2769
J. phys. Chem. 1980, 84, 2769-2774
Temperature Dependence of the Surface Potential at the Mercury/Nonaqueous Solution Interface 2. Borkowska,t W. R. Fawcett;
and S. Anantawan
Guelph-Waterloo Centre for Greduate Work in Ctmmktry (c)ue$h Campus), Dspartment of Chemistry, UntversHy of Que@, Quelph, Ontarb, Canada N 10 2 W 1 (Recehred: February 22, 1980)
The temperature dependence of the surface potential at the mercury/nonaqueous solution interface has been estimated by measuring the temperature dependence of the potential of point of zero charge for the cells Hg/NaC10, in S/Na in Hg and Hg/KPF6 in S/K in Hg where S was methanol, dimethylformamide, N methylformamide,propylene carbonate, or dimethyl sulfoxide. The derived temperature Coefficients are positive for all solvents except methanol. The results are compared with existing data in the literature and discussed within the context of models for solvent structure at polarizable interfaces. It is concluded that the temperature coefficient of the surface potential cannot be used to determine the predominant orientation of solvent dipoles at the point of zero charge.
Introduction Although solvent structure at polarizable interfaces has been the subject of many experimental the preferred orientation of solvent molecules at the uncharged interface is often a subject of controversy. Two experimental methods are generally used to investigate solvent structure at the metal/electrolyte solution interface. In one case the capacity of the inner part of the double layer is estimated by studying the differential capacity of the interface or interfacial tension in the absence of ionicspecific adsorption from the solution. This method is especially informative when data are obtained as a function of temperature. The second method involves estimation of the temperature coefficient of the surface potential at the uncharged interface from data for the dependence of the potential of zero charge (pzc) on temperature. This analysis was first carried out by Randles and Whiteley4 to examine the properties of the mercury/aqueous solution interface. Subsequent studies have involved both aqueous6i6and nonaqueous solutions.&* It has been customary to interpret the estimates of the temperature coefficient of the surface potential on the basis of molecular models for the solvent at the polarizable metal electrode. In the simplest treatment of interfacial solvent structure, the solvent molecules are assumed to have two orientations at the interface, one in which the positive end of the solvent dipole points toward the metal and another with the negative end toward the metal?JO If the pitively oriented dipoles predominate at an uncharged surface, then the surface potential measured across a dipole monolayer from the metal to the solution is positive. Furthermore, as temperature increases, the net preferred orientation decreases and the temperature coefficient of the surface potential is negative. Accordingly, the positive temperature coefficients observed for the mercury/aqueous solution interface4* have been interpreted as evidence that water molecules are preferentially adsorbed on mercury so that the oxygen atoms interact with the metal. Recently, more detailed models for interfacial solvent structure have appeared including a four-state model applicable to associated solvents such as water" and a three-state model which describes the dielectric properties of unassociated solvents at charged boundaries.12 Interpretation of the temperature coefficient of the surface On leave from the Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland. 0022-3654/80/2084-2769$0 1.OO/O
potential on the basis of these models is much more complex, since the concentration of the additional components of the solvent monolayer, clusters in the case of associated solvents" and molecules with their dipole vectors parallel to the surface in the case of unassociated solvents,12also change with temperature. The purpose of the present study was to obtain data for the potential dependence of the surface potential in a wider variety of solvents both protic and aprotic. The solvents were chosen so that representatives of each of the three groups considered by Parsons2 were included. In addition the experiments were designed to avoid ions which are strongly specifically adsorbed at mercury from the chosen solvents.
Experimental Section The pzc at the Hg/solution interface was determined in cells of the configuration HglNaClO,SIHg,Na
(1)
where S represents one of five different solvents: methanol (MeOH), N-methylformamide (NMF), N,N-dimethylformamide (DMF), propylene carbonate (PC), and dimethyl sulfoxide (Me&O). The mercury electrode was a steaming electrode as described by Grahame for pzc det e r m i n a t i o n ~and ~ ~ the amalgam electrode, an all glass dropping electrode similar to those used by Ma~1nnes.l~ The p~tentialdifferences were measured with a precision of 0.1 mV by using a Wavetek Model 201 null voltmeter of high input impedance. Measurements were also made in cells containing KPF6 as electrolyte with a potassium amalgam reference electrode. Activity coefficients for the NaC104 electrolyte in the various solvents were estimated by measuring the electromotive force of the following cells: Na glass(%M NaC104,Slly M NaC104,SINaglass (2) The Na glass electrodes (Fisher Scientific) performed satisfactorily in all of the above sol~ents.'~The liquid junction in the above cell was formed in a wetted partially closed glass stopcock. Measurements were made with a precision of 2 mV for concentrations differing by a factor of 2 in the range 10-4-0.1M. High-purity solvents were prepared by drying the reagent-grade material over 3-A molecular sieves and distilling twice. In the case of MeOH, the distillation was carried out at atmospheric pressure under nitrogen. All 0 1980 American Chemical Society
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Borkowska et al.
