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solution interface
PHYSICAL, CHEMICAL AND STRUCTURAL ASPECTS THE ELECTRODE/SOLUTION INTERFACE*
OF
SERGIO TRASATTI Laboratory
of Electrochermstry
of the Unwewty,
(Receiwd
Abstract-The
29 Nooemhrr
Via Venezian 21, 20133 Milan, Italy 1982)
structure of interfacial water at metal electrodes is inferred and discussed on the basis of the
experimental evidence resulting from the use of such diverse techniques and approaches as thermodesorption spectra, electroreflection spectroscopy, dielectric studies, quantum-chemical calculations, work function measurements, electrochemical adsorption and kinetic investigations, heat of adsorption, potential of zero charge determination. The behaviour at metal surfaces is shown to parallel that at oxides since in both cases the same factors are basically involved. Qualitative and quantitativedifferences are pointed out forsp,d- and sd-metals. Besides the more usuat aspect of water preferential orientation, also the role played by the surface electrons of the metal is illustrated. Structural effects related to the atomic configuration of the surface are finally taken into consideration.
The role of the interfacial region at the electrode/solution boundary can hardly be overemphasized. Any electrochemical event is affected by its structure[i, 21 and this is important not only from the fundamental but also from the practical point of view. The purpose of this paper is to illustrate some of the experimental facts, ascertained by diverse techniques, indicating that structuring of the solvent (in particular, water) occurs near a solid surface (in particular, metal surfaces). Since covering all relevant aspects is impossible here, the treatment will be restricted to uncharged solid surfaces in contact with solutions of not specifically adsorbed electrolytes. Particles in the surface region of two immiscible phases in contact experience unbalanced forces. For this reason, the properties of the phase in the interfacial region differ from those in the bulk. In particular, polar [or polarizablc) molecules usually adopt a preferential orientation[3,4]. Many different techniques are used to investigate the structure and the properties of the interfacial region. These are distinguished into traditional and non-traditional ones[5]. The former are the classical electrochemical methods: measurements of potenlial of zero charge, capacitance, interfacial tension. These techniques enable pieces of direct information to be obtained from the interface. Radiometric measurements may be included in this category. Indirect information on the structure of the interfaclal region can be obtained from electrochemical adsorption and kinetic studies[2,4,6]. The latter group of techniques includes those largely borrowed from the field of surface physlcs[7, 81. They are distinguishable[5,9] into rn situ and rx ;ritu techniques. The former include: Kaman, MBssbaucr, reflection spectroscopics, electron spin resonance, X-ray diffraction, ellipsometry, etc. The latter include Auger spectroscopy, low energy electron diffraction, X-ray photoelectron spectroscopy, electron stimulated desorption, electron energy loss * Plenary lecture delivered at the 33rd ISE Meetmg in Lyon on 6 September 1982.
spectroscopy, work function measurement, etc. A special arrangement for the study of the interfacial region is achieved by so-called “emersed electrodes”. This experimental approach developed by Hansen et a1.[10] consists in partially removing the electrode from the solution at constant electrical polarisation. The “emersed” portion of the electrode surface has been shown to retain its electrical double layer with the same structure and the same composition as the “immersed ” portion. This arrangement provides a unique opportunity of “looking” at the interface without any screening effect by the bulk of the liquid phase[ll]. Each technique reveals a specific property of the interface so that only a partial view of that region is obtained. Experimental studies are supplemented by theoretical approaches based on modeling of the interfacial region[4, 12-151. Before analysing the experimental evidence in the case of electrodes, it may be useful to look at what an interfacial region at a solid surface might be like. Drost-Hansen[l6], on the basis of thermal anomalies for a number of properties, has provided a molecular picture of water in contact with strongly polar (hydrophilic) and non-polar (hydrophobic) surfaces (Fig. I). The former can be regarded as structure-breaking and the latter as structure-making surfaces. The intrinsic structure of water (treated by Drost-Hansen in terms of a mixture model) is enhanced near non-polar surfaces. The enhanced order decays to the bulk one through a disordered region. Water molecules at hydrophilic surfaces are strongly oriented. The ordered region propagates mto the liquid by hydrogen bonds to a depth depending on the nature of the solid-water interaction, and decays to the bulk order through a disordered region. The total thickness of the interface is suggested to be of the order of up to l(t20 molecular layers for some especially hydrophilic solid surfaces. Oxides are relevant examples of polar surfaces. When in contact with water molecules, their surface becomes covered by a “carpet” of OH groups (Fig. 21, whose abihty to dissociate depends on the nature ofthe
1083
SERGIOTRASATTI
Fig. 3. Differential heat of water adsorption on a-Fe,O, as derived from measurements of immersion heat[ 191. The heat decreases with coverage (the number of monolayers is Indicated) and becomes .zero (corresponding to the constant liquefaaction heat of water) only as the layer of water molecules hydrogen-bonded to the surface OH groups has been completed.
(b) Fig. 1. Models of interfacIal water at polar (a) and non-polar (b) surfaces after Drost-Hansen[ 163. Polar surfaces Induce a strong order and behave thus like structure-breaking surfaces. The local order changes into the bulk order through a strongly disordered region whose thickness varies with the depth of the strongly ordered region. At non-polar surfaces, the bulk order of water ia enhanced (structure-making surfaces). A disordered regmn, whose thickness IS much thmner than in cast (a), 1s still envisaged to exist bctwccn bulk and surface water.
metal-OH bond[l7, IS]. In the case of oc-Fe,O, the heat of adsorption of water (derived from the immersion heat and measuring the strength of the oxidewater interaction) decreases to the normal Iiquefaction heat after two molecular layers (Fig. 3) 1191. The layer ofOH groups (chemisorbed water)and the first layer of water molecules hydrogen-bonded to the OH groups are immobile. The apparent dielectric constant rises sharply to the bulk value of water starting from the third adsorbed layer (Fig. 4)[20]. The thickness ofthe interfacial region is expected to depend on the propagation of the surface effect through the strength of the hydrogen bonds. In the case of silica, a strongly acid oxide[21], the entropy of water adsorption supports Drost-Hansen’s picture. The spurface entropy decays to the bulk value through a region of more negative values (enhanced order) cxtcnding for a few molecular layers (Fig. 5)[22].
(a)
Fig. 2. Sketch showing the mechanism of”hydroxylation” of an oxide surface. Sold circles are metal ions; open circles are oxygen atoms. As water molecules come IKIcontact with the surface, they become adsorbed on the metal ions. This situation is however unstable in most of the cases and transfer ofprotons to neighbouring oxygen eons occurs with exlensive formation of surface hydroxyl groups. The ability of these OH to dissociate in solution is expected to depend on the nature of the metal ion-oxygen Interaction at the various SlWS
Fig. 4. Apparent dielectric constant of water adsorbed on xFe203 measured at 100 Hz and different temperatures[20]. The value remams very low and constant as long as the layer of water molecules (2) hydrogen-bonded to the surface OH groups (1) has been completed.
Aspects of electrode/solution
Amount
1085
interface
adsorbed
CW/
mg g-‘1
Fig. 5. Differential (o) and integral (0) entropy of water adsorbed at 25X onto glass powder activated at the same temperature[22].Theentropyofsurfacewaterismorenegatlve thanthat ofbulk watec(-ppp-)overa few monomolecular layers (the 10th is marked) which indicates enhanced structuring.
In the case of metals, experiments show that they behave dif‘ferently dependtng on their electronic structure[2,4]. They can bc split into three groups[23,24]: d-metals (transition metals), sp-metals (non-transition metals), and sd-metals (01, Ag, Au). Classical eiectrochemical techniques give reliable results with .spmetals for which physical techniques have been little employed thus far. Conversely, d-metals have been much more investigated by physical techniques than by electrochemical methods. s&metals are in an intermediate position. Thermodcsorption spectra of water adsorbed on dmetals (Ru, Rh, Pt) show[25S27] the presence of three peaks (Fig. 6). That at the highest temperature has been associated[25] with the chemisorption of the first monolayer, the intermediate one with a second monolayer hydrogen-bonded to the chemisorbed molccules, and the third with ice-like multilayers. Results indicate extensive hydrogen bonding in each of the layers. These metals are known to be hydrophilic and water molecules are expected to be oriented with the oxygen atom towards the surface[2,4]. In fact, the first desorption peak correlates well with the metal-oxygen bond strength in oxides (Table 1). The preferential orientation of chemisorbed water molecules is supported by quantum-chemical calculations[B] (Fig. 7). These also show that water can break into H and OH on the surface and this is easier the stronger the
Table 1. Relation between waler thermal desorptlon and metal-oxygen bond strength -__ High temperature D(M-0) Metal Ref. (kJ mol-‘)* peak (K) Pt(lll) Rh(lll) Ru(001)
180 190 208-230
* From the heat of dissociation
f: 25
11 \\I
A
data
100 110 140
of the oxide MO[75].
