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On the behaviour of molecules of the quinoline group at the water—Mercury interface
J. Electroanal. Chem., 123 (1981) 21--34 Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
21
ON THE BEHAVIOUR OF MOLECULES OF THE QUINOLINE GROUP AT THE WATER--MERCURY INTERFACE PART iI. ELECTROCAPILLARY STUDY OF ISOQUINOLINE IN SOLUTIONS CONTAINING ONLY AN "INERT" ELECTROLYTE
C1. BUESS-HERMAN, N. VANLAETHEM-MEUREE, G. QUARIN and L. G I E R S T
Facult4 des Sciences, Universitd Libre de Bruxelles, 50, avenue F.D. Roosevelt, 1050 Brussels (Belgium) (Received 15th December 1980)
ABSTRACT
Isoquinoline has been examined in a 0.5 M Na2SO4 aqueous solution, by using various electrocapillary methods. In the range of small potentials, and rather surprisingly, isoQ and Q present a very similar behaviour, despite the considerable difference in their dipole orientation with respect to the median carbons. Esin--Markov plots and capacity--potential-concentration maps indicate that there is gradual desorption towards negative potentials. Three distinct states (dilute flat molecules, a mixture of flat and erect molecules, and a m o n o l a y e r with a variable extent of local clusters) have been characterized and their domain of prevalence demarcated. At more negative potentials, a sudden reorientation is observed, which reflects the lack of miscibility between two superficial states internally stabilized by their own set of lateral interactions. The effects of the surfactant concentration, the electrolyte concentration and the t e m p e r a t u r e on the position of the transition potential are interpreted on the basis of the relative position of two 7 = f(E) curves, whose vertical shifts are controlled by the Gibbs isotherm equation, with two distinct discrete values for the superficial excesses corresponding to the two antagonizing monolayers. The most c o m p a c t of these monolayers is stable in the region with stretches between the phase transition potential and the electrolyte potential, both very sharply defined, its structure has been determined on the basis of a n u m b e r of converging facts, such as the limiting value of F, the values of the capacity and the position of the pzc, extrapolated from the o M = f(E) plots.
INTRODUCTION
In the first paper of this series [1], the adsorption of quinoline (Q) at the water--mercury interface has been investigated, by considering the case of a solution containing an inactive supporting electrolyte. Despite this major simplification, the behaviour remains rather involved. A number of distinct superficial states have been characterized as a function of the bulk concentration and the electrode potential, and connected to various orientations of the molecule. In order to check the validity of the conclusions which have been drawn, it is of interest to examine other molecules as equivalent as possible in terms of size, atomic arrangement and electrical properties, but differing by the orientation of their dipole moment. 0 0 2 2 - 0 7 2 8 / 8 1 / 0 0 0 0 - - 0 0 0 0 / $ 0 2 . 5 0 © Elsevier Sequoia S.A.
22
Isoquinoline (isoQ) meets all these requirements. Most of its properties (Table 1) are very similar to those of quinoline. However, the presence of the nitrogen atom in position 2 is now such that the dipole m o m e n t (2.6 D in vacuo) presents an angle of the order of 70 ° [2] with respect to the bond between the two median carbons (C4a--Csa), so that the vector components parallel and normal to this axis are respectively of 0.9 and 2.4 D (Table 2) (for Q, for which the angle is 2 °, the corresponding values were 2.2 and 0.1 D). This change in the direction of the dipole can be expected to promote new stable molecular orientations at the interface. Therefore, comparison between Q and isoQ affords a possible way of disentangling the effects simply related to the overall geometry of the two molecules from those which are more specifically controlled by the interactions between the dipole and the electrical field of the electrode. The most peculiar phenomenon presented by isoQ is the existence of a very sharp phase transition between two condensed layers, which must necessarily be of incompatible structures. The pioneering work of Lambert [ 3] has shown that the extent of inhibition exerted by isoQ on a number of electrode processes is considerably larger for the structure which prevails at potentials negative with respect to this transition. [As a typical example, the reduction of Tl * is characterized by an unaffected reversible wave leading to a fully developed diffusion current, as long as the transition potential ( E T ) iS not reached. As soon as it passed beyond this, the current drops to zero and maintains this value for more than 500 mV.] Lambert made the suggestion that with increasingly negative potentials, the isoQ molecules, which first stay flat, then tilt gradually towards a perpendicular position (with their carbons 4 and 5 adjacent to the surface as for Q) to assume finally, in an abrupt way, a third upright orientation with carbons 6--7 in contact with the metal. This model has been supported in recent work based on ellipsometry [ 4], which has also provided indications pointing to multilayer formation, for solutions close to saturation. The existence of a restructuration transition process raises interesting problems of a kinetic nature. Lambert and co-workers [5 ], using square-wave signals
TABLE 1 Some properties of
isoquinoline
S o l u b i l i t y in w a t e r 2 5 ° C
Boiling point Melting point pK a Refractive index
Dipole moment Dimensions
4.2 × 10 -2 M 243.2°C 26.5°C 5.4 1.62 2.6 D a = 0.7 n m b = 0.87 n m c = 0.3 n m
23 TABLE 2 Basic orientations of isoquinoline and derived quantities Carbons on the surface
Orientation
Area /nm 2
Thickhess /nm
101° Fmax /mol cm -2
Dipole m o m e n t / D Perpendicular to the surface
Parallel to the surface
All
Flat
0.61
0.3
2.7
--
--
4--5
Perpendicular "4--5"
0.26
0.7
6.4
0.88
2.44
5--6
Perpendicular "5--6"
0.267
0.73
6.2
1.99
1.67
Perpendicular "6--7"
0.21
0.87
7.9
2.44
0.88
6--7
of adequate amplitude, reached the conclusion that the process of rearrangement is basically of the "nucleation and g r o w t h " type, as indicated by sharp peaked i = f(t) transients presenting a marked hysteresis. These exploratory experiments were followed by the work of Jenard and Hurwitz [6], who investigated the major characteristics of the rising part of the cathodic chronopotentiometric transients. In these works, which were principally aimed at bringing out phenomenological facts, no a t t e m p t was made to interpret quantitatively the rate of the process in terms of a suitable model and an adequate formulation. The whole problem will be reconsidered in more detail in Part III of this series [ 7]. The present paper is strictly limited to the t h e r m o d y n a m i c aspects of the adsorption of isoQ in the presence of Na2SO4 which, as stated in Part I, can be considered (together with NaOH and NaF) as a suitable "inert" electrolyte in the sense that anionic coadsorption in the inner layer is essentially negligible. Potashnikov [8] has shown that isoQ, in the presence of Na2SO4, is affected by the salting-out effect. Our own determinations (Table 3) confirm quantitatively his results and indicate that NaF, NaOH and Na2SO4 act in a very similar way, in contrast with NaC104 which is notably less active, as is generally the case for organic solutes. Salting-out is easily accounted for by expressing the surfactant concentration with respect to its solubility in the particular electrolyte employed.
TABLE 3 Solubility of isoquinoline in various electrolytes [Na+]/M
Na2SO4
NaOH
NaF
NaC104
0.1 0.5 1.0
3.9 x 10-2M 3.1 x 10-2M 2.1 x 10-2M
3.8 X 10-2M 2.9 x 10-2M 2.0 x 10-2M
3.8 x 1 0 - 2 M 2.7 x 10-2M 2.0 x 10-~M
3.9 x 1 0 - 2 M 3.7 x I 0 - 2 M 3.3 x 10-2M
24 EXPERIMENTAL
Reagents Water, mercury and Na2SO4 were treated as described previously [ 1]. Here, isoQ (UCB " p u r u m " grade) was submitted to two successive fractional distillations under vacuum. The final colourless liquid was kept under nitrogen in a refrigerator, in order to slow down any possible degradation. The saturated concentration is 2.1 × 10 -2 M in ½ M Na2SO4 aqueous solutions kept at 25°C.