The Journal of Fhysical Chemisfty, Vol. 84, No. 21, 1980
other solvents were distilled at reduced pressure, the temperature at the top of the column being as follows: 57 "C for NMF, 30 "C for DMF, and 60 OC for PC and Me2S0. The purity of the solvents was checked by gas chromatography by using a GE SF96 silicone column, 1.8 m long. An organic impurity in PC, assumed to be propylene oxide, could easily be removed by placing the distilled solvent under a vacuum at room temperature. The water content of the purified solvents as determined by Coulometric Karl Fischer titration was -0.01% except in the case of DMF where it was 0.05% . Traces of water were removed from DMF by passing the amalgam through the solution for 15 min before making any potentiometric determinations. Na and K amalgams were prepared by electrolysis from saturated NaC104and KPF6solutions, respectively, in PC and stored in a controlled-atmosphere chamber (Vacuum Atmospheres). The alkali metal concentration was determined by reacting a known weight of amalgam with water at -40 "C with air bubbling through the solution, and titrating the resulting hydroxide with standardized dilute HC1. High-purity electrolytes were prepared by recrystallizaticn of reagent-grade salts. NaC104 was twice recrystallized froin triply distilled water. In the case of KPF6, a recrystallization from water was followed by one from acetone. These salts were dried over P205at 100 "C and at reduced pressure. All measurements and solution preparations were carried out in the controlled-atmosphere chamber under nitrogen. Residual oxygen was removed from the solutions by bubbling purified N2 through them. The temperature of the system was kept constant by pumping methanol from a thermostat (Brinkman) with a precision of 0.1 "C through an external jacket surrounding the cell.
-
Results and Discussion Activity Coefficient Measurements. The emf of the concentration cell with transport using specific Na+ ion electrodes (cell 2) was determined for at least nine different NaC104 concentrations in the left-hand compartment. Analysis of these data requires transport number data, the thermodynamic relationship for the cell emf being (3) where a&) is the mean activity of NaC104 at a concentration of i M and t.. is the mean transport number for the anion in the two solutions. Since the transport-number data for this salt are not available in the nonaqueous solvents considered, two extra thermodynamic assumptions were made in obtaining estimates of the mean activity coefficientsf*. In particular, it was assumed that the mean activity coefficients are given by the extended DebyeHQckel theory and that the mean anionic transport number is a linear function of the square root of mean salt con~entration.'~The mean ionic activity coefficient for a 1-1 electrolyte is then given by eq 4, where c is the (4) electrolyte concentration (molarity), A and B are the Debye-Huckel constants, and A is the ion size parameter; the relationship between t- and the average electrolyte concentration c (C = [ c ( x ) + c(y)]/2) is t- = to- + bell2 (5) where to is the anionic transport number at infinite di-
1
I
c
07
/
6
0O 5
-
io* /*
01
F/0mole 2
1 i ' ) 2I
03
04
Flgwe 1. Varialbn in anionic transport number f- with the square root of sen commtrat&mclMfor NacIo, h a varlety of mqueous sohmts: (X) MeOH; (0) DMF; (A) PC; and (0)NMF.