Temperature
/
K
Fig. 6. Thermal desorption spectra of water adsorbed on Ru (001)at 115 K (upper part)and at 165 K (lower part). Peak A is associated with a first chemlsorbed monolayer, peak B with a second monolayer hydrogen-bonded to the first, and peak C with ice-like multilayer formation. The latter peak is in fact absent as adsorption is carried out at the higher temperature. Reprinted from[25] with permlssion of the American Institute of Physics.
1086
SERGIO
TRASATTI
Fe,+
-1
H,O
no
\ \ \ \
4
H
i \
TransItion state
\ \ ‘\
Adsorbed
Hi0
-36
L
t Adsorbed
H,O n
-4
0
-4
‘I
1
I 55A
H,O
@I
I68
ES
1 Dissociated
H
DissociatedHz0 lb)
Fig. 7. Quantum-chemical calculations of the energies of adsorption states of water on Pt(a) and Fe(b) clusters modelling the (1ll)and the (100)planes, respectively. Adsorption through the oxygen atom is seen to be favoured with respect to adsorption through the hydrogen atoms. Breakage of H,O molecules occurs through an intermediate state which can be overcome the more easily the more stable the final stale of the resulting fragments. Reprinted from[28] with permission by the North-Holland Publishmg Company.
primary chemical interaction (@Fig. 7a, b). For this reason studies of adsorption from the gas phase are usually carried out at very low temperatures, and this is a drawback if comparison with the electrochemical situation is the ultimate scope. On the other hand, thermodesorption studies indicate very Low metalwater bond streugths[29] so that no adsorption can apparently occur at room temperature. This may not reflect unambiguously the situation in the presence of the condensed liquid phase. The evidence from the electrochemical side is in fact in favour of chemisorption. Electroreflectance spectra (Fig. 8a)[30] suggest that water molecules are chemisorbed on d-metals with partial injection of electrons into the d-band of the solid. The bond strength has been found to correlate with the electronic density of states (Fig. 8a). The chemisorption bond may be envisaged to enhance the acidity of adsorbed water molecules[31]. Therefore, the propagation of the surface order may be expected to differ from metal to metal depending on the strength of the metal -water bond. No definite evidence is presently available for this. Direct information on the orientation of interfacial water molecules can be obtained by comparing work functions and potentials of zero charge[32,33]. The basic equation is [ 343 E, = ,, = Q/F + 6~ M - ,qS(dip) + const.,
(1)
where the constant depends on the nature of the reference electrode only, gs(dip) is the cwntribution to the potential associated with any layer of preferentially oriented dipoles on the solution side, and 6~~ is the
Fig. 8. (a) Electroreflectance spectra of MO (I), Ni (2), Cr (31, Co (4), V (5) and Fe (6) in 0.1 M NaOH. Modulation amplitude, 70 mV; frequency, 30 Hr. Potential/V (nhe): -0.8 (1); -0.7 (2); - 1.2 (3); -0.8 (4); - 1.2 (5); -0.8 (6). (h) Correlation between electroreflectance peak intensity of transition metals and density of electronic states at the Fermi level. Reprinted from[30] with permission by Plenum Press.
change in the surface dipole of the metal upon contact with the liquid phase. Since only E, _ 0 and @ are amenable to direct experimental determination, it ensues that the parameter resulting from a graphical comparisdn of E, = 0 and @, besides those related to
Aspects of electrode/solution interface oriented water molecules, includes also electronic effects due to the surface of the metal, If these are first neglected and everything is attributed to adsorbed water molecules, ie everything is included in gS(dip), results (Fig. 9) suggest that water molecules are oriented with the oxygen atoms towards the metal surface thus producing a more negative contribution to the potential of 0.33 V with respect to Hg taken as a reference interface. This contribution appears to he metal independent. This piece of evidence is corroborated by measurements of electron work function change upon water adsorption from the gas phase[35]. Water is again adsorbed at very low temperature to avoid desorption or fragmentation. Results (Table 2) show that the electron work function is decreased by about 1 eV for all transition metals upon completion of l-2 monolayers (Fig. 10). This does not mean that the propagation of the surface effect cannot be stronger for other properties. These data are in qualitative agreement with data of potential of zero charge. From a quantitative point of view the dift‘erence can be
OF0 __a--
___F____.______-__________5_
Pd
Rh
6
Ta
Fig 9. Apparent surface potential (measured with respect to the situation at Hg) associated with oriented water dipoles at transition metal/solution interfaces at the potential of zero charge, as a function of the enthalpy of adsorption of oxygen from the gas phase taken as a relative measure of the afinity
of the melal surface for water molecules. Ag S(dip) has been obtained from (1) and includes electronic effects. The sign of Ags(dip) implies a more negative contribution to the electrode potential with transition metals than with Hg.
Table 2. Maximum work function charlges of metal films upon adsorption of waterat 77 K[35] Metal
@(eV)*
Fe
4.40 4.91 4.88 5.71 5.40 4.64
co Ni Pt AU Cu
Aa( -0.97 ~ 1.20 - 1.10 - 1.02 - 0.60 ~ 0.73
* Work function of the bare film after annaIling at 323 K, except Au annealed al 373 K.
+ Maximum change upon saturation of the surface with water.
1087
Fig. 10. Decrease in electron work function of Fe lilms ir7. 1Onm thick upon adsorption of water at 77 K. In terms of water dipole orientation, the sign of the work function change
implies a preferential attachment of the molecules with the oxygen atoms (viz. the negative end of the molecular dipole) facing the metal surface. Most of the variation occurs during adsorption of the first monolayer. Reprinted from[35] with permission by the North-Holland Publishing Company.
accounted for by partial charge-transfer phenomena[25-301 (taking place even with rare gas adsorption[36]) and possibly by electronic effects related to the perturbation of the tail of the electron distribution at the metal surface. In fact, A@ measurements give absolute values of [ -g S(dip) + dx “1, whereas relative values [ - Ag S(dip) + ASx “1 are obtained from Fig. 9. Thus, any direct comparison of data from Figs 9 and 10 may be misleading, unless the general trend with the nature of the metal is taken into consideration. However, before appraising the importance of the electronic effects, it may be useful to look at the experimental evidence for sp-metals. Data for spmetals are available, besides with water also with acetonitrile (ACN)[37, 381 and dimethylsulphoxide (DMS0)[39,40]. Plols of potentials of zero charge against work functions show the same pattern in aqueous and non-aqueous solutions[34]. Linear relationships have been found between the two quantities. Divergence from the straight line of unit slope can be associated either with electronic effects or with solvent preferential orientation or both. Let us neglect electronic effects first. From (I), water molecules result to be preferentially oriented with the oxygen atoms towards the metal. The orientation has been found to be metal dependent and to increase with the aflinity of the metal for oxygen, as expressed by the enthalpy of formation of the monoxide, MO. Linear correlations between the surface potential and the affinity for oxygen exist for all solvents (Fig. 11). The
difyerent slopes can be related to the different structures of the solvent. Differential capacity measurements suggest that ACN, a strongly asymmetric molecule, is difficult to reorient, while DMSO can be reoriented more easily than water because its molecule is more free[23]. Direct experimental information on the electronic effects is not available. However, evidence for the role played by the surface electrons of the metal is also
c&N
w
I
:
pared with the parameter [SK M- gs(dip)] obtained from the plot of the potential of zero charge vs the work function (cf Fig. 11). Results (Fig. 12) show that calculations are very likely to overestimate ~5%~. However, the role of the metal can hardly be neglected in future work. Despite the undoubted importance of the electronic effects, the purely chemical effects can hardly be overemphasized. Linear correlations exist between the point of zero charge of oxides and the elcctronegativity of the metal ions (Fig. 13)[48]. The point of zero charge becomes more positive in terms of potential as the electronegativity increases. This parallels the findings in the case of metals if the electron work function
/
/
I- lg. 11. Apparent surface potential, measured with respst to Hg, associated with oriented solvent molecules at spmetal/solution interfaces at the potential of Nero charge, as a function of the enthalpy of the reaction M [solid) + f-0, (gas) = MO(solid), taken as a relative measure ofthe affinity of thl: metal surface for solvent molecules. Aes(dip) obtained as in Fig. 9. Data for water (W) from[23]. Data for acetonitrile (ACN) and dimethylsulphoxide (DMSO) from[37-401.
