SYEDM, AHMED
3546
Studies of the Double Layer
at
Oxide-Solution Interface
by Syed M. Ahmed Mineral Sciences Division, Mines Branch, Depavtment of Energy, Mines and Resources, Ottawa, Canada (Received January 98,1969)
Data have been obtained on the following double-layer characteristics of AI&, ZrOz, and Tho, by measuring the equilibrium distribution of the potential-determiningions (H+, OB-) at the oxide-solution interface as 8 function of pH, and of the ionic strength of the solution: (a) the zero point of charge, (b) the surface-charge densities, q*, of 8 1 2 0 3 in KNOg, KCI, and NaC104solutions and q+ of ZrOz and ThOz in KC1 and NaC104solutions, (c) the differentialcapacity of the double layer on AlzO, and (d) the change in the interfacial energy resulting from ionic adsorption. Results of the present work and of the previous work on other oxides (SiOz, ZrOz, Thoz,FesOa,Fe304,SnOz,and TiOz)have been correlated t o elucidate the general mechanisms of surface-charge formation and of ionic adsorption on oxides. The anions or cations may be specifically adsorbed on oxides through a basic or an acidic dissociation of the surface hydroxylgroups, respectively. An isotherm has been derived for the specific adsorption of K+ on the cathodic surfaces of oxides and the AG’K+ has been calculated from this isotherm for different oxides. The variation of AGOK+ with pH, with ( C K N O ~ and , with surface coverage and also the relationship of A G O K + with the t potentials of oxides has been discussed.
Introduction A number of studies have been made of the double layer on 0xidesl-j by measuring the equilibrium distribution of the potential-determining ions (H+, OH-) a t the oxide-solution interface, as a function of their electrochemical potential (pH) and of the ionic strength of the solution. This paper reports studies of the double layer on AlzOa, ZrOz, and ThOz. These studies are a continuation of previous work on a series of oxides (SiOz, ZrOz, Th02,3 Fez03,Fe804,eSnOz, and TiO2‘) reported earlier from this laboratory. The general properties of the oxide-solution interface have also been examined in this paper by correlating the J results obtained on the above aeries of o x i d e ~ . ~ l ~An attempt is also made to obtain the charge densities and the standard free energies, for the specific adsorption of K + on some of the above oxides. The present work on oxides is based on the same well-established principles as those employed for studying the AgI*-lo and the Ag2S11,12systems. However, as the oxide-solution system is much more complex than the AgI and the Ag2S systems, modified experimental techniques have had to he used in the present work. Hence the theory and assumptions involved in this work with oxides mill be outlined first.
Theory The occurrence of neutral and charged (*) surfaces at the oxide-solution interface has been attributeda,’ to the formation of metal aquo complexes, as shown schematically in (I). The negative surface charge originates from acidic dissociation of the surface hydroxyl groups (at p H > epc) and increases with increasing pH and also with increasing cation concentration. The cations (K +, in the present work) may be adsorbed on the oxide layer T h e Journal of Physical Chemistry
E+
[neutral compIex1
E-
either in a hydrated form or in a form dehydrated in the direction norma’l to the surface, The adsorption of K + in the dehydrated form on oxides is comparable to the well-known specific adsorption described in double-layer theories. l S - l 5 (1) G.A. Parks and P. L. De Bruyn, J . Phjls. Chem., 66,967 (1962) (2) J.Lyklema, J . ElectroanaE. Chem., 18,341 (1968). (3) S.M . Ahmed, Can. Z.Chem., 44, 1663,2769(1966). (4) G.Y. Onoda, Jr., and P. L. De Bruyn, Surface Sci., 4,48 (196%). (5) S. M. Ahmed and D. Maksimov, Mines Branch, Research Report R 196, Department of Energy, Mines and Resources, Ottawa, 1968. (6) S. M.Ahmed and D. Maksimov, Can. J . Chem., 46,3841 (1968). (7) 8. M. Ahmed and D. Maksimov, J. Colloid Interfac. Sci., 29, 97 (1969). (8) E.L. Mackor, Rec. Trav. Chim., 70, 663,747,763 (1951). (9) J. Lyklema and J. Th. G. Overbeek, J . Colloid Sci., 16, 595 (1961). (10) B. H.Bijsterbosoh and J. Lyklema, ibid., 20,665 (1965). (11) W.L. Freyberger and P. L. De Bruyn, J . Phys. Chem., 61, 586 (1957). (12) I. Iwasaki and P. L. De Bruyn, ibid., 62,594 (1958). (13) D.C. Grahame, Chem. Rev., 41,441 (1947). (14) P. Delahay, “Double Laysr and Electrode Kinetics,” Interscience Publishers, Inc., New York, N. Y . ,1965,pp 17-80.