TABLE I : Parameters Describing the Variation in the Mean Activity Coefficient f* for NaClO, with Salt Concentration According to the Extended Debye-Huckel Theory (Eq 4 ) in a Variety of Nonaqueous Solvents at 26 C A, dm3l2 E , dm3 solvent mol-l/z nm-l a, 0 nm MeOH 1.895 5.09 0.60 DMF 1.594 4.81 0.66 NMF 0.1443 2.16 0.48a PC 0.686 3.63 0.70 4.26 0.70 1.108 Me,SO a The voalue off* is extremely insensitive to the assumed value of a ; the value quoted here is that estimated by Povarov et a1.21
lution and b is a constant. Having chosen a value of A, one may calculate f k , a*, and thus a value of t- on the basis of the experimental values of 6 (eq 3). The best value of A was assumed to be that which yielded values of t- which were a linear function of (Figure 1). The precision with which A is estimated by this technique is A0.03 nm, values of A outside of this range resulting in t- against C'/' plots which deviate significantly from linearity. The values of the parameters giving activity coefficients for NaC104 according to eq 4 in MeOH, DMF, NMF, PC, and M e 8 0 a t 25 "C are summarized in Table I. The values of the ion size parameter reported in Table I appear quite reasonable and compare favorably with values estimated from Stoke's radii where these data are available.1g~20 In the case of NMF, the estimated activity coefficients are extremely insensitive to the chosen value of A because of the high dielectric constant of the solvent. In this case, A was assumed to be the same as the value found earlier for NaCl in NMF (0.48 nm).'l
Temperature Dependence of the Surface Potential Temperature Dependence of the Surface Potential. Following the analysis of Randles and W h i t e l e ~ ,the ~ temperature coefficient of the potential drop across the mercury/solution interface at the pzc, dArp0/dT, is related to the temperature coefficient of the potential difference across cell 1 by eq 6, where SNa(A) is the partial molar
entropy of Na in the amalgam, ~ N ~ + ( is s ) that for Na+ in solution, and 3,(A) is that for electrons in the amalgam. I t has been customary in treating these data to assume that single ion activity coefficients can be replaced by mean ionic activity coefficients and that S+) is negligibleq4 Partial molar ionic entropies are approximately given by eq 7 in dilute solutions, where Soi is the standard partial
molar ionic entropy of a 1m solution and p is the density of the solvent. Values of goi were taken from the compilation of Criss and Salomon.22 Activity coefficients for NaClO, in DMF, MeOH, NMF, and PC were those determined in the present study; those for NaC104in M@O were assumed to be equal to values of ft reported earlier for KPF6 in the same solvent.23 Finally, the temperature coefficient of the ionic activity was assumed to be negligible at constant salt concentration. The partial molar entropy for Na in the amalgam was calculated from data tabulated in the literature%according to the relationship
= SONa + A5Na(A) = S " N ~ R In X N + ~ A~'N~(A)(8) where SONa is the molar entropy of liquid Na at 25 "C, A5Na(A) is the relative partial molar entropy of Na in the amalgam, AYNa(A) is the excess relative partial molar entropy, and x N a is the mole fraction of Na in the amalgam. In the case that cell 1was constructed with K amalgam and KPF6 in the electrolyte solution, the values of the mean salt activity coefficient were taken from data in the literature. Values of ft for KPF in MezSO and PC were reported by Hills and Reeves.2s In the case of DMF, MeOH, and NMF, f* for KPFe was assumed to be the same as that for NaC104. The partial molar entropy for K in the amalgam was estimated from literature data" by using an equation corresponding to eq 8 with K replacing Na. However, in this case the concentration dependence of A ~ * K ( A )has not been established with precision; thus, ASK(A) was assumed to be independent of concentration.% Because of the uncertainty in fi and the estimates of dA40/dT for the potassium systems are considered to be less reliable than those for the sodium systems. Finally, it should be noted that the estimated values of dAqjo/dT are quite insensitive to the values chosen for f* Therefore, the fact that the activity coefficient data for the potassium salts were not obtained independently is not of quantitative importance. The precision with which the measured potential difference across cell 1 8 and its temperature coefficient are corrected for variation in salt and amalgam activity may be assessed by comparing estimates of t$O, the standard potential difference (at 25 "C and with the concentrations of salt and amalgam equal to 1 m). At the pzc, the standard potential drop is given by eq 9, where gois the
SNa(A)
potential drop a t the pzc, aNa+(s) is the activity of the Na+
The Journal of Physical Chembtty, Vol. 84, No. 21, 1980 2771
TABLE I1 : Summary of Thermodynamic Data for Cell 1 at Pzc and Derived Temperature Coefficients of the Surface Potential at the Hg/S Interface salt amalgam concn, concn, dA@,/ mol mole d8 Jd? dT, tio0,O solvent L-' fraction mV K' mVK-' V Hg/NaC10, in S/Na in Hg MeOH 0.049 0.0077 -0.734 i 0.004 -0.16 1.700 DMF 0.100 0.0024 -0.386 i 0.004 0.50 1.877 0.100 0.0008 -0.451 i 0.003 0.53 1.883 0.057 0.0008 -0.446 i 0.003 0.49 1.880 NMF 0.100 0.0042 -0.222i 0.002 0.22 1.795 0,100 0.0036 -0.238 i 0.003 0.22 1.791 0.100 0.0026 -0.254 i 0.003 0.23 1.792 PC 0.100 0.0042 -0.460 i 0.005 0.19 1.667 0.050 0.0036 -0.426 i 0.003 0.17 1.671 Me,SO 0.056 0.0027 -0.276 i 0.002 0.49 1.852 Hg/KPF, in S/K in Hg 0.12 1.721 MeOH 0.049 0.0069 -0.154 f 0.003 DMF 0.050 0.0065 -0.356 i 0.002 0.55 1.870 NMF .0.049 0.0078 0.213 i 0.002 0.36 1.784 0.23 1.784 PC 0.10 0.0085 -0.079 i 0.003 a The standard state is 1 m concentration for both the salt in solution and the metal in the amalgam;the temperature is 25 "C.