6 t
obtained when comparing adsorption of neutral molcculcs at the air/solution and the Hg/solution Interface[41]. This effect can only be calculated on the basis of the electronic theory of metals. The possible penetration of the field into the metal surface region has been emphasized by the Soviet school[42--14] but the basic idea can be found in a paper by Rice1451 in 1928. This Idea has not been developed until recently. Physicists have now invaded the field of electrochemistry trying to solve this problem. Recent work has been carried out for example by Schmlckler[46], and Badiali rl uI.[47]. The latter authors have calculated the quantity 6~~ explicitly. This value can be cook
P
w Fig. 13. Point of zero charge of oxides as a function of the oxide electronegativity. xox has,been calculated for the generic oxide M,O, from x,, = X+J(x~~~), whcrc X~ and xo are the absolute Mulhken clectronegativlties of the metal and oxygen, respectively. From[48].
.to C6 Pb c i-
.
Ln I?
>
>a-
-ArSx”g
IdrplllV
Fig. 12. Graphical comparison of calculated 66M (from Table 4 of[47]) and the experimental quantity obtamed from (1). (-) Straight lme of umt slope. Smce 9 S(dip) is expected to be positive and 6xM to be negative, the experimental quantity should always exceed the calculated one. The diagram shows that this is never the case suggesting that calculations probably overemphasize SXM.
Fig. 14. Relation between the heat of Immersion in water and the poinl of zero charge of oxides. From[4L)].
Aspects
of electrode/solution Interface
1089
Fig. 15. Change in water molecule population specifically adsorbed electrolyte as deduced from
at the interface of Au, Hg and Pb with a solution of not elcctroreflectance data as a function of charge density on the metal surface. Reprinted from[51] with permission by Else~er Sequoia S.A.
is regarded as the electronegativity of the metal surface[2]. Consistently, the immersion heat of oxides have been found [49] to increase with the point of zero charge (Fig. 14). Accordingly, a parallel increase has been found in the metal-water interaction as the metal becomes more electropositive. It is not advisable to go too far with these parallelisms quantitatively, but they offer a qualitative picture which can well support the results obtained with metals. Independent evidence for the existence of a hydrophilicity scrics for sp-metals is on the other hand available. Thus, electroreflectance measurements [SO, 511 (Fig. 15) can be interpreted in terms of changes in water population at the potential of zero charge increasing in the sequence Au, Hg, Pb. Increase in water populatmn may be associated with enhanced water monomer adsorption[4]. Indirect evidence is provided also by kinetic studies. Perbromate ion reduction proceeds[52] with adsorbed water molecules acting as proton donors: II,O+BrO;+2@+BrO;+20H~.
(2)
Calculations[53] show that the reaction rate should increase wrth increasing metal hydrophilicity. Results (Fig. 16)[54] support this idea and show that the reaction rate increases from Hg to Sb to Bi in agreement with the hydrophilicity series. Models suggested on the basis of the experimental evidence discussed above emphasize the following aspects: (i) the preferential orientation of water is with the oxygen atom towards the metal, and bonds are formed through the lone-pairs with d-metals; (ii) the bond strength is intrinsically low, some time less than that of a hydrogen bond, so that water molcculcs tend to give extended hydrogen bonding both in the adsorbed layer and with the bulk of the liquid phase. Gas phase experiments suggest[2Y] that the structure of rhe adsorbed layer may differ very little from that of a slice of water cut from a tridimensional network of hydrogen bonds (Fig. 17). This Idea has been introduced by Guidelli[ 131 in his statistical mechanical treatment (Fig. 1S), based on a multi-state water model in which dipole+lipole interactions and hydrogenbond formation are both accounted for. It should be
stressed that experiments do not give any evidence for the existence of free monomers at the interface[55, 561, although the calculation of the entropy offormation of the interface does not allow this aspect to be ruled out definitely[57]. The maximum surface potential of 0.4 V on dmetals* is understandableC2, 591 in terms of preferential orientation of water molecules with an average angle of the dipole to the surface of about 35” (Fig. 19). This is in agreement with gas phase experiments [25,29] and with calcu~dtions[60, 613, and corresponds to molecules with one of the lone-pairs directed perpendicularly to the surface. However, there is evidence that not all of the molecules in the monolayer adopt this limiting position[29]. The point is that the Helmholtz formula to calculate the potential drop across a layer of dipoles requires that a definite value should be given to the apparent dielectric constant within the layer. If the value of S-7[4] is used, then the resulting average orientation is that depicted in Fig. 19. However, dielectric studies with oxldes have shown[20] that for immobile adsorbed water the permittivity may be very close to that of a vacuum. If this is the case also for metals, the value of about 0.4 V can be compatible with the structure shown in Fig. 17. Thus, a very small change in orientation could produce a large change in potential drop. Therefore, very small energy changes are expected to be involved in going from a hydrophobic to a hydrophilic metal surface. This conclusion is supported by some indirect experimental evidence with sp-metals. If adsorption is envisaged as a water replacement reaction on the solid surface[62], the free energy change, along a series of metals, upon adsorption of organic substances may be related to the change in metalLwater bond strength, with the metal organic adsorbate interaction regarded as approximately constant[2, 631. Kcsults[4] for npentanol adsorption show (Fig. 20) that the surface activity decreases as the hydrophilicity increases. but
* This v~lur IS obtained from the relntlve ~alut: of 0.33 V (dFig. 9) by adopting the value of 0.07 V[4, 581 for the surface potential of water at Hg.
SERCIO TRASATTI
decreases in solution along the same scr~es of metals.* The difference could be related to the strength of the bond between the metal that must be displaced.
E-
surface
and water
molecules
The magnitude of this difference seems however too high to be interpreted simply in terms of heat of water desorption. More probably, some adsorption modes in the gas phase resulting in especially strong bonds arc no longer available in solution[Z, 66,671. This idea can be corroborated by observing that hydrogen adsorption in solution run parallel with the strength of the metal-H bond in metal aquo-hydride complexes along the same series of metals[64]. Structural effects come into light as single crystal faces are used as electrodes. Potentials of zero charge change with crystal face orientation mainly because of a change in the work function of the face[2]. But the difference in potential of zero charge between the (I 11) and the (110) face of face-centered cubic metals (Cu[68], Ag[69], Au[70]) appears to always exceed the difference in electron work function (Table 3). This -* If may be thought that the comparison between the gas
$,(Vl
Fig. 16. Corrected Tafel plots for the reduction of 2 x 10e4 M NaBrO, in aqueous solutions of various NaF concentrations (different symbols, from 5.2 x 10-j to 0.1 M) at different metal electrodes. Reprmted from[541 with per-
mission by Plenum Press. the free energy change from the most hydrophilic to the most hydrophobic metal amounts to only 4&XkJ mol- I. This kind of evidence is more ambiguous with dmetals[64]. The heat of adsorption of hydrogen can be Teen (Fig. 21) to increase m the gas phase whereas it
phase and the electrochemical data were to be made at the potential of zero charge to reproduce the apparent situation of uncharged surface in both cases. But in fact this approach would be physically ambiguous. First of all, for most of these metals the potential of zero charge is more negative than the reversible potential of the hydrogen reaction[65]. Thus, the heat of adsorption would be measured at coverages close to saturation thus including lateral interaction effects which are absent in the gas phase at H + 0. Second, as hydrogen is adsorbed on a bare metal surface, charge transfer occurs also in the gas phase to an extent which depends on the work function of the metal Therefore, the surface is really uncharged in the gas phase, and the work function change (ie the potential in the electrochemical situation) is different for different metals.
Table 3. Potential ofzcro charge and electron work function of smgle crystal facts of face-centered cubic metals E,_,(V,
Metal _._~._.___ Ag* AU Cu
(111, -046 0.58 -037
n/w)
Ref. (110) (100) --..--.__...____.__~__ -0.61 -077 69 0.35 019 70 -il.46 -063 68
A[(lll)-(IlO)] AE, ,,(V) A@(eV) 0.31 0.39 030
0 22 0 19 0 24
* The values of the potenttal of 7ero charge include a small ascertained specific adsorption of F-[76. 771.