SYMPOSIUM ON INTERFACIAL PHENOMENA
3547
The behavior of oxide surfaces on the anodic side of the zpc is best explained on the basis of proton addition to the neutral aquo complex, together with or without the replacement of the surface hydroxyl groups by the anions present. Thus in mechanism A, the anions do not replace the surface hydroxyl groups and stay as counterions outside the primary hydration is then shell of the surface. The oxide electr~de’~,~i’ reversible to H + and the positive-surface-charge densities (q+) depend primarily on U H + and not on the anion concentration as such. However, on increasing the anion concentration, at a given pH, any increase in q+ will be determined solely by the effect of the increased ionic strength on the structure of the diffuse double layer, and on u H + itself. From the theory of the diffuse double layer,13-15 the increase in p+ values on increasing the ionic strength of the solution from 0.001 M to I M of a 1: 1 salt is only 1-2 pC/cm2 for a double-layer potential of about 100 mV.18 As this variation in q+ is within the present experimental error (=t1 &), the q+ values will depend solely on
system in presence of 1: 1 potassium salts, a t constant temperature and pressure, is written as
dr =
-(rOH-
-
rH+)d,EHx - (rK+’dPKX) (rX-’dpKX) (3)
where X is an anion and other terms have their usual meaning. From the above equation, the following standard relationships for the differential capacity, C ( h ) , of the double layer and for the change in the interfacial energy (y) have been derived.1,3,8,9,12,1a dy
=
- (Q-dE),EKx
(4)
hence (5) ’KX
and
aH+.
I n mechanism B, in addition to the chemisorption of H + on the neutral, aquo complex, the anions replace the surface hydroxyl groups through a basic dissociation and are themselves chemisorbed on the metal atoms of the surface. This process, in the acid-base titrations (see Experimental Section), would liberate more OH- from the surface and result in higher (apparent) q+ values in comparison to those in mechanism A. These apparent q + values depend not only on aH+ but also markedly on the anion concentration. The above behavior appears to be characteristic of all the transition metal oxides. For other oxides, such as quartzj3 neither the complex formation with H + nor the basic dissociation of the surface hydroxyl groups may occur. Hence, no excess of positive surface charge would result on such oxides. All the above possibilities have been encountered in the series of oxides investigated in this work.
Thermodynamics If the surface structure at the zpc is taken as the reference plane, then ( r H + - roH-)zpo = 0, where rHT or r O H - is the surface excess, in mo1/cm2, of the potential-determining ion adsorbed on the surface. At any other pH the effective or excess surface charge is
4’ = F(rH+-
rOH-)
(1)
Thus q is positive or negative depending on whether r H + > or < r O H - . At a given pH, the potential, E, of the double layer on oxides relative to the zpc (the reference plane) is calculated from the Nernst equation
E
=
-0.059(pH - pH,,,) at 2 5 . 0 ”
(2)
The Gibbs adsorption equation, applied to the above
where yo is the interfacial energy at the reference plane (zpc). The significance of these parameters will be discussed later in relation to the present work. At constant pH or half-cell potential, E*
at the cathodic surfaces and
at the anodic surfaces. The excess surface charge, q+ or q-, is obtained directly from the present experiments, so that, from the principle of electroneutrality q* = -p(ri
+
rd)
(9)
where Ti and rdrefer to the relative, ionic ( x = 1) excess in the inner and in the diffuse double layers, respectively. As evidence for the validity of the above relationship, the q- values for quartz obtained in this on work3 by potentiometric titrations, and the I”&+ quartz, obtained by Li and de Bruynlg using 22Naas a tracer, are in good agreement (-9 pC/cm2, a t pH 10 in 0.001 M KN03 or NaC1). The double-layer charge due (15) M.A. V. Devanathan and B. V. K. S. R. A. Tilak, Chem. Rev., 65, 635 (1965). (16) G. Korttim and J. O’M. Bockris, “Textbook of Electrochemistry,” Vol. I, Elsevier Publishing Co., London, 1951, p 293. (17) J. T. Stock, W. C. Purdy, and L. M . Garcia, Chem. Rev., 58, 611 (1958). (18) C. D. Russel, J . Electroanal. Chem., 6,486 (1963). (19) H. C. Li and P. L. De Bruyn, Surface Sci., 5,203 (1966).