ion in solution, and aNa(A) is that for Na in the amalgam, concentrations being expressed as molality. The activity coefficients for the alkali metals in the amalgam were taken from the tabulation of Korshunov et aleB The data obtained for both the Na and K systems in the various solvents are summarized in Table 11. In cases where measurements were made at various salt and amalgam concentrations, good agreement is found between estimated values of dAdo/dT and Eo. These results indicate the precision of the experiments and the validity of the assumptions made in estimating the concentration-dependent terms involved in the analysis. In the case of the Na system, dA4,/dT is positive for all solvents except MeOH. For the K system, dAbo/dT is always slightly more positive than for the Na system and, in the case of MeOH, has a different sign. This difference is attributed to the uncertainty in the partial molar entropy for K in the amalgam,24but it could also be partly due to the fact that activity coefficient data for KPF6 were not available for most of the solvents considered. For these reasons, the results for the K system are considered much less reliable. The results obtained here for methanol may be compared with those reported earlier by Garnish and Parsons7 and Koczorowski and Figaszewski? The calculated value" of dA40/dT based on data for dilute LiCl solutions in MeOH with a calomel reference electrode7is -0.4 mV K-' mol-'. In the case of NaCl solutions in the concentration range 0.02-0.1 M,s dA4,/dT varied from -0.074 to -0.28 mV K mol-'. Although the specific adsorption of C1- at the Hg electrode is expected to affect these results somewhat? they support the above results obtained with different reference electrode systems. The mean value of dA40/dT for the Hg/MeOH interface, if one considers all results with the assumption of equal precision and ignores the role of ionic specific adsorption, is -0.15 f 0.13 mV K-' mol-'. Although the scatter is large, it is clear that the temperature coefficient of the surface potential is negative for this interface. The other nonaqueous solvent for which data are reported in the literature8 is formamide (F);in this case the recalculated value" of dA40/dTis 1.0 mV K-' mol-'. This solvent has the largest value of the temperature coefficient, its magnitude being considerably greater than that for
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The Journal of Physical Chemistry, Vol. 84, No. 21, 1980
NMF (0.2 mV K-’mol-’). F, IWF, and water are strongly associated solvents in the bulk. This feature is probably reflected to some extent at the phase boundary, but its variation with the nature of the molecule and its capacity for hydrogen bonding are difficult to assess. Soluent-Molecule Orientation at the Phase Boundary. The analysis of interfacial capacity data for a variety of solvents in the absence of ionic-specific adsorption has led to the conclusion that features of the inner-layer capacity against charge-density curves can be attributed to solvent properties in the inner In the molecular treatment of the dielectric properties of the inner layer, it is assumed to consist of a monolayer or solvent molecules represented as polarizable hard spheres with permanent dipole vectors oriented in selected directions with respect to the phase boundary. The simple models which can be applied to aprotic solvents for which intermolecular hydrogen bonding is unimportant, and their application to estimating dA40/dT are now considered. According to the two-state model as detailed by Levine, Bell, and Smith,27the molecules in the monolayer may occupy two orientations with respect to the electrode’s field, namely, “up” with the positive end of the dipole vector pointing towards the metal, and “down” with the negative end towards the metal. The resulting expressionB for the potential drop across the monolayer is3lz7eq 10, Ce
where u is the charge density on the metal, d is the diameter of a solvent molecule and, thus, the thickness of the monolayer, X is the reaction field due to dipole orientation in the direction of the electrode’s field, c, is an effective coordination number for a given dipole with its nearest neighbors in the monolayer and dipole images in the electrode and diffuse layer, and NT= 2/3lI2d2,the number density of solvent molecules in the hexagonally closepacked monolayer. The reaction field is given by eq 11,3”
X=
cebV1 - f2)
- 4ra11~1
d3 + c , ~ I /
(11)
where fl is the fraction of molecules in the “up” direction, f2 that in the “down” direction, p the magnitude of the molecular dipole moment, and a11ita polarizability along the direction of the dipole vector. Since the monolayer is completely filled by solvent molecules in the two orientations (12) f l + fz = 1 Combining eq 10-12 and differentiating with respect to temperature a t constant electrode charge density, one obtains
It is readily apparent that, in this simple case, the sign of dA40/dT reflects the sign of dfl/dT at the pzc. Thus, when dA40/dT is positive, df,/dT is positive and the number of “up” dipoles increases and that of “down” dipoles decreases with an increase in temperature. This situation is expected when “down” dipoles predominate at the pu: at lower temperatures and the surface potential is negative. Considering the variation in dipole orientation with electrode charge density at constant temperature, one expects that the numbers of molecules in the two orientations will be equal at a sufficiently negative charge
Borkowska et at.
