma/dT)PAn        

 

 

that

[1519]

(dA/dV)TPn=(Pi + P)/(gamma - T(dgamma/dT)PAn) and (dgamma/dP)TAn=(dA/dV)TPn=(gamma - T(dgamma/dT)PAn/(Pi + P)

 

As Tα/κT is zero at 3.984 °C, so is (Πi + P) and both are negative below this temperature, as must be (dA/dV)TPn; Πi is zero at 3.99 °C when cohesive and repulsive components of the hydrogen bonding are equal. It follows that the densities of surface and bulk water are equal at 3.984 °C as, below this temperature, the surface density contracts relative to the bulk density (rather like what happens at the surface of hexagonal ice [937] ). Thus below 3.984 °C, bulk liquid is less dense than the surface liquid whilst above this temperature the bulk liquid is more dense than the surface liquid. [Back] [Back to Top to top of page]

 

d These clusters are apparently built up from ~100 H2O molecule clusters; the same size that forms the core to the ES clusters in the icosahedral model of water [738b] and as found by X-ray analysis in Mo-based nanodrops. [Back]

 

interfacial water surface density and charge profiles

e The charge on the surface of just theoretical water (H2O, modeled without dissociation), gives a change in the charge across the surface dependent on the depth of the surface examined, as indicated opposite [1348a]. Thus overall it is negative (relative to a positive bulk) but where the very outer layer of the interface (next to the gas) is more positive [1348c]. It is probable that the outer (gas-facing) positive contribution is due to the almost-free singly linked water molecules compensating for the negative quadrupole and dipole contributions in the denser part of the surface layer [1348b]. A recent explicit ab initio electronic charge density study shows that a negative exterior surface may be present due to electron density 'leaking' into the gas phase (suggested by the dotted line opposite) and not because of any preferential orientation of water molecules [1721], although this may have a consequential orientation effect similar to the description above.

 

It should be noted that there is some dispute over the charge at aqueous surfaces [1205, 2628]., with much of the extensive earlier work indicating that it is negative [1517]. Molecular modeling is restrictive to the number of molecules used and usually have just a single ion (H3O+ or OH-) with a few (e.g. a 1000 or so) water molecules. It should be recognized that such situations are far from realistic as there are no counter ion effects, the pH's are effectively extreme (pH ~1 for H3O+) and the particles are charged; such models usually give the H3O+ on the surface with the OH- ion buried [2577]. Also, simulations have shown that NaOH solution surfaces are negative [1677]. Stable nanoscopic droplets of water have been formed with 0.005 e- nm-2 negative charge [2298]. The view prominent here [1591] is greatly influenced by the undisputedly negative zeta potential. Also there are good reasons to suppose that the surface negative region is actually a different aqueous phase, perhaps similar to a liquid crystal, that extends for up to hundreds of nanometers or even further [1328a]. It is possible that this aqueous phase is created by the higher osmotic pressure created at the surface (e.g. see [1669]) by an excess of longer-lived water clusters created there.

 

It should be noted that it is extremely difficult to keep the surface of water uncontaminated during experiments Tiny amounts of contaminants in solution or gas phase tend to gather here. [Back]

 

Zeta potential of water over a range of pH, [1853] f There is some dispute over the charge on the surface at pH 7 [1205]. Part of this dispute may be due to the ambiguity in what is meant by 'surface'; is it the very outside or is it several molecular dimensions deep as found with the 'slipping plane'. Shown opposite is the zeta potential in water with no added ions except the necessary H3O+ and OH- [1853]. Oxonium ions (H3O+) may well favor the extreme outside surface layer as they are anisotropic (as also are hydroxide ions). However this layer has low molecular density and any net acidity would be diluted by several orders of magnitude if the complete surface layer (that is, the molecules from the surface down to those showing 'bulk' properties) was to be considered. Also they may attract hydroxide ions and eliminate their charge by forming water. Surface H3O+ ions may (preferably) evaporate from the surface [1883] leaving a negatively charged surface behind. Throughout this relatively large volume, hydroxide ions are preferred over oxonium ions. In small water cluster modeling H3O+ seems to prefer the outside of clusters [854] whereas OH- prefers a more central position [1828]. However such a small cluster modeling approach may not be applicable to extensive surfaces as optimal small clusters of just H2O do not seem to be present in bulk water. Using a continuum model (where the explicit structure of water is not included) also prefers H3O+ towards the outside of the cut-off surface rather than OH- [2360]. Such a model, however, does not work well in the majority of water studies and does not explain the presence of negatively charged interfaces found by electrophoresis. Studies with surface active pH sensitive dyes (e.g. [1947] can show the 'apparent' pH at the surface as different from the bulk pH but care must be taken over the conclusions as differences may be due to artifacts inherent in the determination as the dye completely covers the surface. [Back]


g A water layer with a thickness of up to 35 nm exists at hydrophobic surfaces. This layer is characterized by a more
ordered network of hydrogen bonds than liquid water [1714]. [Back]

superoxide with 4 water molecules


h Superoxide (O2·-) hydrates in a planar manner (see right) to H-O protons on four water molecules [2007]. It decays by reaction with its conjugate acid

HO2·,

O2·- + HO2· + H2O → H2O2 + O2 + OH-

 

in water at a rate constant of about 108 M-1 s-1 which gives a half life-period of about a second, when O2·- and HO2· are at micromolar concentration. [Back]

 

i The opposite effect occurs at metal or uncharged conductive interfaces, where ions are attracted towards the surface due to oppositely charged 'image' charges. [Back]

j Evaporation and condensation

The mechanism and rates of evaporation and condensation have proved very difficult to model, with orders-of-magnitude different results being produced [2759]. An important complexity in this (generally) non-equilibrium phenomenon is that there is a difference in temperature between the liquid surface and the interfacial vapor, depending on the vapor pressure and the temperatures of bulk liquid and bulk vapor. Even the direction of the discontinuity has proved contentious. The two empirical parameters used are the condensation and evaporation coefficients, originally defined as:

 

condensation coefficient =   number of molecules absorbed by the liquid phase

                                            number of molecules impinging on the liquid phase

 

      evaporation coefficient = number of molecules transferred to the vapor phase

                                               number of molecules emitted from the liquid phase

 

The experimental values for these coefficients depend on the underlying theory but may lie orders of magnitude below unity and not be equal to each other. The primary reasons for the condensation coefficient being less than unity is that impinging molecules may knock out an existing molecule from the liquid or they simply bounce off. The primary reasons for the evaporation coefficient being less than unity is that leaving molecules may bounce back from molecules in the vapor or simply be replaced by a vapor molecules [2760]. [Back]

 

k Although sometimes the word surface is limited to its geometrical two-dimensional meaning whilst the word interface is used to describe the thin three-dimensional layer between the phases in contact, at this site the two words are used interchangeably to describe the thin three dimensional layer unless further clarified. [Back]

 

l   Such confined meniscus water has been shown to possess 106-107 times greater viscosity than bulk water [1304]. [Back]

 

 

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This page was established in 2007 and last updated by Martin Chaplin on 30 August, 2017


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