Volume 7’9, Number 11 November 1969
3548
to Ca2+ adsorption on quartz20obtained by using 45Caas a tracer is also found to agree closely with the Q- values of quartz3. Hence, if rdis known experimentally, or calculated from the diffuse double-layer theory, the density of the specifically adsorbed ions as a function of pH and the activity CCKX may be obtained by difference using eq 9. From the adsorption data, it is then possible to obtain the standard free energy (AGO) for the specific adsorption of ions on oxides using a suitable adsorption i ~ o t h e r m . 1 ~ ~These 2 ~ AGO values (apparent), for a given composition of the solution, represent the work done in transferring the ion from the bulk solution to the interface and, therefore, also provide information on the effective potential of the compact layer (relative to the bulk solution) that the adsorbate ion encounters as it approaches the adsorption site. Hence, these standard free energy values are expected to be closely related to the ( potentials of oxides. A study of the variation of A G O with pH, with electrolyte concentration, and with surface coverage mould be particularly helpful in understanding the oxidesolution equilibria and also in comparing the reactivity of different oxides. The use of Langmuir adsorption isotherm in calculating AC" for the adsorption of K + on oxides will be considered in the last section.
Experimental Section Materials. Crystalline a-AlzOa in a powder form and ZrOz and T h o z in fused forms were supplied by the Aluminum Co. of Canada and the Norton Co., Chipawa, Ont., respectively. These oxides in - 100 150 mesh size were passed through a Frans magnetic separator to remove any magnetic impurities and were subsequently leached with hot 20% HE02 in an allglass Soxhlet. The resultant material was washed free of excess acid, d e ~ l i m e d and , ~ stored in degassed conductivity water that was replaced frequently. From spectrographic analysis, the A1203 and ThOz were found to be 99.9% pure, while the ZrOz contained 0.5% trace impurities whose interference was assumed to be negligible. From X-ray diffraction analysis, ZrOz and Thoz were found to be single-phase baddelyite (monoclinic) and thorianite (cubic), respectively. a-&03 may be considered either as rhombohedral or hexagonal in structure, where each A1 is surrounded by six oxygen atoms arranged octahedrally. RecrystalIized AR grade KNOa, KC1, or NaC104 were used to &just the ionic strength of the solution while the pH was adjusted, as required, using CO2-free KOH, Hx03, KC1, or HC104 depending on the salt used. The electrolyte solutions were freshly prepared before each set of experiments using degassed conductivity water and were kept under an argon atmosphere before use. Any dissolved oxides of nitrogen in the "03 were removed by bubbling argon through a hot solution of the acid that was subsequently preserved in a dark battle a t low temperatures.
+
The Journal of Physical Chemistry
SYED1c1. AHMED
Method, The adsorption studies were carried out in a small Pyrex glass cell that was provided with a water jacket for maintaining constant temperature (25.0') and with two side tubes for circulating argon during the experiment. The cell was also closed to the air by a flexible covering. The pH of the solution was measured with a Beckman Research Model pH meter using a small glass electrode and a saturated calomel reference electrode. The pH meter was enclosed in a constant-humidity (50%) box, leaving the electrodes outside, and could be read with a relative accuracy of 0.002 pH unit. Beckman buffers, prepared according to the KBS standardsjZ2were used for standardizing the pH meter. The pH is henceforth referred to as pH,. The calomel electrode was separated from the KKOa solution using a KN09saturated solution bridge with a fiber junction. When using NaC104 solutions, a solution bridge containing KN03 and XaN03 in the molar ratio of 1:7 was used.8 The liquid-junction potentials in both these cases are small. The calibration of pH, against molar conHC1, or HC104 was carried centrations of KOH, "03, out separately in 0.001, 0.1, and 1 M solutions of the respective potassium salts (n'aClO4 used for HC104 and KSOS for KOH). These calibration curves were used later in calculating the charge densities. As these curves were obtained under constant conditions of electrode combinations, small shifts may arise in the potential scale on increasing the ionic strength of the solution, but the calculations of the charge densities should remain unaffected. The