density. Examination of the predictions of the two-state model with respect to the differential capacity of the monolayer reveals that a capacity maximum occura at the charge density where fl = fz and that this is the only extremum on the capacity curve? MezSO,DMF, and PC all possess capacity maxima a t positive charge densities: shallow minima also being observed at negative charge densities in the case of PC30 and DMF.3’ If one assumes that the minima can be attributed to another physical effect such as electrostriction,B the presence of capacity maxima at positive charge densities requires that fl > f2 at the pzc and that dAdo/dT < 0. Thus, the observed maximum on the inner-layer capacity curves and the temperature coefficient of the surface potential cannot be explained on the basis of the two-state model. A more satisfactory description of interfacial capacity behavior in aprotic solvents is provided by a three-state model12in which a third state involving solvent molecules with their dipole vectors parallel to the interface is added to the model described above. The expression for the potential drop across the monolayer (eq 12) remains the same, whereas that for the reaction field X becomes
where q = aIl(fi + fz) + alf3, f3 being the fraction of molecules in the monolayer with a “parallel” orientation and a, being the dipole polarizability in a direction perpendicular to the dipole vector. The expression for complete coverage is
fi + fz + f 3 = 1
(15)
If one combines 10 and 14, it is easily shown that
d3 + c,q
Two limiting cases may be distinguished with respect to the interfacial capacity curve8 predicted by the three-state m ~ d e l . ~When J ~ the surface concentration of the “parallel” orientation is low, the capacity curves are similar to those predicted by the two-state model with a maximum occurring at the charge density where fl = f2. In this case one may assume that df3/dT = 0 at the pzc. Equation 16 then reduces to eq 13, and the predictions of the threestate model with respect to the temperature dependence of the surface potential are identical with those of the two-state model. On the other hand, when the surface concentration of “parallel” dipoles is high, the capacity curve has three extrema, a minimum at the charge density where fl fz and maxima when fl = f3 and fz N f3.”12 When the minimum occurs at far-negative charge densities, one may assume that fl and dfl/dT are both negligible a t the pzc. It follows that (dfZ/dT), = -(df3/dT),,, and the temperature coefficient of the surface potential is given by eq 17,where
the subscript “0”refers to the pzc. When the maximum occurs at positive charge densities and anisotropy of the polarizability can be neglected (all N aL),dipoles in the “parallel” orientation predominate at the pzc and (dfi/dT) is positive. Thus, the three-state model predicta a negative value of dA40/dT under these circumstances (eq 17). This
Temperature Dependence of the Surface Potential
situation is apparently that for the Hg/MeOH However, the majority of aprotic solvents which poasess interfacial capacity curves with a minimum a t negative charge densities and a maximum at positive charges have positive temperature coefficients of the surface potential (Table 11). The latter situation can be explained by the three-state model when the solvent molecule is characterized by signifcant anisotropy of its polarizability. When q > al,the capacity maximum at positive charge densities occurs at more positive charge densities than the point where f2 = f3, that is, at a charge density where f i > f p It follows that the surface concentration of “down” dipoles can be greater than that of “parallel” dipoles, and, therefore, (df2/dr), can be negative at the pzc. According to eq 17, one then expects to observe a positive temperature coefficient of the surface potential. In fact, experimental values of the Kerr constants for a variety of amide solvents including DMF determined as solutes in dioxane at 25 0C32show that the component of the polarizability tensor along the dipolar axis is greater than those in perpendicular