Instrumentation The various techniques which have been used to measure interfacial tensions, charge densities and differential capacities are identical to those already reported [ 1 ]. However, drop times at very negative potentials are poorly reproducible, so that superficial tensions have been computed in that range by integration of the charge density--potential plots. Electrode potentials are always referred to the saturated calomel electrode (SCE). RESULTS The adsorption of isoQ has been investigated under the same conditions as for Q, in a potential range ( 0 . 0 - 1 . 0 V) which is again limited here by the gradual onset of interfering faradaic processes involving the surfactant itself. Adsorption is evident from the progressive lowering of the interfacial tension 7 with increasing concentrations. Figure 1, which features a family of electrocapillary curves A'y = f(E), shows the existence of distortions at negative potentials for the most concentrated solutions. Complete desorption is not observable, since it occurs, in both directions, at potentials where the behaviour is no longer that of an ideally polarizable electrode. However, the desorption at negative charges can be detected as a distinct break, at medium concentrations of isoQ and in strongly alkaline solutions, for which the interference of reduction processes is less troublesome. The charge density--potential plots (Fig. 2) show that the potential of zero charge (pzc) shifts markedly towards less negative values with increasing concentration without reaching any defined limiting value. This trend is reversed for the most concentrated solutions at small negative charge densities {about f r o m - - 2 to --6 pC cm -2 ). The existence of an ill-defined nodal point suggests the presence of an extremum in the F = f(E) function. At potentials around --0.65 V, the charge densities tend to values which are independent of the bulk concentration. Qualitatively at least, all these experimental facts are not drastically different from those observed with Q. However, in striking contrast with the latter, a vertical step of more than 1.5 pC cm -2 disjoints the charge~potential plot. The discontinuity shifts towards more negative potentials with dilution, according to a function which is approximately logarithmic in concentration (Fig. 2, inset). Beyond the transition, the charge densities are essentially concentration independent until desorption is finally reached. The data on double-layer capacities, obtained by tensammetry, are presented
25
A ~ ' / m N . m -1
1
60
/
7.10 -~
/ , /
5.10
/ /
40
~
-~
3.10-~ 2.10-~
10 -1 7.10 -2 20 5 . 1 0 -2
3.10 -2
"
I
o.o
05
2.10 -2 ' 10-~ --~.~
1:o -E/v (SCE)
Fig. 1. Influence of the potential on the lowering of interfacial tensions in ½M Na2SO4 for the relative concentrations c/cs, as indicated.
under the form of isocapacity c o n t o u r curves, plotted as a function of the potential and the logarithm of the concentration (Fig. 3). At potentials close to 0.0 V, and within an extended concentration range, fairly large values of the capacity are observed. Two crest lines (thick solid lines), which are the locus of the capacitive peaks, have approximately the same origin at E = --0.25 V and log c/ cs = --0.55. When the isoQ concentration is gradually increased, the peak shifts first to less and then to more negative potentials, a fact which suggests a qualitative change in the properties of the surface state which prevails in the left-hand side of the diagram. At more negative potentials (zone II) the capacities decrease and reach a constant value of the order of 5.8 p F cm -2 , except for the m o s t c o n c e n t r a t e d solution. At even more negative potentials, a sudden discontinuity {broken line a in Fig. 3) demarcates an extended, new, distinct region (zone I) which is similarly characterized by a constant capacity of lower value (4.5 p F cm -2 ). Capacity and charge steps are coincident. The absence of any capacity peak, even at the lowest frequencies, indicates that the related kinetic process can be considered as infinitely slow, at least for a potential amplitude of +5 mV. Desorption occurs at negative potentials which are shifted negatively with increasing concentrations (curve/~). A triple point ( d o t t e d circle) is observed for a bulk concentration of the order of 5 × 10 -2 c, and at a potential of ~ 1 . 2 6 V. Values of the superficial excesses have been calculated by graphical differentiation of the F = f ( R T in c) plots, with an absolute standard deviation of the
26
/
/~JC.cm 2
/ -10 i
i
i i
t i
. i
b
-5-
/
p
t J
1 1
log c/c s
-1
/ /
/ I I
I oo
tb - E/v0.5
1D
-E/V(SCE)
-
Fig. 2. Plots of charge density vs. potential for the following relative isoQ concentrations: (1) 0.01; (2) 0.02, (3) 0.03; (4) 0.05; (5) 0.07; (6) 0.1; (7) 0.2; (8) 0.3; (9) 0.5, (10) 0.7; (11) 1. The broken line refers to the electrolyte alone. Dotted and dashed lines (a) and (b) show the extrapolation of the two ranges of common charge densities. A, B, C, I and II refer to the zones observed in Fig. 7. Inset: variation of the transition potential with the logarithm of the relative concentration.
order of 0.3 × 10 -'0 mol cm -2. For the sake of subsequent comparison [9], they have been plotted as a function of both the electrode potential (Fig. 4) and the charge density (Fig. 5a). The transition already observed on the charge and capacity curves appears here as a conspicuous step in the superficial excess, which increases discontinuously from about 6.4 + 0.3 X 10 -'o to 8.1 + 0.3 X 10 -~° mol cm -2, and depends very little on the concentration except for the saturated solution, for which it is slightly smaller. The experimental isotherms have been constructed by plotting the surface excess as a function of the logarithm of the bulk concentration, for several selected values of the potential (Fig. 5b) and the charge density (Fig. 5c). The accuracy and precision of all experimental data have been checked, by
27
iogC/c s
C
Y,
I 4.5
\
/
//3
\ /
\
-1
isoQ
/
/
( \?'~1
/
H20 18- 20
12
-2
40
30 25 20
15
O~
~:o
=
-E/V(SCE)
Fig. 3. I s o c a p a c i t y c o n t o u r map as a f u n c t i o n of the isoQ c o n c e n t r a t i o n . The two thick lines are the "crest lines". B r o k e n lines 0~ and/3 c o r r e s p o n d respectively to the t r a n s i t i o n and d e s o r p t i o n . T h e d o t t e d circle locates the triple point. A, B, C, I and Ii refer to the zones observed in Fig. 7.
101° ["/tool cm -2 10
I-
75
:I J
I.
9
5
2
2.5
0.0
O.5
1£)
-E/V(SCE)
Fig. 4. Superficial excess as a f u n c t i o n of the electrode p o t e n t i a l for the same relative conc e n t r a t i o n s as in Fig. 2. T h e b r o k e n line refers to the " c r e s t lines". A, B, C, I and II refer to the zones observed in Fig. 7.
28
IO+m'cm2 Tr
Ir
.,,
l,
5
t 5
-+'
5
(105
5
06
_ .
log C/c
0
: d5 o
-ElV ~
0
d5 b
-E/V v
-I
0
0
io9 C/cs
,
-1
0
c
Fig. 5. Plot of the superficial excess as a function of: (a) the potential; (b) the charge density (relative concentrations as indicated); ic) the logarithm of the relative concentration for various values of the potential; (d) the logarithm of the relative concentration for various values of the charge density.