Ref. 78,79 80 81
effect due to
Fig 17 Model oftheproposed strucrllreofwateradsorhed on transition metalsurfaces basedonan ~rstudy. Water molecules are hydrogen-bonded in hexagonal rings. Three mnlecules over six have one of their lonepair orhitals directed towards the metal surface. Solid segments represent hydrogen atums (LIfthe sketch on the right-hand side). Reprinted from[29] with permissmn by the North-Holland Publishing Company.
Aspects of electrode/solution
interface
1091
Fig. 18. Sketch of a tridimensional network of hydrogenbonded watermolecules terminating at a plane metal surface, resulting in a multi-state model for adsorbed water. Reprinted from[l3] with permission by Elsevier Sequoia S.A.
E,
/V
Lrhe)
Fig. 21. Heat of adsorption of hydrogen in the gas phase (0) solution (0) of not specifically adsorbed electrolytes as a function of the potential of incipient hydrogen deposition. Data from[2,67]. E, is a parameter whose unique function is to linearise the plots. The verbcal distance between the points for the same given metal measures the incidence of electronic, solvent and structural effects characteristic of the electrochemical situation.
and in
Fig. 19. Average orientation of water molecules on transition metal surfaces as resulting from the application of the Helmholtz formula ys(dip) = r~,,,~ x P”,,J x sin O/s, wnh yS(dtp) = 0.4 V (cfFlg. 9 with 0.07 V for the absolute value of the surface potential of water at Hg), ,tn,o = 10’ 5 molecule cm 2 pt, o = 6.12 x 10es3 Cm, and e = 6. With a lower value of a,*the measured &dip) could be consistent with the model depicted in Fig, 1’7,smce not all of the molecules would be required to possess the maximum orientation shown here.
indicates that the electronic effect and/or the specific orientation of water molecules vary with the crystal face[71]. In principle, since the electron work function measures the electronegativity of the metal surface, a stronger orientation is expected on the face with the lower work function. Thus, the affinity for oxygen is also expected to be face specific as any other chemical property. The same is true for the electronic effect. However, the picture of the metal-water interaction is presently less clear for sd-metals than for the other two groups[24,72,73]. It may be of interest to extend this analysis to solid metals other than silver in order to assess how much the situation at a polycrystalline surface may be complicated. Experimental data (Table 4) are scarce but it is possible to see that the “inhomogeneity” of double layer parameters shows a trend to increase with the melting point[73], ie structural effects are important only with really “frozen” surfaces. Low-melting metals exhibit a not negligible atomic self-diffusion even at room temperatureC741. Therefore, surfaces are somewhat “fluid’ and structural effects are smoothed down. Working with polycrystalline surfaces can still be meaningful with only very few low-melting metals,
Table 4. Range !O. Standard free energy of tt-pentanol adsorption from aqueous soluttons ot not specrficahy adsorbed electrolytes on various metals, as a function of the enthalpy of ibrmation of the oxtde MO assumed to be a relatrve measure of the affinity ofthe metal surface for water molecules.~l he point for Ag 15in brackets because the value of AGY, has been extrapolated from data for n-hexanol. n-Pentanol is seen to be more adsorbed on less hydrophilic metals. Data of AG:, from[4].
Metal Bi Pb Ln Ag
of potential of mro charge crystal faces of metals ~~___~ Meltmg pomt ( Cl 271 327 420 961
0.05 0.04 0.15 0.31
for ditTerent ..-
82 83 84 69
1092
SERGIOTRASATTI
but it is certainly much less useful and sometimes misleading with Cu, Ag and Au. The situation is even worse with d-metals whose degree of surface inhomogeneity is very high. Acknowfedyement-Financial support National Research Council (CNR, acknowledged.
ta this work by the Rome) is gratefully
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Tobias) (1961). 2. S. Trasatti, trockrmtcal
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(Edited by H. Gerischer and C. 10,~. 213. Wiley-Interscience, New York
Engineering
W. Tobias)Val. (1977). 3. 9. V. Derjagin, 1. pkys. Chem. 1, 29 (1932). 4. S. Trasatti, in Modern Aspects ofElectrochemistry (Edited by 9. E. Conway and J. O’M. Bock&) Vol. 13. p. 81. Plenum Press, New York (1979). 5. Proc. lnt. Conf. Non-craditionul Approaches to the So/id/Elec~rulyte Interface, SW/Y Sri. 101 (1980). 6. S. Trasatti, J. elrcrmunol. Chem. 65, 815 (1975). 7. H. Gerischer, D. M. Kolb and J. K. Sass, Adu. Phys. 27, 437 (197X). 8. S. R. Morrison, Chemical Physics of Suflaces. Plenum Press, New York (1977). 9. Extended Abstracts. Int. ConJ Electronic and Moleculur Structure o/EleetrodrfElrctroiyte Interfaces, Logan, UT, July (1982). 10. W. N. Hansen, C. L. Wang and T. W. Humperys, J. dectroanal.Chem. 90, 137 (1978). 11. D. L. Rath and D. M. Kolb, Sur$ Sci. 109, 641 (1981). 12. W. R. Fawcett, Israel J. Chem. 18, 3 (1979). 13. R. Guidelli, J. elecrroanal. Chem. 123, 59 (1981). 14. D. Henderson and L. Blum, J. electroanol. Chem. 132, 1 (1982). 15 R. Parsons and R. M. Reeves, J. electroanal. Ckem. 123, 141 (1981). 16. W. Drost-Hansen, Ind. Engng Ckem. 61, IO (1969). 17. D. N. Furlong, D. E. Yates and T. W. H&y, in Electrodes ofConductite Metallic Oxides, B (Edited by S. Trasatti) p, 367. Elsevier, Amsterdam (1981). 18. A. Daghetti, G. Lodi, and S. Trasatti, Mater. Chem. 8, 1 (1983). 19 E. McCafTerty and A. C. Zettlcmoyer. Discuss. Faraday Sac. 52. 239 (1971). 20 E. McCafferty, V. Pravdic and A. C. Zettlemoyer, Trans. Faraday Sot. 66, I720 (1970). 21. Th. F. Tadros and J. Lyklema, J. elecrroanal. Chem. 22, 1 (1969). 22. E. Nyilas, T. H. Chm, D. M. Lederman and F. J. Miwle, Colloid lnterjhce Sri. 3, 471 (1976). 23. S. Trasatti, J. rlectroanol Ckem. 123, 121 (1981). 24. S. Trasatti, .I. rleclroonal. Chem. 138, 449 (I 982). 25. P. A. Thiel, F. M. Hoffmann and W. H. Weinberg, J. &em. Phys. 75, 5556 (1981). 26. J. J. Zinck and W. H. Weinberg, J. Vat. Sci. Tecknol. 17, 188 (1980). 27. G. B. Fischer and J. L. Gland, Surf. Sei. 94, 446 (1980). 28. A. B. Anderson, Sur/. Sri. 105, 159 (1981). J. K. Sass, A. M. Bradshaw and S. 20. K. Kretzschmar, Holloway, Surf: Sci. 115, 183 (1982). V. G. 30. Ya. M. Kolotyrkin, R. M. Lazorenko-Manevich, Plotnikov and L. A. Sokolova, EIekrrokhimiyu 13. 695 (1977). L. A. Sokolova and Ya. M. 31. R. M. Lazorenko-Manevich, Kolotyrkin, Elektrokkimiya 14, 1779 (1978). 32. S. Trasatti, J. electroana[. Chem. 33, 351 (1971).