directions. In conclusion, the three-state model can account for both positive and negative values of dAdo/dT in the cases of aprotic solvents possessing interfacial capacity curves with the same general characteristics. However, the sign of the surface potential Ado is always negative, the sign of dAbo/dT being determined by whether dipoles in the “down” or “parallel” orientation predominate at the pzc. Therefore, in general, one may not draw a conclusion regarding the sign of A40 on the basis of the experimental value of dAdo/dT. A similar conclusion was reached by Nedermeijer-Denessen and de Limy38 on the basis of an analysis of estimates of the surface potential at the solvent/air interface and its temperature coefficient. The analysis was based on the twostate model of Levine, Bell, and Smith27but with the assumption that the dipole vectors in the two states are not equal in magnitude. The situation with protic solvents such as water, F, and NMF is much more complex. Values of dAdo/dT are positive for all solvents, the largest coefficient being that for F 8 p B (1.0 mV K-l) and the smallest that for NMF (0.2 mV K-l). According to the cluster models used to describe the interfacial behavior of associated so1vents,11*33,34 one should consider molecular associates with dipolar properties determined by intermolecular hydrogen bonding. Although simpler cluster models which ignore dipole-dipole interactions” or treat them in an approximate fashion%are able to account for the inner-layer capacity observed at the Hg/H20 interface, they cannot account for the positive value of dAdo/dT (0.6 mV K-1).4 However, when one treats dipole-dipole interactions in detail,% all electrostatic properties of the inner layer are explained. At 25 “C, the calculations predict that the surface monolayer is composed of 15% monomers in the “up” orientation, 20% monomers in the “down” orientation, 46% clusters “up”, and 19% clusters ‘bdown”at the pzc. Since the permanent dipole moment of a cluster is assumed to be small (0.53 D) compared to that of a monomer (1.84 D), the net polarization of the monolayer is negative, leading to a predicted value of A40 = -17 mV.34 As temperature increases, there is a tendency for clusters to break up and for monomers to reorient from the predominant “down” orientation to the “up” orientation. When dipole-dipole interactions are considered in detail, the latter effect predominates and the net polarization remains negative but reduced in magnitude, thereby leading to a positive value of dAdo/dT. When one considers the facts that dAdo/dT is positive and a capacity maximum is observed
The Journal of Physkal Chembtry, Vol. 84, No. 21, 1980 277a
a t negative potentials for both F and NMF’, the distribution of solvent species at the Hg interface is probably qualitatively the same as for the Hg/H20 interface. However, because of the complex nature of theae systems, a detailed analysis of interfacial dielectric properties should be made before these data are assessed. It should be noted that the analysis presented here, which is based on that originally used by Randles and Whiteley,’ has been questioned by T r a ~ a t t i .More ~~~~ specifically, Trasatti has pointed out that the work function of metals is not independent of temperature on the basis of available data.3BIt follows that the partial molar entropy of the electron, S,, is not zero. Unfortunately, data for the temperature dependence of the work function of Hg are not available in the literature; however, if the temperature coefficient is nonzero, the interpretation of dAdo/dT in terms of the models discussed above would obviously be different. In conclusion, it is clear that the simple two-state model used to interpret interfacial-capacity data, together with experimental estimates of dAdo/dT, is completely inadequate. It follows that one cannot reach a conclusion regarding the sign of the surface potential by determining dAdo/dT. Additional experimental evidence is necessary to assess solvent behavior at a polarizable interface, specifically, the dependence of inner-layer capacity on temperat~re.~~J+=