resorting to the various procedures mentioned in Part I [ 1] (verification of the Lippmann equation, integration of the capacities extrapolated at zero frequency, cross-derivation of the Gibbs equation). As gradual interference by slow mass transfer develops for the lowest concentrations (c ~< 5 × 10 -2 Cs), the corresponding data will be only considered as representative of tendencies. DISCUSSION With respect to Q, the major peculiarity shown by isoQ is the existence of a discontinuous transition which involves, on its negative potential side, a new, remarkable surface state (delimited by region I, in Figs. 2--4) which will be examined first. The invariability of the differential capacity within an extended range of concentration and potential constitutes a clue to the presence of an organic monolayer. This conclusion is supported by the constancy of the superficial excess. The value of the limiting area per molecule (0.21 nm 2) can only be explained if it is assumed that all molecules are aligned perpendicularly to the electrode surface, in such a way that one aromatic ring faces the electrode and the other the solution. Evidence that all the molecules are oriented with their nitrogen atom pointing towards the solution (Table 2, position 6--7) is afforded by the value of the capacity which is at the same time constant and low (4.50 ttF cm -2 ). These properties can be interpreted in terms of an organic monolayer which interacts only by electronic polarizability with the field generated by the electrode. Coexistence of " u p " and " d o w n " orientations, to an extent depending on the potential, would result in variable and much larger values of the capacity. The proposed orientation is consonant with a thickness of about 0.6--0.7 nm and a dielectric constant of the order of 3--3.5, a set of values which are in fair agreement with the molecular dimensions and the refractive index of the surfactant. The zero potential, evaluated by extrapolating the
29 charge--potential curve (Fig. 2, line a), and referred to the pzc observed in the absence of surfactant, is of the order of +1.70 + 0.05 V, a value close to +1.67 V which is the theoretical Rideal dipole potential, and which should be developed by a compact layer of fully oriented molecules with P = 8.1 X 10 -~° mol cm -2, a normal dipole component of 2.4 D and a dielectric constant of 2.62. A third indication, arguing for an invariant local structure, is given by the good linearity of the Tafel slopes in region I of various faradaic processes which are inhibited by isoQ [3]. An array of parallel dipoles at the electrode surface constitutes a configuration which is intrinsically unfavourable. However (and as for Q), the dipole-dipole repulsion is obviously more than compensated by the combined action of the electrode field governed by the electrode, the interplay of short-range lateral dispersion forces and the "chemical" polarity of the molecules of the surfactant, which spontaneously tend to direct their hydrophilic nitrogen group towards the bulk of the solution. The residual dipole c o m p o n e n t parallel to the surface (0.9 D) is probably disposed in an antiparallel way, which adds to the stability of the structure. It should be noted that the lateral stabilization by all these factors extends to very negative potentials (Fig. 3), at which the sudden occurrence of desorption corresponds to an isotherm of the Frumkin class, with a large, positive, attractive parameter. Although it is much more confined, the region adjacent to the transition line on its positive potential side (Figs. 2~4, region II) is also characterized by a set of const~ant values for the charge density at a given potential (~6.6 pC cm -2 at ~ 0 . 7 0 0 V), the capacity (C = 5.8 pF cm -2) as well as the superficial excess (F H = 6.4 + 0.3 X 10 -~° mol cm -~) ~ exception being made for the most extreme concentrations which will be briefly examined at a later stage. Extrapolation of the linear segment common to all the pertaining charge~potential curves (Fig. 2, line b) gives a pzc potential of +0.47 + 0.1 V. The projected area per molecule is of the order of 0.26 nm 2, a value which clearly excludes the flat orientation, but matches satisfactorily positions such as " 4 ~ 5 " or " 5 ~ 6 " {see Table 2). The related normal c o m p o n e n t of the molecular dipole moment, as can be estimated from the Rideal equation, amounts to 1.7 + 0.4 D. This value can hardly be reconciled at all with the 4 ~ 5 position for which the normal vector would not exceed 0.9 D. The slanting position 5--6 (for which the perpendicular m o m e n t is 1.99 D) appears to be the best-fitting one, inasmuch as validity of the Rideal equation is taken for granted, and as far as it is assumed that (1) all molecules must necessarily be parallel to each other, and (2) tilting of the molecular plane is excluded. The observed dipole potential could also possibly result from the existence of a mixed homogeneous structure, which ought then to retain a definite stoichiometry throughout the whole of region II. Stoichiometric arrays of 4 ~ 5 and 6 ~ 7 molecules can be ruled out as they cannot simultaneously fit the experimental values found for the superficial excess and the dipole potential. Besides, such a surface state conflicts with evidence of a kinetic nature, deduced from study of the transition kinetics [ 7]. The electrocapillary parameters of region II can be evaluated much more accurately, by turning to account the reasonable precision in the determination of the transition potential E w. If each of the two adjacent regions I and iI are described in terms of their own set of P = f(E, c) parabolas, transition neces-
30
sarily occurs at the intersection of the two curves, the prevailing configuration being fixed, at any potential, by the condition of minimal superficial excess free energy (Fig. 6). Increasing the bulk concentration from a to b (arrows) brings about vertical shifts, the amplitude of which depends on the values of the two superficial excesses F ~ and F n, according to the Gibbs adsorption equation: d7 = --aM d E -
R TF d In c
The resulting displacement of the transition potential (from ET, a to ET, b), which occurs at the expense of the less dense monolayer, can easily be calculated from the equations of the two individual parabolas. A further way of cross-checking the data is afforded by the experimental values of the charge step Ao T at the transition potential (Table 4), which must be equal to the difference between the slopes of the two curves at their intersection. In the present case, the parameters defining region I are known with good accuracy (Fim~x = 8.1 -+ 0.3 X 10 -~° mol cm -2, E~z~ = +1 .21 -+ 0.05 V, C I = 4.50 #F cm -~ and 3'Iax = 435.4 mN m -I for c = 0.5G). These data have been used, together with E T and A(I T (Table 4), in order to optimize the parameters proper to region II. The best fit has been obtained with the following set of values" Fm -2 , ~ mii~ x - - _ ~I~ x = 6.7 + 0.1× 10 -~° mol cm -2 , E ~ c = +0.47V, c II = 5 . 8 p F c m 390.3 mN m -~ (for c = 0.5c~). The slight deviations which have been observed at both extremes of the concentration range can be explained, respectively, by the fact that F II increases slightly over its limiting value for the highest concentration (Fig. 4), and that mass transfer equilibrium is not rigorously achieved for the most dilute solutions. These second-order factors improve the fortuitous apparent linearity of the ET = f(log c) plot (Fig. 2, inset), which results, however, chiefly from the remoteness from the vertices of the two electrocapillary curves. The new values found for EI~c and the related dipole potential reconfirm the previous values, with narrower limits, so that all conclusions drawn previously remain basically unaffected (except perhaps that one of the simple orientations perfectly matches the data within the new error margins; however,
1
, .
.
.
.
I
t
ET, b ET,a
E
Fig. 6. Schematic shift of the transition potential with the surfactant concentration.