B. Damaskin, 1. Bagotskaya and N. Grigoryev, Elecrrochim. Acra 19, 75 (1974).
33. A. Frumkin,
34. S. Trasattl, Co&ids Surf: 1, 173 (1980). 35. J. M. Heras’and L. Vi&do, Appl. Sur/ Sci. 4, 238 (1980). 36. B. E. Nieuwenhuys and W. M. H. Sachtler, J. Colioid Interface Sci. 58, 65 (1977). 37. I. A. Bagotskaya and A. M. Kalyuzhnaya, Elektrokhimiya 12, 1043 (1976). 38. U. PaIm and V. Past, Usp. Khim. 44, 2035 (1975). 39. L. M. Dubova and I. A. Bagotskaya, EIektrokkimiya 13, 64 (1977). 12, 806 40. E. K. Petyarv and U. V. Palm, Elrktrokhimiya (1976). Z. Koczorowski, 41. A. Daghetti, S. Trasatti, I. Zagbrskaand 1. electrounal. Chem. 129, 253 (198 I). 42. V. A. Kiryanov, Elektrokkimiya 17, 253 (1981). 43. A. A. Kornyshev and M. A. Vorotyntev, Can. J. Chem. 59, 2031 (1981). 14, 38 1 (1978). 44. R. N. Kuklin, Elekrrokhimiya 4s. 0. K. Rice, Phys. Keu 31, 1051 (1928). 46. W. Schmickler, in Ref.[9]. J. 47. J. P. Badiali, M. L. Rosinberg and J. Goodisman, elertrounul. Chrm. 130, 31 (1981). Sot. 125, 48. M. A. Butler and D. S. Ginley, J. electrochem. 228 (1978). 49. T. W. Healy and D. W. Fuerstnau, J. Colloid Sci. 20, 376 (1965). 50. A. Bewick and J. Robinson, J. eleetroanaL Chem. 60, 163 (1975); 71, 131 (1976). 51. C. Hinnen, C. Nguyen van Huong, A. Rousseau and 1. P. Dalbera, J. electroanal. Chem. 85, 3 I (I 979). 52. N. V. Fedorovich and F. S. Sarbash, D&l. Aknd. Nauk SSSR 243, 1231 (1978). 53. N. V. Fedorovich and F. S. Sarbash, Do&i. Akad. Nauk SSSR 255,923 (1980). 54. N. V. Fedorovich, F. S. &bash, l? I. Mikhailova, V. K. Isupovand V. V. Gavnlov, Elektrokhimiya 15,587 (1979). 55. A. Bewick, in Ref.[Y]. 56. A. Bewick and J. W. Russell, J. elrctroona/. Ckrm. 132, 329 (1982). 57. G. Hills and F. Silva, Can. J. Chem. 59, 1835 (1981). Ckem. 64, 128 (1975). 58. S. Trasatti, J. electroanaL 59. S. Trasatti, J. Ckim. pkys. 72, 561 (1975). 60. F. Flares, I. Gabbay and N. H. March, Surfi Sci. 107, 127 (1981). 61. B C. Khanra, Phys. Letr. 7bA, 194 (1980). 62. J. O’M. Bock&, M. A. Devanathan and K. Miiller, Proc. R. Sot. 274A, SS (1963). 63. N. S. Polyanovskaya and B B. Damaskin, Elcktrokhimiya 16,
64.
65.
66. 67. 68. 69.
70. 71. 72. 73. 74. 75. lb.
77. 78.
531
(1980).
Trasatti, in Elecrrocatalysis (Edited by W. E. O’Grady, P. N. Ross and F. G. W111)p, 73. The Electrochemical Society. Pennington, NJ (1982). A. N. Frumkin and 0. A. Petril, Electrochim. Acta 20,347 (1975). D. J. Barclay, J. electroonai. Chem. 44, 47 (1973). S. Trasatti. Z. phys. Cllem. N. F. 98, 75 (1975). H. Hennig and V. V. Batrakov, E[ectrokkzmi)zr 15, 1833 (1979). G. Valetteand A. Hamelin. J. e[ectroana/. Chum. 45. 301 (1973). A. Hamelin and J. Lecoeur, Sur/: Ser. 57, 771 (1976). J. Lecoeur, J. Andro and R. Parsons, Surf .%,I. 114. 320 (1982). I. A. Bagotskaya and A. V. Shlepakov, Elekrrokhimr~u 16, 565 (1980). G. Valette, J. electroanal. Chem. 139, 285 (1982). R. E. Reed-Hill, Physicui Merailury? Prirrc~ples, p, 171. Van Nostrand Reinhold, New York (1964). C. Glidewell, fnorg. Chim. Actn 24, 149 (I 977). G. Valette. J. elecrroanaL Chem. 138, 37 (1982). G. Valette, J. ~lurrruanul. Chem. 122, 285 (19Si). A. W. Dweydari and C. H. B. Mee. Phgs Stnt. Sol. (u) 17, 247 (1973). S.
Aspects of electrode/solutmn 79. A. W. Dweydarl and C. H. tl. Mee, Ph.yA. St&. Sol. (a) 27, 223 (1975). 80. H. C. Potter and J. M. Blakely, J. b’rrc. Sci. Teclmol. 12, 635 (1975). 81. L. Peralta, Y. Berthier and J. Oudar, Vidr Colbq. 83 (197X).
mterface
1093
82. U. V. Palm, M. P. Pyarnoya and M. A. Salve, Elektrokhimi_w 13, X73 (1977). 83. L. P. Khmelevaya, A. V. Chirhov and B. B. Damaskin, Elektrokhmiya 14, 1304(I 978). 84. V. V. Batrakov, B. B. Damaskin and Yu. P. Ipatov, Eiektrokhrmiyn IO, 14.4 (1974).
OF
SERGIO TRASATTI Laboratory
of Electrochermstry
of the Unwewty,
(Receiwd
Abstract-The
29 Nooemhrr
Via Venezian 21, 20133 Milan, Italy 1982)
structure of interfacial water at metal electrodes is inferred and discussed on the basis of the
experimental evidence resulting from the use of such diverse techniques and approaches as thermodesorption spectra, electroreflection spectroscopy, dielectric studies, quantum-chemical calculations, work function measurements, electrochemical adsorption and kinetic investigations, heat of adsorption, potential of zero charge determination. The behaviour at metal surfaces is shown to parallel that at oxides since in both cases the same factors are basically involved. Qualitative and quantitativedifferences are pointed out forsp,d- and sd-metals. Besides the more usuat aspect of water preferential orientation, also the role played by the surface electrons of the metal is illustrated. Structural effects related to the atomic configuration of the surface are finally taken into consideration.
The role of the interfacial region at the electrode/solution boundary can hardly be overemphasized. Any electrochemical event is affected by its structure[i, 21 and this is important not only from the fundamental but also from the practical point of view. The purpose of this paper is to illustrate some of the experimental facts, ascertained by diverse techniques, indicating that structuring of the solvent (in particular, water) occurs near a solid surface (in particular, metal surfaces). Since covering all relevant aspects is impossible here, the treatment will be restricted to uncharged solid surfaces in contact with solutions of not specifically adsorbed electrolytes. Particles in the surface region of two immiscible phases in contact experience unbalanced forces. For this reason, the properties of the phase in the interfacial region differ from those in the bulk. In particular, polar [or polarizablc) molecules usually adopt a preferential orientation[3,4]. Many different techniques are used to investigate the structure and the properties of the interfacial region. These are distinguished into traditional and non-traditional ones[5]. The former are the classical electrochemical methods: measurements of potenlial of zero charge, capacitance, interfacial tension. These techniques enable pieces of direct information to be obtained from the interface. Radiometric measurements may be included in this category. Indirect information on the structure of the interfaclal region can be obtained from electrochemical adsorption and kinetic studies[2,4,6]. The latter group of techniques includes those largely borrowed from the field of surface physlcs[7, 81. They are distinguishable[5,9] into rn situ and rx ;ritu techniques. The former include: Kaman, MBssbaucr, reflection spectroscopics, electron spin resonance, X-ray diffraction, ellipsometry, etc. The latter include Auger spectroscopy, low energy electron diffraction, X-ray photoelectron spectroscopy, electron stimulated desorption, electron energy loss * Plenary lecture delivered at the 33rd ISE Meetmg in Lyon on 6 September 1982.
spectroscopy, work function measurement, etc. A special arrangement for the study of the interfacial region is achieved by so-called “emersed electrodes”. This experimental approach developed by Hansen et a1.[10] consists in partially removing the electrode from the solution at constant electrical polarisation. The “emersed” portion of the electrode surface has been shown to retain its electrical double layer with the same structure and the same composition as the “immersed ” portion. This arrangement provides a unique opportunity of “looking” at the interface without any screening effect by the bulk of the liquid phase[ll]. Each technique reveals a specific property of the interface so that only a partial view of that region is obtained. Experimental studies are supplemented by theoretical approaches based on modeling of the interfacial region[4, 12-151. Before analysing the experimental evidence in the case of electrodes, it may be useful to look at what an interfacial region at a solid surface might be like. Drost-Hansen[l6], on the basis of thermal anomalies for a number of properties, has provided a molecular picture of water in contact with strongly polar (hydrophilic) and non-polar (hydrophobic) surfaces (Fig. I). The former can be regarded as structure-breaking and the latter as structure-making surfaces. The intrinsic structure of water (treated by Drost-Hansen in terms of a mixture model) is enhanced near non-polar surfaces. The enhanced order decays to the bulk one through a disordered region. Water molecules at hydrophilic surfaces are strongly oriented. The ordered region propagates mto the liquid by hydrogen bonds to a depth depending on the nature of the solid-water interaction, and decays to the bulk order through a disordered region. The total thickness of the interface is suggested to be of the order of up to l(t20 molecular layers for some especially hydrophilic solid surfaces. Oxides are relevant examples of polar surfaces. When in contact with water molecules, their surface becomes covered by a “carpet” of OH groups (Fig. 21, whose abihty to dissociate depends on the nature ofthe
1083
SERGIOTRASATTI
Fig. 3. Differential heat of water adsorption on a-Fe,O, as derived from measurements of immersion heat[ 191. The heat decreases with coverage (the number of monolayers is Indicated) and becomes .zero (corresponding to the constant liquefaaction heat of water) only as the layer of water molecules hydrogen-bonded to the surface OH groups has been completed.