Acknowledgment. We thank Dr. Roger Parsons for helpful discussions related to this work. The financial assistance of the Natural Science and Engineering Research Council of Canada is gratefully acknowledged. References and Notes 8. E. Damasktn, 0. A. Petry, and V. V. Batrakov, “The Adsorptkn of Organic Compounds on Electrodes”, Plenum Press, New Yorlc, 1971, Chapter 1. R. Parsons, Electrmhlm. Acta, 21, 681 (1976). W. R. Fawcett, Isr. J. Chem., 18, 3 (1979). J. E. 8. Randles and K. S. Wh te le l y, Trans. Faraday Soc.. 52, 1509 (1956). Woonkle Pelk, T. N. Anderson, and H. E m , J. Fbys. Chem., 71, 1691 (1967). 2. Koczorowskland Z. Flgeszewski, Rocz. chem.,44, 191 (1970). J. D. Qarnish and R. Persons, Trans. Feredey Soc., 63, 1764 (1967). R. Parsons, J. Eklroanal. Chem., 53, 479 (1974). R. J. Watts-Tobln. PnH4s. Mea.. 6. 133 119611. J. OM. ECCMS, M. A. v. Devaiahin, ad^. ~ h ROC. , a. L d n , Ser. A , 274, 55 (1963). R. Parsons, J. Elecbpenal. Chem., 59, 229 (1975). W. R. Fawcett, J . MYS. Chem.. 82, 1385 (1978). D. C. Qahame, R. P. krsen, and M. A. Poth,. J. Am. Chem. Sac., 71, 2978 (1949). D. A. MacInnes, “The Rhciples of Electrochemlsby”,Dover, New York, 1961, Chapter 8. NaglasselecbodeshevebeentestedprevbuslyandBhownto1pve SattsfactOrily in MeOH,”’ DMF,” Me,S0,17 PC,“ and NMF. J. Tarasrewska. J. Eklroanal. Chem., 49, 443 (1974). J. E. McChne and 1. 8. Reddy, Anal. Chem., 40, 2064 (1968). D. KomJnska, Doctoral Thesis, InetihRe of phvelcal awwnleby, polsh Academy of S-ces, Warsaw, 1977. R. FemanderPrlnl In A. K. Covington and T. Mdtlnson, “Physical Chemktry of Organic solvent Systems,” plenum Press, New Yorlc, 1973. -. Chanter - - r - - 5. -M. Matsuva. K. umemoto, and Y. Takeda. Bull. Chem. Soc. &n., 48, 2253 (1975). Yu. M. Povarov, Yu. M. Kessler, A. I. Gorbanev, and I.V. Sefmva, Dokl. Akad. Narrk SSSR.155. 1411 119641. c. M. crlss and M. sakmon in'^. R. &i&itm and T. DICM~WI, “Physical chemtatry of Organlc Solvent Systems”, Plenum Press, New York, 1973, p 286. 0. J. H#ts and R. M. Reeves, J . Ekclroanal. Chem., 38, 1 (1972). R. Huttgren, R. L. Orr, P. D. Anderson, and K. K. Kelky, “Selected Values of lhwrmdynarnic Ropertles of Metals and Alloys”,Why, New York, 1963. V. N . K o r s h u n o v , A . B . ~ w , a n d IP..-, ,6, 1204 (1970). QamMand Parsons’ estmated dAd J d Tfor the WMeoH krterface
a.
(23) (24) (25) (26)
2774
J. Phys. Chem. 1980, 84, 2774-2779
molar entropy for Hg, So used In the calculation was that for the gaseous metal (174.8 J Kq mor'), the appropriate value being that for the HquM metal (76.0 J K-' mor'). As a result, the value of dAq5ddTrepated in ref 7 Is high by 1.02 mV K-' md-', the conect value being -0.40 mV K-' mot'. A similar error was made in ref 8 where the gasqhase molar entropy of Ag (182.9 J K-' md-')was used instead of that for the sdkl state (42.55 J K-' mol-'). The corrected estimate of dAq5 d d Tfor fwmamkle on the basis of the data discussed is 1.0 mV K-I mol-' S. Levine, G. M. MI, and A. L. s m k , J. phys. C b m . , 73,3534 (1969). Although a more general case was consklered In the odginal paper of L e h e , Bell, and ~mlth,*'the version presented hare assumes that the fn@tude of the dipole moments in the two orientations are qUei. J. R. Macdonald, J . chem. phys., 22, 1857 (1954).
(30) Nguyen Huu Cuong, A. Jenard, and H. D. Hurwltz, J . Electroanel. Chem., 103, 399 (1979). (31) W. R. Fawcett, 6. M. Ikeda, and J. 6. Sellan, Can. J. Chem., 67, 2268 (1979). (32) M. J. Aroney, R. J. W. LeFBvre, and A. N. Singh, J. Chem. Soc., 3179 (1965). (33) 6. B. Damaskh and A. N. Frurw1,€&ctmdh. Acta, 19, 173 (1974); 6. B. Damaskin, J . Electrosnal. Chem., 75, 359 (1977). (34) W. R. Fawcett, S. Levine, R. M. delrloklge, and A. C. McDoneld, J. Ekfroenal. Chem., in press. (35) 2.Borkowska and W. R. Fawcett, EleMrokhlmiye, in press. (36) H.J. M. NedermeiJer-DeNessenand C. L. deligny, J. E l e c f r a e ~ l . Chem., 59, 1 (1975). (37) S. Trasatti, J . Ekfroansl. Chem., 33, 351 (1971). (38) S. Trasatti, J. Electroenel. Chem., 82, 391 (1977). (39) J. C. Rivlbe, Sol@ Sfate Surf. Scl., 1, 180 (1969).