31 TABLE 4 Transition potential (ET) and amplitude of the charge step (~o T) for various isoQ concentrations c/c s
ET/V
AtrT/#C cm -2
1.0 0.7 0.5 0.3
-0.730 -0.790 -0.860 -0.970
1.7 1.7 1.6 1.5
this can of course be expected, considering all the shortcomings which are implicit in the Helmholtz--Rideal equation). The procedure which has been described above is based on electrocapillary parabolas expressed in terms of the total electrode potential difference. This method ignores, therefore, the presence of the diffuse part of the double layer, whose potential contribution should be deducted, since the capacity of the inner layer is controlled only by E -- ¢2. In ½M Na~SO4, ¢2 potentials are small and proportional to the charge density, so that no significant error results from the omission of the correction - - e x c e p t that the overall electrode capacities axe smaller than the inner values by <2%, which roughly compares with the experimental precision. This is, however, no longer the case when the concentration of the supporting electrolyte is lowered. The position of the transition potential then depends in a complex way on the simultaneous variation of the ¢2 potential and of the activity coefficient of the surfactant (as for all other electrocapillary properties [ 10,11 ] ). The activity factor can be kept constant if the concentrations are maintained proportional to each saturation value. Under such standardized conditions, the shift of E w with the concentration of the supporting electrolyte (Table 5) practically coincides with the corresponding variation of ¢~ calculated from the Gouy--Chapman theory applied to the 1--2 electrolyte. The fact that E T is exclusively controlled by E -- ¢2 calls for the following observations: (1) The "ideal" transition potentials which would be observed in the total absence of a diffuse double layer can be easily and accurately evaluated by a short-range linear extrapolation. (2) The fact that ¢2 is controlled only by aM demonstrates the absence of any detectable c o m p o n e n t of charge inside the organic monolayer, and supports the initial assumption that there is no significant coadsorption when OH-, For SO4~- are the anions of the supporting electrolyte. The charge balance is of course different when coadsorption is allowed [ 12]. (3) More fundamentally, the determination of ¢2 potentials from ET shifts may constitute a new method of investigating diffuse layers which is simpler and more direct than most other classical methods. As has been already observed, the transition potential is markedly temperature dependent. The coefficient--17 mV/°C given by Greef [4] refers to solutions kept at a constant concentration of isoQ. Jenard and Hurwitz [6], using
32 TABLE 5 Transition potential corrected for the corresponding variation of ~2 when the concentration of the electrolyte is changed Electrolyte concentration [Na2SO4 ]/M
ET/V
~2/V
(ET--~2)/V
0.5 0.05 0.005
-0.730 -0.785 --0.840
-0.028 -0.079 -0.132
-0.702 -0.708 -0.708
saturated solutions of isoQ in 1 M KCI and in 0.1 M Na2CO3, have observed a slope o f - - 1 4 mV/°C, which is very similar to our own value, obtained in ½M Na2SO4 between 5 and 35°C. All other conditions being maintained constant, the variation of the interfacial tension with temperature is given by d3' = --s dT. As for the variation of the transition potential with temperature, this may be expressed in terms of the relative vertical shift of the electrocapillary curves which characterize the two distinct molecular films. The translation of E w is thus an indirect function of the difference of superficial entropy between the two monolayers. It has been found that for c = 0.5c~ the superficial entropy difference amounts to +2.3 × 10 -a J cm -2 K -~, when the transition is crossed from positive to negative potentials. This trend meets the intuitive idea that organization increases with compactness. The existence of a phase transition can be explained in the simplest terms by stating that there is total "incompatibility" between two different superficial states. In other words, this means that each of the two stable states possesses a free energy lower than that of any possible mixed structure. In the present case the transition involves two adsorption layers, each of them obeying its own Frumkin-type isotherm, associated with a large, positive interaction coefficient. Such large coefficients reflect the compact assemblage of rigid molecules which, being given the electrical field imposed by the electrode, can organize themselves through the interplay of various lateral orientation-promoting forces. The resulting layers constitute defined, condensed, superficial phases, either liquid, mesomorphic or solid (more precise evidence for the existence of given states will be available in the papers dealing with the kinetics of the transition and the inhibition power). On the other hand, interaction between molecules representative of regions I and II must be considerably weaker in view of the steric and electrical mismatch at the molecular level, and the impaired ability to mobilize forces strong enough to overcome the greater stability of the two prevailing condensed structures. As a result, mixed films are highly unstable, and can only be formed (and their decay observed), by resorting to sophisticated instrumentation. They correspond to the forbidden part of an S-shaped isotherm which, instead of describing the competitive behaviour of a surfactant and a solvent as is usually the case, here refers to the antagonism between two different monolayers, each of them made of the same molecules.
33 In the range of less negative potentials and lower concentrations, compact films are no longer present, and the behaviour becomes much more complex, as demonstrated by Figs. 2--5. Rather unexpectedly, the trends shown in the C, OM and F plots are strikingly similar to those found with Q. For both surfactants, at concentrations of medium and low values, there is a gradual desorption towards positive potentials, while, for concentrations close to saturation, the superficial excesses increase over the limiting value characterizing the adjacent compact film. These two opposite types of behaviour explain why, with increasing concentrations, the capacity peak shifts first towards less negative potentials, to reverse its deplacement as soon as the concentration is sufficiently large {about 0.3Cs). As was the case for Q, the main characteristics of the adsorption can be made most salient by constructing the Esin--Markov diagram (Fig. 7), which plots the potential shift as a function of the superficial excess, for discrete charge density values. Regions I and II are confined to the two vertical lines I and iI, which are separated by a space which corresponds to the forbidden part of the isotherm. As analysis of the left-hand side of the diagram can be based on the same arguments which have been fully developed for Q, they do not need to be repeated at great length. The major conclusions are as follows: (1) The fan-like array of curves originating from the O--O coordinates covers a region (zone A) which consists of an aqueous dilute film of molecules lying
+2
,4 *6
/,
+ 0.2-
z,//i o-,
o.o
0
~
1010I-'/MOi CM-2
Fig. 7. Plot of the shift of potential as a function o f the superficial excess for various charge densities, as indicated.
34 fiat. There is no appreciable potential shift with F when the charge density is close to zero, as can be expected for a surfactant having no dipolar component perpendicular to the surface. (2) Zone B can be explained in terms of a mixed structure composed of water and variable amounts of fiat and standing molecules (to the exclusion of intermediate-tilted orientations, as indicated by the breaks in the curves). (3) The upper left-hand side of the plot is consistent with the gradual development of a new state, characterized by abnormally high superficial excesses, lower capaci£y values and a conspicuous reversal of the potential shift. The latter effect suggests that the molecules embedded in the multilayer fraction of the film organize themselves in a way which tends to cancel (or at least to lower) their dipole contribution to A ( E - ¢2). The existence of this zone has recently been established independently [4]. That there is a nearly complete loss of identity between Q and isoQ (and also, as will be shown in Part IV, for most of their mono-methyl derivatives) may at first seem puzzling. It can, however, be understood by considering that the smaller the normal vector of the dipole moment, the less distinct is the differentiation, so that aromatic surfactants tend to display a common behaviour, as far as they are lying flat and dilute on a surface, and behave as a dielectric with a common thickness. The gradual change from this state to a full compact layer requires about the same charge density for both molecules, and since their dipole moments are commensurate, the potential span required for the reorientation process (and therefore the height and width of the capacity peaks) is also nearly identical. Closer examination of the diagrams representative of the two surfactants reveals, however, some slight differences. Superadsorption is somewhat larger for Q, for which zone C is more extended in the Esin--Markov plot and corresponds to larger A~, and smaller differential capacities. At intermediate and lower coverages, Q still remains more adsorbed than isoQ, as indicated by comparison between their respective isotherms and by additional evidence of a kinetic nature, which will be presented at a later stage
[71. As for quinoline [ 1], it has been preferred to avoid undue repetitions by deferring the general conclusions to Parts IV and V, as the complementary study of several methyl-derivatives of Q and isoQ {coupled with the elucidation of the inhibition processes which are observed in their presence) will then be at hand, with new elements of discussion. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12
CI. Buess-Herman, L. Giexst and N. Vanlaethem-Meur~e, J. Electroanal. Chem., 123 (1981) 1. H. Wefler-Feflchenfeld, Yuh-Lin Mao and E.D. Bergmann, Isr. J. Chem., 9 (1971) 111. J.P. L a m b e r t , D o c t o r a t e Thesis, Univer~it~ Libre de BruxeHes, 1973. M.W. H u m p h r e y s and R. Parsons, J. Electroanal. Chem., 82 (1977) 369. N. V a n l a e t h e m , J.P. Lambert and L. Gierst, Chem. Ing. Tech., 44 (1972) 219. A. Jenard and H.D. Hurwitz, J. Electroanal. Chem., 70 (1976) 27. G. Quarin, CI. Buess-Herman and L. Gierst, J. Electroanal. Chem., 123 (1981) 35. M.M. Potashnikov and I.G. Belavina, Zh. Prikl. K h i m . , 38 (1965) 1585. CI. Buess-Herman, G. Quarin and L. Gierst, J. Electroanal. Chem., Part IV, to be published. A.A. Survila and B.B. Damaskin, E l e k t r o k h i m i y a , 3 (1967) 1138. B.B. Damaskin, E.V. Stenina, V.A. Yusupova and N.V. Fedorovich, E l e k t r o k h i m i y a , 8 (1972) 1409. C1. Buess-Herman, G. Quarin and L. Gierst, J. Electroanal. Chem., Part VI, to be published.