(b) Fig. 1. Models of interfacIal water at polar (a) and non-polar (b) surfaces after Drost-Hansen[ 163. Polar surfaces Induce a strong order and behave thus like structure-breaking surfaces. The local order changes into the bulk order through a strongly disordered region whose thickness varies with the depth of the strongly ordered region. At non-polar surfaces, the bulk order of water ia enhanced (structure-making surfaces). A disordered regmn, whose thickness IS much thmner than in cast (a), 1s still envisaged to exist bctwccn bulk and surface water.
metal-OH bond[l7, IS]. In the case of oc-Fe,O, the heat of adsorption of water (derived from the immersion heat and measuring the strength of the oxidewater interaction) decreases to the normal Iiquefaction heat after two molecular layers (Fig. 3) 1191. The layer ofOH groups (chemisorbed water)and the first layer of water molecules hydrogen-bonded to the OH groups are immobile. The apparent dielectric constant rises sharply to the bulk value of water starting from the third adsorbed layer (Fig. 4)[20]. The thickness ofthe interfacial region is expected to depend on the propagation of the surface effect through the strength of the hydrogen bonds. In the case of silica, a strongly acid oxide[21], the entropy of water adsorption supports Drost-Hansen’s picture. The spurface entropy decays to the bulk value through a region of more negative values (enhanced order) cxtcnding for a few molecular layers (Fig. 5)[22].
(a)
Fig. 2. Sketch showing the mechanism of”hydroxylation” of an oxide surface. Sold circles are metal ions; open circles are oxygen atoms. As water molecules come IKIcontact with the surface, they become adsorbed on the metal ions. This situation is however unstable in most of the cases and transfer ofprotons to neighbouring oxygen eons occurs with exlensive formation of surface hydroxyl groups. The ability of these OH to dissociate in solution is expected to depend on the nature of the metal ion-oxygen Interaction at the various SlWS
Fig. 4. Apparent dielectric constant of water adsorbed on xFe203 measured at 100 Hz and different temperatures[20]. The value remams very low and constant as long as the layer of water molecules (2) hydrogen-bonded to the surface OH groups (1) has been completed.
Aspects of electrode/solution
Amount
1085
interface
adsorbed
CW/
mg g-‘1
Fig. 5. Differential (o) and integral (0) entropy of water adsorbed at 25X onto glass powder activated at the same temperature[22].Theentropyofsurfacewaterismorenegatlve thanthat ofbulk watec(-ppp-)overa few monomolecular layers (the 10th is marked) which indicates enhanced structuring.
In the case of metals, experiments show that they behave dif‘ferently dependtng on their electronic structure[2,4]. They can bc split into three groups[23,24]: d-metals (transition metals), sp-metals (non-transition metals), and sd-metals (01, Ag, Au). Classical eiectrochemical techniques give reliable results with .spmetals for which physical techniques have been little employed thus far. Conversely, d-metals have been much more investigated by physical techniques than by electrochemical methods. s&metals are in an intermediate position. Thermodcsorption spectra of water adsorbed on dmetals (Ru, Rh, Pt) show[25S27] the presence of three peaks (Fig. 6). That at the highest temperature has been associated[25] with the chemisorption of the first monolayer, the intermediate one with a second monolayer hydrogen-bonded to the chemisorbed molccules, and the third with ice-like multilayers. Results indicate extensive hydrogen bonding in each of the layers. These metals are known to be hydrophilic and water molecules are expected to be oriented with the oxygen atom towards the surface[2,4]. In fact, the first desorption peak correlates well with the metal-oxygen bond strength in oxides (Table 1). The preferential orientation of chemisorbed water molecules is supported by quantum-chemical calculations[B] (Fig. 7). These also show that water can break into H and OH on the surface and this is easier the stronger the
Table 1. Relation between waler thermal desorptlon and metal-oxygen bond strength -__ High temperature D(M-0) Metal Ref. (kJ mol-‘)* peak (K) Pt(lll) Rh(lll) Ru(001)
180 190 208-230
* From the heat of dissociation
f: 25
11 \\I
A
data
100 110 140
of the oxide MO[75].
Temperature
/
K
Fig. 6. Thermal desorption spectra of water adsorbed on Ru (001)at 115 K (upper part)and at 165 K (lower part). Peak A is associated with a first chemlsorbed monolayer, peak B with a second monolayer hydrogen-bonded to the first, and peak C with ice-like multilayer formation. The latter peak is in fact absent as adsorption is carried out at the higher temperature. Reprinted from[25] with permlssion of the American Institute of Physics.
1086
SERGIO
TRASATTI
Fe,+
-1
H,O
no
\ \ \ \
4
H
i \
TransItion state
\ \ ‘\
Adsorbed
Hi0
-36
L
t Adsorbed
H,O n
-4
0
-4
‘I
1
I 55A
H,O
@I
I68
ES
1 Dissociated
H
DissociatedHz0 lb)
Fig. 7. Quantum-chemical calculations of the energies of adsorption states of water on Pt(a) and Fe(b) clusters modelling the (1ll)and the (100)planes, respectively. Adsorption through the oxygen atom is seen to be favoured with respect to adsorption through the hydrogen atoms. Breakage of H,O molecules occurs through an intermediate state which can be overcome the more easily the more stable the final stale of the resulting fragments. Reprinted from[28] with permission by the North-Holland Publishmg Company.
primary chemical interaction (@Fig. 7a, b). For this reason studies of adsorption from the gas phase are usually carried out at very low temperatures, and this is a drawback if comparison with the electrochemical situation is the ultimate scope. On the other hand, thermodesorption studies indicate very Low metalwater bond streugths[29] so that no adsorption can apparently occur at room temperature. This may not reflect unambiguously the situation in the presence of the condensed liquid phase. The evidence from the electrochemical side is in fact in favour of chemisorption. Electroreflectance spectra (Fig. 8a)[30] suggest that water molecules are chemisorbed on d-metals with partial injection of electrons into the d-band of the solid. The bond strength has been found to correlate with the electronic density of states (Fig. 8a). The chemisorption bond may be envisaged to enhance the acidity of adsorbed water molecules[31]. Therefore, the propagation of the surface order may be expected to differ from metal to metal depending on the strength of the metal -water bond. No definite evidence is presently available for this. Direct information on the orientation of interfacial water molecules can be obtained by comparing work functions and potentials of zero charge[32,33]. The basic equation is [ 343 E, = ,, = Q/F + 6~ M - ,qS(dip) + const.,
(1)
where the constant depends on the nature of the reference electrode only, gs(dip) is the cwntribution to the potential associated with any layer of preferentially oriented dipoles on the solution side, and 6~~ is the
Fig. 8. (a) Electroreflectance spectra of MO (I), Ni (2), Cr (31, Co (4), V (5) and Fe (6) in 0.1 M NaOH. Modulation amplitude, 70 mV; frequency, 30 Hr. Potential/V (nhe): -0.8 (1); -0.7 (2); - 1.2 (3); -0.8 (4); - 1.2 (5); -0.8 (6). (h) Correlation between electroreflectance peak intensity of transition metals and density of electronic states at the Fermi level. Reprinted from[30] with permission by Plenum Press.