total shifts in the half-cell potentials, as measured by the glass electrode, on increasing the ionic strength from 0.001 to 0.1 and 1 M were also determined and taken into account where necessary. The stored oxide samples were taken directly out of water, rewashed, and almost all the excess water filtered out under suction in an atmosphere of argon. The resultant moist material (water content -2.5% by wt of oxide) was used for further work. A 10-ml sample of the electrolyte solution was placed in the reaction vessel and any dissolved COZwas expelled by rapid stirring in a current of argon and the variation of pH, on stirring the blank solution was then followed at intervals of 2 min. After the solution attained a steady pH,, an oxide sample (2-3 g), prepared as above, was added to the solution, and the variation of pH, with time was again followed, normally at intervals of 2 min but at intervals of 1 min near the zpc. From kinetic studies of the oxide-solution interface, it has been shown6 that almost all the change in pH, occurs during the first 2-6 min after adding the oxide to the (20) S. M. Ahmed and A. B. Van Cleave, Can. J . Chem. Eng., 43, 23
(1965). (21) E. Bloingren and J. O'M. Bockris, J . Phys. Chem., 63, 1475 (1959). (22) R.G.Bates, J . Res. Nut. Bur. Stand., A66, 179 (1962)
SYMPOSIUM ON INTERFACIAL PHENOMENA solution. This initial change was followed by a' slow variation in pH, that was attributed to the solubility of the oxide and to slime formation. The rapid change in pH, occurring in the first few minutes after addition of the oxide to the solution has been attributed6J to the reaction of the oxide surfaces with the electrolyte. The amount of oxide dissolved in solution during the course of each experiment (2-6 min) was found to be below detection limits using colorimetric methods (sensitivity 1 pmo1/1.). It should be emphasized, however, that although the amount of oxide in solution a t any given time during the experiment may be very small because of the low solubility product of the metal hydroxides, the total pH, changes due to the cumulative solubility effects (such as dissolution, dissociation, and reprecipitation) could be far beyond this limit and such as to mask the primary surface effects. In order to avoid the above solubility effects, fresh samples of oxides (coarse and crystalline) and of solutions were used for each of the charge-density values obtained, so that the oxides were not allowed t o contact the test solution for more than 2-6 min. This procedure is comparable to using fresh surfaces of mercury in electrocapillary work. The surface areas of the oxides were determined by the krypton gas adsorption methodz3 and were 4414 cm2/g for A1203,1438 cm2/g for ZrOz, and 192 cm2/g for Thoz. These areas were determined on oxides which had first been given a blank stirring treatment in water under the same conditions as used in the adsorption experiments.
3549 -60
-50
-40
-
-
0
IM
V
0.IM
X
DOOIM
-30 -
.0 NaCI04
1
3 t too
I
I
I
I
16 0
PH
Figure 1. Variation in charge density (q*) on AlzOawith final pH, and potential difference relative to the zpc in KNOI, NaC1O4, and KCl solutions.
Table I : Zpc of Oxides in KN03 Solutions KNOa,
Results The addition of oxides to the electrolyte solutions resulted in a decrease or an increase in pH, depending on whether the initial pH, of the solution was greater than or less than the zpc of the oxides. Thus, the oxides acquire a negative or a positive surface charge in solutions whose pH, is greater than or less than the the pH, of the zpc, respectively, according to mechanisms already discussed. I n order to calculate the surface-charge densities, the initial and final values of a H t (below pH, 7) or U O H - (above pH, 7) were corrected for the pH, variations of the blank and for the solubility effects. The a H t and U O H - values were converted into concentration terms (CH+ and COH-) using the corrcsponding calibration curves. From these concentration values, (ACH+ or ACoH-)/lO ml/cm2 of oxide and hence the surface-charge densities, q*, were calculated. The results of q* for A1203 in 0.001, 0.1, and 1 M KX03, KC1, and NaC104 solutions are shown plotted against the final pH, and the potential relative to the zpc in Figure 1. Similar results of q+ for ZrOz and ThOz in KCl and NaC104 solutions are shown in Figures 2 and 3, respectively. The results of q* for the same ZrO2 and ThOz samples in KN03 solutions have also been reported earlier.8
I
6 1 7 81 9 I 10 -100 -200 -300mV pH AND POTENTIAL IN mV RELATIVE TO Z.PC.