21
ON THE BEHAVIOUR OF MOLECULES OF THE QUINOLINE GROUP AT THE WATER--MERCURY INTERFACE PART iI. ELECTROCAPILLARY STUDY OF ISOQUINOLINE IN SOLUTIONS CONTAINING ONLY AN "INERT" ELECTROLYTE
C1. BUESS-HERMAN, N. VANLAETHEM-MEUREE, G. QUARIN and L. G I E R S T
Facult4 des Sciences, Universitd Libre de Bruxelles, 50, avenue F.D. Roosevelt, 1050 Brussels (Belgium) (Received 15th December 1980)
ABSTRACT
Isoquinoline has been examined in a 0.5 M Na2SO4 aqueous solution, by using various electrocapillary methods. In the range of small potentials, and rather surprisingly, isoQ and Q present a very similar behaviour, despite the considerable difference in their dipole orientation with respect to the median carbons. Esin--Markov plots and capacity--potential-concentration maps indicate that there is gradual desorption towards negative potentials. Three distinct states (dilute flat molecules, a mixture of flat and erect molecules, and a m o n o l a y e r with a variable extent of local clusters) have been characterized and their domain of prevalence demarcated. At more negative potentials, a sudden reorientation is observed, which reflects the lack of miscibility between two superficial states internally stabilized by their own set of lateral interactions. The effects of the surfactant concentration, the electrolyte concentration and the t e m p e r a t u r e on the position of the transition potential are interpreted on the basis of the relative position of two 7 = f(E) curves, whose vertical shifts are controlled by the Gibbs isotherm equation, with two distinct discrete values for the superficial excesses corresponding to the two antagonizing monolayers. The most c o m p a c t of these monolayers is stable in the region with stretches between the phase transition potential and the electrolyte potential, both very sharply defined, its structure has been determined on the basis of a n u m b e r of converging facts, such as the limiting value of F, the values of the capacity and the position of the pzc, extrapolated from the o M = f(E) plots.
INTRODUCTION
In the first paper of this series [1], the adsorption of quinoline (Q) at the water--mercury interface has been investigated, by considering the case of a solution containing an inactive supporting electrolyte. Despite this major simplification, the behaviour remains rather involved. A number of distinct superficial states have been characterized as a function of the bulk concentration and the electrode potential, and connected to various orientations of the molecule. In order to check the validity of the conclusions which have been drawn, it is of interest to examine other molecules as equivalent as possible in terms of size, atomic arrangement and electrical properties, but differing by the orientation of their dipole moment. 0 0 2 2 - 0 7 2 8 / 8 1 / 0 0 0 0 - - 0 0 0 0 / $ 0 2 . 5 0 © Elsevier Sequoia S.A.
22
Isoquinoline (isoQ) meets all these requirements. Most of its properties (Table 1) are very similar to those of quinoline. However, the presence of the nitrogen atom in position 2 is now such that the dipole m o m e n t (2.6 D in vacuo) presents an angle of the order of 70 ° [2] with respect to the bond between the two median carbons (C4a--Csa), so that the vector components parallel and normal to this axis are respectively of 0.9 and 2.4 D (Table 2) (for Q, for which the angle is 2 °, the corresponding values were 2.2 and 0.1 D). This change in the direction of the dipole can be expected to promote new stable molecular orientations at the interface. Therefore, comparison between Q and isoQ affords a possible way of disentangling the effects simply related to the overall geometry of the two molecules from those which are more specifically controlled by the interactions between the dipole and the electrical field of the electrode. The most peculiar phenomenon presented by isoQ is the existence of a very sharp phase transition between two condensed layers, which must necessarily be of incompatible structures. The pioneering work of Lambert [ 3] has shown that the extent of inhibition exerted by isoQ on a number of electrode processes is considerably larger for the structure which prevails at potentials negative with respect to this transition. [As a typical example, the reduction of Tl * is characterized by an unaffected reversible wave leading to a fully developed diffusion current, as long as the transition potential ( E T ) iS not reached. As soon as it passed beyond this, the current drops to zero and maintains this value for more than 500 mV.] Lambert made the suggestion that with increasingly negative potentials, the isoQ molecules, which first stay flat, then tilt gradually towards a perpendicular position (with their carbons 4 and 5 adjacent to the surface as for Q) to assume finally, in an abrupt way, a third upright orientation with carbons 6--7 in contact with the metal. This model has been supported in recent work based on ellipsometry [ 4], which has also provided indications pointing to multilayer formation, for solutions close to saturation. The existence of a restructuration transition process raises interesting problems of a kinetic nature. Lambert and co-workers [5 ], using square-wave signals
TABLE 1 Some properties of
isoquinoline
S o l u b i l i t y in w a t e r 2 5 ° C
Boiling point Melting point pK a Refractive index
Dipole moment Dimensions
4.2 × 10 -2 M 243.2°C 26.5°C 5.4 1.62 2.6 D a = 0.7 n m b = 0.87 n m c = 0.3 n m
23 TABLE 2 Basic orientations of isoquinoline and derived quantities Carbons on the surface
Orientation
Area /nm 2
Thickhess /nm
101° Fmax /mol cm -2
Dipole m o m e n t / D Perpendicular to the surface
Parallel to the surface
All
Flat
0.61
0.3
2.7
--
--
4--5
Perpendicular "4--5"
0.26
0.7
6.4
0.88
2.44
5--6
Perpendicular "5--6"
0.267
0.73
6.2
1.99
1.67
Perpendicular "6--7"
0.21
0.87
7.9
2.44
0.88
6--7
of adequate amplitude, reached the conclusion that the process of rearrangement is basically of the "nucleation and g r o w t h " type, as indicated by sharp peaked i = f(t) transients presenting a marked hysteresis. These exploratory experiments were followed by the work of Jenard and Hurwitz [6], who investigated the major characteristics of the rising part of the cathodic chronopotentiometric transients. In these works, which were principally aimed at bringing out phenomenological facts, no a t t e m p t was made to interpret quantitatively the rate of the process in terms of a suitable model and an adequate formulation. The whole problem will be reconsidered in more detail in Part III of this series [ 7]. The present paper is strictly limited to the t h e r m o d y n a m i c aspects of the adsorption of isoQ in the presence of Na2SO4 which, as stated in Part I, can be considered (together with NaOH and NaF) as a suitable "inert" electrolyte in the sense that anionic coadsorption in the inner layer is essentially negligible. Potashnikov [8] has shown that isoQ, in the presence of Na2SO4, is affected by the salting-out effect. Our own determinations (Table 3) confirm quantitatively his results and indicate that NaF, NaOH and Na2SO4 act in a very similar way, in contrast with NaC104 which is notably less active, as is generally the case for organic solutes. Salting-out is easily accounted for by expressing the surfactant concentration with respect to its solubility in the particular electrolyte employed.
TABLE 3 Solubility of isoquinoline in various electrolytes [Na+]/M
Na2SO4
NaOH
NaF
NaC104
0.1 0.5 1.0
3.9 x 10-2M 3.1 x 10-2M 2.1 x 10-2M
3.8 X 10-2M 2.9 x 10-2M 2.0 x 10-2M
3.8 x 1 0 - 2 M 2.7 x 10-2M 2.0 x 10-~M
3.9 x 1 0 - 2 M 3.7 x I 0 - 2 M 3.3 x 10-2M
24 EXPERIMENTAL
Reagents Water, mercury and Na2SO4 were treated as described previously [ 1]. Here, isoQ (UCB " p u r u m " grade) was submitted to two successive fractional distillations under vacuum. The final colourless liquid was kept under nitrogen in a refrigerator, in order to slow down any possible degradation. The saturated concentration is 2.1 × 10 -2 M in ½ M Na2SO4 aqueous solutions kept at 25°C.