change in the surface dipole of the metal upon contact with the liquid phase. Since only E, _ 0 and @ are amenable to direct experimental determination, it ensues that the parameter resulting from a graphical comparisdn of E, = 0 and @, besides those related to
Aspects of electrode/solution interface oriented water molecules, includes also electronic effects due to the surface of the metal, If these are first neglected and everything is attributed to adsorbed water molecules, ie everything is included in gS(dip), results (Fig. 9) suggest that water molecules are oriented with the oxygen atoms towards the metal surface thus producing a more negative contribution to the potential of 0.33 V with respect to Hg taken as a reference interface. This contribution appears to he metal independent. This piece of evidence is corroborated by measurements of electron work function change upon water adsorption from the gas phase[35]. Water is again adsorbed at very low temperature to avoid desorption or fragmentation. Results (Table 2) show that the electron work function is decreased by about 1 eV for all transition metals upon completion of l-2 monolayers (Fig. 10). This does not mean that the propagation of the surface effect cannot be stronger for other properties. These data are in qualitative agreement with data of potential of zero charge. From a quantitative point of view the dift‘erence can be
OF0 __a--
___F____.______-__________5_
Pd
Rh
6
Ta
Fig 9. Apparent surface potential (measured with respect to the situation at Hg) associated with oriented water dipoles at transition metal/solution interfaces at the potential of zero charge, as a function of the enthalpy of adsorption of oxygen from the gas phase taken as a relative measure of the afinity
of the melal surface for water molecules. Ag S(dip) has been obtained from (1) and includes electronic effects. The sign of Ags(dip) implies a more negative contribution to the electrode potential with transition metals than with Hg.
Table 2. Maximum work function charlges of metal films upon adsorption of waterat 77 K[35] Metal
@(eV)*
Fe
4.40 4.91 4.88 5.71 5.40 4.64
co Ni Pt AU Cu
Aa( -0.97 ~ 1.20 - 1.10 - 1.02 - 0.60 ~ 0.73
* Work function of the bare film after annaIling at 323 K, except Au annealed al 373 K.
+ Maximum change upon saturation of the surface with water.
1087
Fig. 10. Decrease in electron work function of Fe lilms ir7. 1Onm thick upon adsorption of water at 77 K. In terms of water dipole orientation, the sign of the work function change
implies a preferential attachment of the molecules with the oxygen atoms (viz. the negative end of the molecular dipole) facing the metal surface. Most of the variation occurs during adsorption of the first monolayer. Reprinted from[35] with permission by the North-Holland Publishing Company.
accounted for by partial charge-transfer phenomena[25-301 (taking place even with rare gas adsorption[36]) and possibly by electronic effects related to the perturbation of the tail of the electron distribution at the metal surface. In fact, A@ measurements give absolute values of [ -g S(dip) + dx “1, whereas relative values [ - Ag S(dip) + ASx “1 are obtained from Fig. 9. Thus, any direct comparison of data from Figs 9 and 10 may be misleading, unless the general trend with the nature of the metal is taken into consideration. However, before appraising the importance of the electronic effects, it may be useful to look at the experimental evidence for sp-metals. Data for spmetals are available, besides with water also with acetonitrile (ACN)[37, 381 and dimethylsulphoxide (DMS0)[39,40]. Plols of potentials of zero charge against work functions show the same pattern in aqueous and non-aqueous solutions[34]. Linear relationships have been found between the two quantities. Divergence from the straight line of unit slope can be associated either with electronic effects or with solvent preferential orientation or both. Let us neglect electronic effects first. From (I), water molecules result to be preferentially oriented with the oxygen atoms towards the metal. The orientation has been found to be metal dependent and to increase with the aflinity of the metal for oxygen, as expressed by the enthalpy of formation of the monoxide, MO. Linear correlations between the surface potential and the affinity for oxygen exist for all solvents (Fig. 11). The
difyerent slopes can be related to the different structures of the solvent. Differential capacity measurements suggest that ACN, a strongly asymmetric molecule, is difficult to reorient, while DMSO can be reoriented more easily than water because its molecule is more free[23]. Direct experimental information on the electronic effects is not available. However, evidence for the role played by the surface electrons of the metal is also
c&N
w
I
:
pared with the parameter [SK M- gs(dip)] obtained from the plot of the potential of zero charge vs the work function (cf Fig. 11). Results (Fig. 12) show that calculations are very likely to overestimate ~5%~. However, the role of the metal can hardly be neglected in future work. Despite the undoubted importance of the electronic effects, the purely chemical effects can hardly be overemphasized. Linear correlations exist between the point of zero charge of oxides and the elcctronegativity of the metal ions (Fig. 13)[48]. The point of zero charge becomes more positive in terms of potential as the electronegativity increases. This parallels the findings in the case of metals if the electron work function
/
/
I- lg. 11. Apparent surface potential, measured with respst to Hg, associated with oriented solvent molecules at spmetal/solution interfaces at the potential of Nero charge, as a function of the enthalpy of the reaction M [solid) + f-0, (gas) = MO(solid), taken as a relative measure ofthe affinity of thl: metal surface for solvent molecules. Aes(dip) obtained as in Fig. 9. Data for water (W) from[23]. Data for acetonitrile (ACN) and dimethylsulphoxide (DMSO) from[37-401.
6 t
obtained when comparing adsorption of neutral molcculcs at the air/solution and the Hg/solution Interface[41]. This effect can only be calculated on the basis of the electronic theory of metals. The possible penetration of the field into the metal surface region has been emphasized by the Soviet school[42--14] but the basic idea can be found in a paper by Rice1451 in 1928. This Idea has not been developed until recently. Physicists have now invaded the field of electrochemistry trying to solve this problem. Recent work has been carried out for example by Schmlckler[46], and Badiali rl uI.[47]. The latter authors have calculated the quantity 6~~ explicitly. This value can be cook
P
w Fig. 13. Point of zero charge of oxides as a function of the oxide electronegativity. xox has,been calculated for the generic oxide M,O, from x,, = X+J(x~~~), whcrc X~ and xo are the absolute Mulhken clectronegativlties of the metal and oxygen, respectively. From[48].
.to C6 Pb c i-
.
Ln I?
>
>a-
-ArSx”g
IdrplllV
Fig. 12. Graphical comparison of calculated 66M (from Table 4 of[47]) and the experimental quantity obtamed from (1). (-) Straight lme of umt slope. Smce 9 S(dip) is expected to be positive and 6xM to be negative, the experimental quantity should always exceed the calculated one. The diagram shows that this is never the case suggesting that calculations probably overemphasize SXM.
Fig. 14. Relation between the heat of Immersion in water and the poinl of zero charge of oxides. From[4L)].
Aspects
of electrode/solution Interface
1089
Fig. 15. Change in water molecule population specifically adsorbed electrolyte as deduced from
at the interface of Au, Hg and Pb with a solution of not elcctroreflectance data as a function of charge density on the metal surface. Reprinted from[51] with permission by Else~er Sequoia S.A.
is regarded as the electronegativity of the metal surface[2]. Consistently, the immersion heat of oxides have been found [49] to increase with the point of zero charge (Fig. 14). Accordingly, a parallel increase has been found in the metal-water interaction as the metal becomes more electropositive. It is not advisable to go too far with these parallelisms quantitatively, but they offer a qualitative picture which can well support the results obtained with metals. Independent evidence for the existence of a hydrophilicity scrics for sp-metals is on the other hand available. Thus, electroreflectance measurements [SO, 511 (Fig. 15) can be interpreted in terms of changes in water population at the potential of zero charge increasing in the sequence Au, Hg, Pb. Increase in water populatmn may be associated with enhanced water monomer adsorption[4]. Indirect evidence is provided also by kinetic studies. Perbromate ion reduction proceeds[52] with adsorbed water molecules acting as proton donors: II,O+BrO;+2@+BrO;+20H~.