Ultrahigh-Vacuum Techniques in the Measurement of Contact Angles. 5. LEED Study of the Effect of Structure on the Wettability of Graphite'** Malcolm E. Schrader David W. T a y b Naval Sh@ Research and Devebpmnf Center, Annapolis, Merylend 21402 (Received: November 5, 1979; In Final F m : June 16, 1980)
It was previously found that bakeout and ultrahigh evacuation of the (OOO1) plane of oriented graphite produced a clean surface (as determined by AES) which yielded a water contact angle of -35". Ion bombardment of that evacuated surface reduced the water contact angle to 0". The present work seeks to determine whether the ion bombardment removed undetectable (to AES) residual contamination or merely disordered an already clean surface. The (OOO1) plane of ZYE3 oriented graphite was examined by LEED after vacuum heating to -800 "C, and the water contact angle measured in situ. The surface was then put through several cycles of ion bombardment, LEED analysis, and water contact angle measurement. The original heated surface showed LEED patterns characteristic of clean graphite (OOO1) and yielded water contact angles of 38 f 3". The LEED patterns gradually disappeared with increasing ion bombardment, accompanied by decreasing water contact angles. The water contact angle did not reach zero before the LEED pattern had completely disappeared. It is concluded that the contact angle of 38 f 3" represents a clean (OOO1) surface of !ZYB oriented graphite while the 0" contact angle results from formation on the surface of a disordered (amorphous) layer. The value of 38 f 3" found on ordered ZYB is not necessarily that for a perfect (OOO1) surface. The contact angle of the latter is estimated to be in the range of 42 f 7". The results are discussed in terms of values reported in the literature for the surface energy of graphite (Ooal), and the wettability of surface carbon in general.
Introduction The surfaces of carbon adsorbents have been regarded as high energy because of their efficacy as adsorbents for gases and vapors. Until recently, however, all reported contact angles for water on the surface of nominally elementary carbon have been greater than 0". Values for the graphite (OOO1) surface have generally ranged from 60 to 85" with considerable hystere~is.~-~ In recent work on oriented graphite, air-cleaved (O001) surfaces were cleaned by heating in ultrahigh vacuum.2 Auger electron spectroscopic (AES)measurements indicated that the surfaces were free of contaminants with the possible exception of hydrocarbons or chemisorbed hydrogen. The contact angle of water on this surface, measured in situ, was found to be ~ 3 5 When ~ . the surface was ion bombarded, however, water spread upon it spontaneously with a 0" contact angle. Evidence was presented that the high contact angles (60to 85") obtained when these measurements are made in air are due to carbonaceous contamination. It was not clear, however, whether the vacuum-baked surface yielding the 35" contact angle was clean, ordered graphite (OOO1) or a surface containing residual hydrocarbons (or chemisorbed hydrogen). If the former were true, the reduction in contact angle after ion bombardment would be the result of disordering of the (O001) structure, whereas if the latter
were true, the Oo contact angle obtained after ion bombardment could be the value characteristic of the clean, ordered (OOO1) surface. In this work, various types of graphite surface prepared in vacuum were examined by LEED (in addition to AES) for evidence as to which represented clean, ordered graphite (OOO1). Contact angles of water were measured in situ on each of the structures thus characterized.
Experimental Section Grade ZYB of a high-purity pyrolytic graphite called oriented graphite (Union Carbide) was used in these experiments. The sample was cleaved in air and mounted on a sample manipulator in a bell jar with porta for AES, electron bombardment heating, and ion bombardment as described previously? A stainless steel cold finger was wed to condense water vapor for in situ sessile drop contact angle measurements as described previously? In addition, a port with LEED capability was utilized in the present experiments. Results Effect of Ion Bombardment on LEED and Contact Angle (in situ), A ZYB oriented graphite surface which, after cleavage, had been tamped down with Kimwipe tissue
This article not subject to US. Copyright. Publkhed 1980 by the American Chemical Society