Instrumentation The various techniques which have been used to measure interfacial tensions, charge densities and differential capacities are identical to those already reported [ 1 ]. However, drop times at very negative potentials are poorly reproducible, so that superficial tensions have been computed in that range by integration of the charge density--potential plots. Electrode potentials are always referred to the saturated calomel electrode (SCE). RESULTS The adsorption of isoQ has been investigated under the same conditions as for Q, in a potential range ( 0 . 0 - 1 . 0 V) which is again limited here by the gradual onset of interfering faradaic processes involving the surfactant itself. Adsorption is evident from the progressive lowering of the interfacial tension 7 with increasing concentrations. Figure 1, which features a family of electrocapillary curves A'y = f(E), shows the existence of distortions at negative potentials for the most concentrated solutions. Complete desorption is not observable, since it occurs, in both directions, at potentials where the behaviour is no longer that of an ideally polarizable electrode. However, the desorption at negative charges can be detected as a distinct break, at medium concentrations of isoQ and in strongly alkaline solutions, for which the interference of reduction processes is less troublesome. The charge density--potential plots (Fig. 2) show that the potential of zero charge (pzc) shifts markedly towards less negative values with increasing concentration without reaching any defined limiting value. This trend is reversed for the most concentrated solutions at small negative charge densities {about f r o m - - 2 to --6 pC cm -2 ). The existence of an ill-defined nodal point suggests the presence of an extremum in the F = f(E) function. At potentials around --0.65 V, the charge densities tend to values which are independent of the bulk concentration. Qualitatively at least, all these experimental facts are not drastically different from those observed with Q. However, in striking contrast with the latter, a vertical step of more than 1.5 pC cm -2 disjoints the charge~potential plot. The discontinuity shifts towards more negative potentials with dilution, according to a function which is approximately logarithmic in concentration (Fig. 2, inset). Beyond the transition, the charge densities are essentially concentration independent until desorption is finally reached. The data on double-layer capacities, obtained by tensammetry, are presented
25
A ~ ' / m N . m -1
1
60
/
7.10 -~
/ , /
5.10
/ /
40
~
-~
3.10-~ 2.10-~
10 -1 7.10 -2 20 5 . 1 0 -2
3.10 -2
"
I
o.o
05
2.10 -2 ' 10-~ --~.~
1:o -E/v (SCE)
Fig. 1. Influence of the potential on the lowering of interfacial tensions in ½M Na2SO4 for the relative concentrations c/cs, as indicated.
under the form of isocapacity c o n t o u r curves, plotted as a function of the potential and the logarithm of the concentration (Fig. 3). At potentials close to 0.0 V, and within an extended concentration range, fairly large values of the capacity are observed. Two crest lines (thick solid lines), which are the locus of the capacitive peaks, have approximately the same origin at E = --0.25 V and log c/ cs = --0.55. When the isoQ concentration is gradually increased, the peak shifts first to less and then to more negative potentials, a fact which suggests a qualitative change in the properties of the surface state which prevails in the left-hand side of the diagram. At more negative potentials (zone II) the capacities decrease and reach a constant value of the order of 5.8 p F cm -2 , except for the m o s t c o n c e n t r a t e d solution. At even more negative potentials, a sudden discontinuity {broken line a in Fig. 3) demarcates an extended, new, distinct region (zone I) which is similarly characterized by a constant capacity of lower value (4.5 p F cm -2 ). Capacity and charge steps are coincident. The absence of any capacity peak, even at the lowest frequencies, indicates that the related kinetic process can be considered as infinitely slow, at least for a potential amplitude of +5 mV. Desorption occurs at negative potentials which are shifted negatively with increasing concentrations (curve/~). A triple point ( d o t t e d circle) is observed for a bulk concentration of the order of 5 × 10 -2 c, and at a potential of ~ 1 . 2 6 V. Values of the superficial excesses have been calculated by graphical differentiation of the F = f ( R T in c) plots, with an absolute standard deviation of the
26
/
/~JC.cm 2
/ -10 i
i
i i
t i
. i
b
-5-
/
p
t J
1 1
log c/c s
-1
/ /
/ I I
I oo
tb - E/v0.5
1D
-E/V(SCE)
-
Fig. 2. Plots of charge density vs. potential for the following relative isoQ concentrations: (1) 0.01; (2) 0.02, (3) 0.03; (4) 0.05; (5) 0.07; (6) 0.1; (7) 0.2; (8) 0.3; (9) 0.5, (10) 0.7; (11) 1. The broken line refers to the electrolyte alone. Dotted and dashed lines (a) and (b) show the extrapolation of the two ranges of common charge densities. A, B, C, I and II refer to the zones observed in Fig. 7. Inset: variation of the transition potential with the logarithm of the relative concentration.
order of 0.3 × 10 -'0 mol cm -2. For the sake of subsequent comparison [9], they have been plotted as a function of both the electrode potential (Fig. 4) and the charge density (Fig. 5a). The transition already observed on the charge and capacity curves appears here as a conspicuous step in the superficial excess, which increases discontinuously from about 6.4 + 0.3 X 10 -'o to 8.1 + 0.3 X 10 -~° mol cm -2, and depends very little on the concentration except for the saturated solution, for which it is slightly smaller. The experimental isotherms have been constructed by plotting the surface excess as a function of the logarithm of the bulk concentration, for several selected values of the potential (Fig. 5b) and the charge density (Fig. 5c). The accuracy and precision of all experimental data have been checked, by
27
iogC/c s
C
Y,
I 4.5
\
/
//3
\ /
\
-1
isoQ
/
/
( \?'~1
/
H20 18- 20
12
-2
40
30 25 20
15
O~
~:o
=
-E/V(SCE)
Fig. 3. I s o c a p a c i t y c o n t o u r map as a f u n c t i o n of the isoQ c o n c e n t r a t i o n . The two thick lines are the "crest lines". B r o k e n lines 0~ and/3 c o r r e s p o n d respectively to the t r a n s i t i o n and d e s o r p t i o n . T h e d o t t e d circle locates the triple point. A, B, C, I and Ii refer to the zones observed in Fig. 7.
101° ["/tool cm -2 10
I-
75
:I J
I.
9
5
2
2.5
0.0
O.5
1£)
-E/V(SCE)
Fig. 4. Superficial excess as a f u n c t i o n of the electrode p o t e n t i a l for the same relative conc e n t r a t i o n s as in Fig. 2. T h e b r o k e n line refers to the " c r e s t lines". A, B, C, I and II refer to the zones observed in Fig. 7.
28
IO+m'cm2 Tr
Ir
.,,
l,
5
t 5
-+'
5
(105
5
06
_ .
log C/c
0
: d5 o
-ElV ~
0
d5 b
-E/V v
-I
0
0
io9 C/cs
,
-1
0
c
Fig. 5. Plot of the superficial excess as a function of: (a) the potential; (b) the charge density (relative concentrations as indicated); ic) the logarithm of the relative concentration for various values of the potential; (d) the logarithm of the relative concentration for various values of the charge density.