(2)
Calculations[53] show that the reaction rate should increase wrth increasing metal hydrophilicity. Results (Fig. 16)[54] support this idea and show that the reaction rate increases from Hg to Sb to Bi in agreement with the hydrophilicity series. Models suggested on the basis of the experimental evidence discussed above emphasize the following aspects: (i) the preferential orientation of water is with the oxygen atom towards the metal, and bonds are formed through the lone-pairs with d-metals; (ii) the bond strength is intrinsically low, some time less than that of a hydrogen bond, so that water molcculcs tend to give extended hydrogen bonding both in the adsorbed layer and with the bulk of the liquid phase. Gas phase experiments suggest[2Y] that the structure of rhe adsorbed layer may differ very little from that of a slice of water cut from a tridimensional network of hydrogen bonds (Fig. 17). This Idea has been introduced by Guidelli[ 131 in his statistical mechanical treatment (Fig. 1S), based on a multi-state water model in which dipole+lipole interactions and hydrogenbond formation are both accounted for. It should be
stressed that experiments do not give any evidence for the existence of free monomers at the interface[55, 561, although the calculation of the entropy offormation of the interface does not allow this aspect to be ruled out definitely[57]. The maximum surface potential of 0.4 V on dmetals* is understandableC2, 591 in terms of preferential orientation of water molecules with an average angle of the dipole to the surface of about 35” (Fig. 19). This is in agreement with gas phase experiments [25,29] and with calcu~dtions[60, 613, and corresponds to molecules with one of the lone-pairs directed perpendicularly to the surface. However, there is evidence that not all of the molecules in the monolayer adopt this limiting position[29]. The point is that the Helmholtz formula to calculate the potential drop across a layer of dipoles requires that a definite value should be given to the apparent dielectric constant within the layer. If the value of S-7[4] is used, then the resulting average orientation is that depicted in Fig. 19. However, dielectric studies with oxldes have shown[20] that for immobile adsorbed water the permittivity may be very close to that of a vacuum. If this is the case also for metals, the value of about 0.4 V can be compatible with the structure shown in Fig. 17. Thus, a very small change in orientation could produce a large change in potential drop. Therefore, very small energy changes are expected to be involved in going from a hydrophobic to a hydrophilic metal surface. This conclusion is supported by some indirect experimental evidence with sp-metals. If adsorption is envisaged as a water replacement reaction on the solid surface[62], the free energy change, along a series of metals, upon adsorption of organic substances may be related to the change in metalLwater bond strength, with the metal organic adsorbate interaction regarded as approximately constant[2, 631. Kcsults[4] for npentanol adsorption show (Fig. 20) that the surface activity decreases as the hydrophilicity increases. but
* This v~lur IS obtained from the relntlve ~alut: of 0.33 V (dFig. 9) by adopting the value of 0.07 V[4, 581 for the surface potential of water at Hg.
SERCIO TRASATTI
decreases in solution along the same scr~es of metals.* The difference could be related to the strength of the bond between the metal that must be displaced.
E-
surface
and water
molecules
The magnitude of this difference seems however too high to be interpreted simply in terms of heat of water desorption. More probably, some adsorption modes in the gas phase resulting in especially strong bonds arc no longer available in solution[Z, 66,671. This idea can be corroborated by observing that hydrogen adsorption in solution run parallel with the strength of the metal-H bond in metal aquo-hydride complexes along the same series of metals[64]. Structural effects come into light as single crystal faces are used as electrodes. Potentials of zero charge change with crystal face orientation mainly because of a change in the work function of the face[2]. But the difference in potential of zero charge between the (I 11) and the (110) face of face-centered cubic metals (Cu[68], Ag[69], Au[70]) appears to always exceed the difference in electron work function (Table 3). This -* If may be thought that the comparison between the gas
$,(Vl
Fig. 16. Corrected Tafel plots for the reduction of 2 x 10e4 M NaBrO, in aqueous solutions of various NaF concentrations (different symbols, from 5.2 x 10-j to 0.1 M) at different metal electrodes. Reprmted from[541 with per-
mission by Plenum Press. the free energy change from the most hydrophilic to the most hydrophobic metal amounts to only 4&XkJ mol- I. This kind of evidence is more ambiguous with dmetals[64]. The heat of adsorption of hydrogen can be Teen (Fig. 21) to increase m the gas phase whereas it
phase and the electrochemical data were to be made at the potential of zero charge to reproduce the apparent situation of uncharged surface in both cases. But in fact this approach would be physically ambiguous. First of all, for most of these metals the potential of zero charge is more negative than the reversible potential of the hydrogen reaction[65]. Thus, the heat of adsorption would be measured at coverages close to saturation thus including lateral interaction effects which are absent in the gas phase at H + 0. Second, as hydrogen is adsorbed on a bare metal surface, charge transfer occurs also in the gas phase to an extent which depends on the work function of the metal Therefore, the surface is really uncharged in the gas phase, and the work function change (ie the potential in the electrochemical situation) is different for different metals.
Table 3. Potential ofzcro charge and electron work function of smgle crystal facts of face-centered cubic metals E,_,(V,
Metal _._~._.___ Ag* AU Cu
(111, -046 0.58 -037
n/w)
Ref. (110) (100) --..--.__...____.__~__ -0.61 -077 69 0.35 019 70 -il.46 -063 68
A[(lll)-(IlO)] AE, ,,(V) A@(eV) 0.31 0.39 030
0 22 0 19 0 24
* The values of the potenttal of 7ero charge include a small ascertained specific adsorption of F-[76. 771.
Ref. 78,79 80 81
effect due to
Fig 17 Model oftheproposed strucrllreofwateradsorhed on transition metalsurfaces basedonan ~rstudy. Water molecules are hydrogen-bonded in hexagonal rings. Three mnlecules over six have one of their lonepair orhitals directed towards the metal surface. Solid segments represent hydrogen atums (LIfthe sketch on the right-hand side). Reprinted from[29] with permissmn by the North-Holland Publishing Company.
Aspects of electrode/solution
interface
1091
Fig. 18. Sketch of a tridimensional network of hydrogenbonded watermolecules terminating at a plane metal surface, resulting in a multi-state model for adsorbed water. Reprinted from[l3] with permission by Elsevier Sequoia S.A.
E,
/V
Lrhe)
Fig. 21. Heat of adsorption of hydrogen in the gas phase (0) solution (0) of not specifically adsorbed electrolytes as a function of the potential of incipient hydrogen deposition. Data from[2,67]. E, is a parameter whose unique function is to linearise the plots. The verbcal distance between the points for the same given metal measures the incidence of electronic, solvent and structural effects characteristic of the electrochemical situation.
and in
Fig. 19. Average orientation of water molecules on transition metal surfaces as resulting from the application of the Helmholtz formula ys(dip) = r~,,,~ x P”,,J x sin O/s, wnh yS(dtp) = 0.4 V (cfFlg. 9 with 0.07 V for the absolute value of the surface potential of water at Hg), ,tn,o = 10’ 5 molecule cm 2 pt, o = 6.12 x 10es3 Cm, and e = 6. With a lower value of a,*the measured &dip) could be consistent with the model depicted in Fig, 1’7,smce not all of the molecules would be required to possess the maximum orientation shown here.
indicates that the electronic effect and/or the specific orientation of water molecules vary with the crystal face[71]. In principle, since the electron work function measures the electronegativity of the metal surface, a stronger orientation is expected on the face with the lower work function. Thus, the affinity for oxygen is also expected to be face specific as any other chemical property. The same is true for the electronic effect. However, the picture of the metal-water interaction is presently less clear for sd-metals than for the other two groups[24,72,73]. It may be of interest to extend this analysis to solid metals other than silver in order to assess how much the situation at a polycrystalline surface may be complicated. Experimental data (Table 4) are scarce but it is possible to see that the “inhomogeneity” of double layer parameters shows a trend to increase with the melting point[73], ie structural effects are important only with really “frozen” surfaces. Low-melting metals exhibit a not negligible atomic self-diffusion even at room temperatureC741. Therefore, surfaces are somewhat “fluid’ and structural effects are smoothed down. Working with polycrystalline surfaces can still be meaningful with only very few low-melting metals,
Table 4. Range !O. Standard free energy of tt-pentanol adsorption from aqueous soluttons ot not specrficahy adsorbed electrolytes on various metals, as a function of the enthalpy of ibrmation of the oxtde MO assumed to be a relatrve measure of the affinity ofthe metal surface for water molecules.~l he point for Ag 15in brackets because the value of AGY, has been extrapolated from data for n-hexanol. n-Pentanol is seen to be more adsorbed on less hydrophilic metals. Data of AG:, from[4].
Metal Bi Pb Ln Ag
of potential of mro charge crystal faces of metals ~~___~ Meltmg pomt ( Cl 271 327 420 961
0.05 0.04 0.15 0.31
for ditTerent ..-
82 83 84 69
1092
SERGIOTRASATTI
but it is certainly much less useful and sometimes misleading with Cu, Ag and Au. The situation is even worse with d-metals whose degree of surface inhomogeneity is very high. Acknowfedyement-Financial support National Research Council (CNR, acknowledged.
ta this work by the Rome) is gratefully
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