resorting to the various procedures mentioned in Part I [ 1] (verification of the Lippmann equation, integration of the capacities extrapolated at zero frequency, cross-derivation of the Gibbs equation). As gradual interference by slow mass transfer develops for the lowest concentrations (c ~< 5 × 10 -2 Cs), the corresponding data will be only considered as representative of tendencies. DISCUSSION With respect to Q, the major peculiarity shown by isoQ is the existence of a discontinuous transition which involves, on its negative potential side, a new, remarkable surface state (delimited by region I, in Figs. 2--4) which will be examined first. The invariability of the differential capacity within an extended range of concentration and potential constitutes a clue to the presence of an organic monolayer. This conclusion is supported by the constancy of the superficial excess. The value of the limiting area per molecule (0.21 nm 2) can only be explained if it is assumed that all molecules are aligned perpendicularly to the electrode surface, in such a way that one aromatic ring faces the electrode and the other the solution. Evidence that all the molecules are oriented with their nitrogen atom pointing towards the solution (Table 2, position 6--7) is afforded by the value of the capacity which is at the same time constant and low (4.50 ttF cm -2 ). These properties can be interpreted in terms of an organic monolayer which interacts only by electronic polarizability with the field generated by the electrode. Coexistence of " u p " and " d o w n " orientations, to an extent depending on the potential, would result in variable and much larger values of the capacity. The proposed orientation is consonant with a thickness of about 0.6--0.7 nm and a dielectric constant of the order of 3--3.5, a set of values which are in fair agreement with the molecular dimensions and the refractive index of the surfactant. The zero potential, evaluated by extrapolating the
29 charge--potential curve (Fig. 2, line a), and referred to the pzc observed in the absence of surfactant, is of the order of +1.70 + 0.05 V, a value close to +1.67 V which is the theoretical Rideal dipole potential, and which should be developed by a compact layer of fully oriented molecules with P = 8.1 X 10 -~° mol cm -2, a normal dipole component of 2.4 D and a dielectric constant of 2.62. A third indication, arguing for an invariant local structure, is given by the good linearity of the Tafel slopes in region I of various faradaic processes which are inhibited by isoQ [3]. An array of parallel dipoles at the electrode surface constitutes a configuration which is intrinsically unfavourable. However (and as for Q), the dipole-dipole repulsion is obviously more than compensated by the combined action of the electrode field governed by the electrode, the interplay of short-range lateral dispersion forces and the "chemical" polarity of the molecules of the surfactant, which spontaneously tend to direct their hydrophilic nitrogen group towards the bulk of the solution. The residual dipole c o m p o n e n t parallel to the surface (0.9 D) is probably disposed in an antiparallel way, which adds to the stability of the structure. It should be noted that the lateral stabilization by all these factors extends to very negative potentials (Fig. 3), at which the sudden occurrence of desorption corresponds to an isotherm of the Frumkin class, with a large, positive, attractive parameter. Although it is much more confined, the region adjacent to the transition line on its positive potential side (Figs. 2~4, region II) is also characterized by a set of const~ant values for the charge density at a given potential (~6.6 pC cm -2 at ~ 0 . 7 0 0 V), the capacity (C = 5.8 pF cm -2) as well as the superficial excess (F H = 6.4 + 0.3 X 10 -~° mol cm -~) ~ exception being made for the most extreme concentrations which will be briefly examined at a later stage. Extrapolation of the linear segment common to all the pertaining charge~potential curves (Fig. 2, line b) gives a pzc potential of +0.47 + 0.1 V. The projected area per molecule is of the order of 0.26 nm 2, a value which clearly excludes the flat orientation, but matches satisfactorily positions such as " 4 ~ 5 " or " 5 ~ 6 " {see Table 2). The related normal c o m p o n e n t of the molecular dipole moment, as can be estimated from the Rideal equation, amounts to 1.7 + 0.4 D. This value can hardly be reconciled at all with the 4 ~ 5 position for which the normal vector would not exceed 0.9 D. The slanting position 5--6 (for which the perpendicular m o m e n t is 1.99 D) appears to be the best-fitting one, inasmuch as validity of the Rideal equation is taken for granted, and as far as it is assumed that (1) all molecules must necessarily be parallel to each other, and (2) tilting of the molecular plane is excluded. The observed dipole potential could also possibly result from the existence of a mixed homogeneous structure, which ought then to retain a definite stoichiometry throughout the whole of region II. Stoichiometric arrays of 4 ~ 5 and 6 ~ 7 molecules can be ruled out as they cannot simultaneously fit the experimental values found for the superficial excess and the dipole potential. Besides, such a surface state conflicts with evidence of a kinetic nature, deduced from study of the transition kinetics [ 7]. The electrocapillary parameters of region II can be evaluated much more accurately, by turning to account the reasonable precision in the determination of the transition potential E w. If each of the two adjacent regions I and iI are described in terms of their own set of P = f(E, c) parabolas, transition neces-
30
sarily occurs at the intersection of the two curves, the prevailing configuration being fixed, at any potential, by the condition of minimal superficial excess free energy (Fig. 6). Increasing the bulk concentration from a to b (arrows) brings about vertical shifts, the amplitude of which depends on the values of the two superficial excesses F ~ and F n, according to the Gibbs adsorption equation: d7 = --aM d E -
R TF d In c
The resulting displacement of the transition potential (from ET, a to ET, b), which occurs at the expense of the less dense monolayer, can easily be calculated from the equations of the two individual parabolas. A further way of cross-checking the data is afforded by the experimental values of the charge step Ao T at the transition potential (Table 4), which must be equal to the difference between the slopes of the two curves at their intersection. In the present case, the parameters defining region I are known with good accuracy (Fim~x = 8.1 -+ 0.3 X 10 -~° mol cm -2, E~z~ = +1 .21 -+ 0.05 V, C I = 4.50 #F cm -~ and 3'Iax = 435.4 mN m -I for c = 0.5G). These data have been used, together with E T and A(I T (Table 4), in order to optimize the parameters proper to region II. The best fit has been obtained with the following set of values" Fm -2 , ~ mii~ x - - _ ~I~ x = 6.7 + 0.1× 10 -~° mol cm -2 , E ~ c = +0.47V, c II = 5 . 8 p F c m 390.3 mN m -~ (for c = 0.5c~). The slight deviations which have been observed at both extremes of the concentration range can be explained, respectively, by the fact that F II increases slightly over its limiting value for the highest concentration (Fig. 4), and that mass transfer equilibrium is not rigorously achieved for the most dilute solutions. These second-order factors improve the fortuitous apparent linearity of the ET = f(log c) plot (Fig. 2, inset), which results, however, chiefly from the remoteness from the vertices of the two electrocapillary curves. The new values found for EI~c and the related dipole potential reconfirm the previous values, with narrower limits, so that all conclusions drawn previously remain basically unaffected (except perhaps that one of the simple orientations perfectly matches the data within the new error margins; however,
1
, .
.
.
.
I
t
ET, b ET,a
E
Fig. 6. Schematic shift of the transition potential with the surfactant concentration.
31 TABLE 4 Transition potential (ET) and amplitude of the charge step (~o T) for various isoQ concentrations c/c s
ET/V
AtrT/#C cm -2
1.0 0.7 0.5 0.3
-0.730 -0.790 -0.860 -0.970
1.7 1.7 1.6 1.5
this can of course be expected, considering all the shortcomings which are implicit in the Helmholtz--Rideal equation). The procedure which has been described above is based on electrocapillary parabolas expressed in terms of the total electrode potential difference. This method ignores, therefore, the presence of the diffuse part of the double layer, whose potential contribution should be deducted, since the capacity of the inner layer is controlled only by E -- ¢2. In ½M Na~SO4, ¢2 potentials are small and proportional to the charge density, so that no significant error results from the omission of the correction - - e x c e p t that the overall electrode capacities axe smaller than the inner values by <2%, which roughly compares with the experimental precision. This is, however, no longer the case when the concentration of the supporting electrolyte is lowered. The position of the transition potential then depends in a complex way on the simultaneous variation of the ¢2 potential and of the activity coefficient of the surfactant (as for all other electrocapillary properties [ 10,11 ] ). The activity factor can be kept constant if the concentrations are maintained proportional to each saturation value. Under such standardized conditions, the shift of E w with the concentration of the supporting electrolyte (Table 5) practically coincides with the corresponding variation of ¢~ calculated from the Gouy--Chapman theory applied to the 1--2 electrolyte. The fact that E T is exclusively controlled by E -- ¢2 calls for the following observations: (1) The "ideal" transition potentials which would be observed in the total absence of a diffuse double layer can be easily and accurately evaluated by a short-range linear extrapolation. (2) The fact that ¢2 is controlled only by aM demonstrates the absence of any detectable c o m p o n e n t of charge inside the organic monolayer, and supports the initial assumption that there is no significant coadsorption when OH-, For SO4~- are the anions of the supporting electrolyte. The charge balance is of course different when coadsorption is allowed [ 12]. (3) More fundamentally, the determination of ¢2 potentials from ET shifts may constitute a new method of investigating diffuse layers which is simpler and more direct than most other classical methods. As has been already observed, the transition potential is markedly temperature dependent. The coefficient--17 mV/°C given by Greef [4] refers to solutions kept at a constant concentration of isoQ. Jenard and Hurwitz [6], using
32 TABLE 5 Transition potential corrected for the corresponding variation of ~2 when the concentration of the electrolyte is changed Electrolyte concentration [Na2SO4 ]/M
ET/V
~2/V
(ET--~2)/V
0.5 0.05 0.005
-0.730 -0.785 --0.840
-0.028 -0.079 -0.132
-0.702 -0.708 -0.708
saturated solutions of isoQ in 1 M KCI and in 0.1 M Na2CO3, have observed a slope o f - - 1 4 mV/°C, which is very similar to our own value, obtained in ½M Na2SO4 between 5 and 35°C. All other conditions being maintained constant, the variation of the interfacial tension with temperature is given by d3' = --s dT. As for the variation of the transition potential with temperature, this may be expressed in terms of the relative vertical shift of the electrocapillary curves which characterize the two distinct molecular films. The translation of E w is thus an indirect function of the difference of superficial entropy between the two monolayers. It has been found that for c = 0.5c~ the superficial entropy difference amounts to +2.3 × 10 -a J cm -2 K -~, when the transition is crossed from positive to negative potentials. This trend meets the intuitive idea that organization increases with compactness. The existence of a phase transition can be explained in the simplest terms by stating that there is total "incompatibility" between two different superficial states. In other words, this means that each of the two stable states possesses a free energy lower than that of any possible mixed structure. In the present case the transition involves two adsorption layers, each of them obeying its own Frumkin-type isotherm, associated with a large, positive interaction coefficient. Such large coefficients reflect the compact assemblage of rigid molecules which, being given the electrical field imposed by the electrode, can organize themselves through the interplay of various lateral orientation-promoting forces. The resulting layers constitute defined, condensed, superficial phases, either liquid, mesomorphic or solid (more precise evidence for the existence of given states will be available in the papers dealing with the kinetics of the transition and the inhibition power). On the other hand, interaction between molecules representative of regions I and II must be considerably weaker in view of the steric and electrical mismatch at the molecular level, and the impaired ability to mobilize forces strong enough to overcome the greater stability of the two prevailing condensed structures. As a result, mixed films are highly unstable, and can only be formed (and their decay observed), by resorting to sophisticated instrumentation. They correspond to the forbidden part of an S-shaped isotherm which, instead of describing the competitive behaviour of a surfactant and a solvent as is usually the case, here refers to the antagonism between two different monolayers, each of them made of the same molecules.
33 In the range of less negative potentials and lower concentrations, compact films are no longer present, and the behaviour becomes much more complex, as demonstrated by Figs. 2--5. Rather unexpectedly, the trends shown in the C, OM and F plots are strikingly similar to those found with Q. For both surfactants, at concentrations of medium and low values, there is a gradual desorption towards positive potentials, while, for concentrations close to saturation, the superficial excesses increase over the limiting value characterizing the adjacent compact film. These two opposite types of behaviour explain why, with increasing concentrations, the capacity peak shifts first towards less negative potentials, to reverse its deplacement as soon as the concentration is sufficiently large {about 0.3Cs). As was the case for Q, the main characteristics of the adsorption can be made most salient by constructing the Esin--Markov diagram (Fig. 7), which plots the potential shift as a function of the superficial excess, for discrete charge density values. Regions I and II are confined to the two vertical lines I and iI, which are separated by a space which corresponds to the forbidden part of the isotherm. As analysis of the left-hand side of the diagram can be based on the same arguments which have been fully developed for Q, they do not need to be repeated at great length. The major conclusions are as follows: (1) The fan-like array of curves originating from the O--O coordinates covers a region (zone A) which consists of an aqueous dilute film of molecules lying
+2
,4 *6
/,
+ 0.2-
z,//i o-,
o.o
0
~
1010I-'/MOi CM-2
Fig. 7. Plot of the shift of potential as a function o f the superficial excess for various charge densities, as indicated.
34 fiat. There is no appreciable potential shift with F when the charge density is close to zero, as can be expected for a surfactant having no dipolar component perpendicular to the surface. (2) Zone B can be explained in terms of a mixed structure composed of water and variable amounts of fiat and standing molecules (to the exclusion of intermediate-tilted orientations, as indicated by the breaks in the curves). (3) The upper left-hand side of the plot is consistent with the gradual development of a new state, characterized by abnormally high superficial excesses, lower capaci£y values and a conspicuous reversal of the potential shift. The latter effect suggests that the molecules embedded in the multilayer fraction of the film organize themselves in a way which tends to cancel (or at least to lower) their dipole contribution to A ( E - ¢2). The existence of this zone has recently been established independently [4]. That there is a nearly complete loss of identity between Q and isoQ (and also, as will be shown in Part IV, for most of their mono-methyl derivatives) may at first seem puzzling. It can, however, be understood by considering that the smaller the normal vector of the dipole moment, the less distinct is the differentiation, so that aromatic surfactants tend to display a common behaviour, as far as they are lying flat and dilute on a surface, and behave as a dielectric with a common thickness. The gradual change from this state to a full compact layer requires about the same charge density for both molecules, and since their dipole moments are commensurate, the potential span required for the reorientation process (and therefore the height and width of the capacity peaks) is also nearly identical. Closer examination of the diagrams representative of the two surfactants reveals, however, some slight differences. Superadsorption is somewhat larger for Q, for which zone C is more extended in the Esin--Markov plot and corresponds to larger A~, and smaller differential capacities. At intermediate and lower coverages, Q still remains more adsorbed than isoQ, as indicated by comparison between their respective isotherms and by additional evidence of a kinetic nature, which will be presented at a later stage
[71. As for quinoline [ 1], it has been preferred to avoid undue repetitions by deferring the general conclusions to Parts IV and V, as the complementary study of several methyl-derivatives of Q and isoQ {coupled with the elucidation of the inhibition processes which are observed in their presence) will then be at hand, with new elements of discussion. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12
CI. Buess-Herman, L. Giexst and N. Vanlaethem-Meur~e, J. Electroanal. Chem., 123 (1981) 1. H. Wefler-Feflchenfeld, Yuh-Lin Mao and E.D. Bergmann, Isr. J. Chem., 9 (1971) 111. J.P. L a m b e r t , D o c t o r a t e Thesis, Univer~it~ Libre de BruxeHes, 1973. M.W. H u m p h r e y s and R. Parsons, J. Electroanal. Chem., 82 (1977) 369. N. V a n l a e t h e m , J.P. Lambert and L. Gierst, Chem. Ing. Tech., 44 (1972) 219. A. Jenard and H.D. Hurwitz, J. Electroanal. Chem., 70 (1976) 27. G. Quarin, CI. Buess-Herman and L. Gierst, J. Electroanal. Chem., 123 (1981) 35. M.M. Potashnikov and I.G. Belavina, Zh. Prikl. K h i m . , 38 (1965) 1585. CI. Buess-Herman, G. Quarin and L. Gierst, J. Electroanal. Chem., Part IV, to be published. A.A. Survila and B.B. Damaskin, E l e k t r o k h i m i y a , 3 (1967) 1138. B.B. Damaskin, E.V. Stenina, V.A. Yusupova and N.V. Fedorovich, E l e k t r o k h i m i y a , 8 (1972) 1409. C1. Buess-Herman, G. Quarin and L. Gierst, J. Electroanal. Chem., Part VI, to be published.