4
concn, M
Oxide
Quartz ZrOz ZrOz ThOn Tho2 Fez03 Fez03 FezOa FeaOa 8nOz SnOn Ti02 Ti02 TiOz a-AlzOa a-AlzOa ~-AlaOa
Dry samples Dry samples Dry samples Dry samples Dry samples Specular hematite] moist Specular hematite, moist Specular hematite, moist Magnetite, moist Cassiterite, moist Cassiterite] moist Rutile, moist Rutile, moist Rutile, moist Moist Moist Moist
1
PIIS of ZPC
0,001
3 . 6 i 0.05 5 . 5 i: 0.05 6 . 2 i 0.05 5 . 9 rrt 0.05 6 . 8 Z!Z 0 . 0 6 5.3 f 0.05
0.1
5 . 4 f 0.05
1
5 . 7 i: 0 . 1
0,001 0.001 0.1,1 0,001
6 . 4 i0.1 5 . 5 zk 0 . 1 5.4 5 0.1 5 . 3 f 0.05 5.0 f 0.05 4.8 f 0.1 4.7 i0 . 1 4.8 5 0 . 1 5.0 & 0.1
0.001 1 0 001 I
1
0,l 1 0,001 0.1 1
The differential capacity, C ( & ) , of the double layer was obtained by the graphical differentiation of the (23) S. M.Ahmed, Mines Branch Technical Bulletin TB 84, Department of Energy, Mines and Resources, Ottawa, 1966.
Volume 73, Number 11
November 1060
SYEDM, AHMED
3550
30
20
IO
T
N N
z
o
\ V
V
=I 30
-z
KCI
W
c3
IL
U
5
20
IO
0 .
4
5
6
PHS
Figure 2. Variation in charge density (p+) on ZrOl with final pH, in NaClOd and KC1 solutions.
smoothed charge-density plots and these results for
A1203, and also for SnOz and TiOz (from previous work'), are shown in Figure 4. The (y - yo) values for A1203 were obtained by graphial integration of the charge-density plots and are shown in Figure 5 . Discussion
The Zero Point of Charge. The pH, values obtained in this work315-7 for the zpc of various oxides are summarized in Table I. These zpc values were obtained for the primary equilibrium (2-6 min) between the oxide surfaces and their potential-determining ions in electrolyte solutions that were practically free of their dissolved metal complexes. The values listed in Table I for the zpc of oxides, as shown previously,et7 are in close agreement with the pH of the zero streaming potentials (f potentials) for the same oxides where fresh solutions, free of dissolved metal complexes, are brought into contact with the oxides. The zpc of oxides obtained during prolonged oxide-solution equilibria (-15 hr or more) are, however, different from the above values and are reported to be identical with the isoelectric points of the dissolved metal complexes. Further problems concerning the zpc of oxides have been discussed in detail elsewhere.
-'
The Journal of Physical Chemistry
On increasing the KNOs concentration from 0.001 to 1 M in a blank experiment, the half-cell potential was found to shift to a lower pH, by 0.3 rt 0.05 pH, unit near the zpc. I n order to compensate for the above shift in the half-cell potential, a corresponding shift in the zpc of oxides to a higher pH, is expected, Such a shift in the apc of A1203 and FezO3to a higher pH, may be seen in Table I and this shift does not indicate any specific adsorption of K t or NOa- on these two oxides at the zpc. However, similar shifts in the zpc of ZrOz and ThOz to a higher pH, (exceeding 0.3 pH, unit) and in the zpc of SnO2 and Ti02 to a lower pH, indicate specific adsorption of NOa- on ZrOz and Tho2 and of E(+ on SnOt and TiOz. Further evidence of the above effects is available from th, charge-density and differential-capacity values shown in the next section. The Positive Surface Charye Densities. The pi surface charge densities of the following oxideE been found to be practically independent (i1 the concentration of anions under investigation : (Figure 1) and TiOZ7in NOa-, C1- and Clod-, and and Fez036 in NOa- and Clod- solutions. Hen the anodic region, the behavior of these oxides fc mechanism A (see Introduction) where H + are c sorbed on the neutral surface giving rise to a posii charged surface, and the anions stay as count6 outside the primary hydration shell of the surface.23 terms of double-layer theories,la-l6 the anions ir above cases may be said to be nonspecifically adsc on the metal atoms of the surface. Although tl values, at a given pH,, may differ from one oxic another in the above cases, depending on the na of the coordination of the surface complexes, increase in the q+ values with the increasing a concentration at a given pH, (cJ eq 8) will, howe , be determined by the diffuse double-layer theory. The behavior of the semiconductor and magnetic oxides in the q + region was, however, found to be exceptional. Thus, although the q+ values of specular hematite (Fe203,an n-type semiconductor) in NO3- and Clodsolutions6 were found to be independent of the NO,and Clod- concentrations, these q+ values were much higher (-30 pC/cm2 130 mV) than the qf values for other oxides. These high values of q+ indicate a much closer packing of anions (NOa- and Clod-) on the Fez03than for other oxides. The behavior of magnetite6 (an n-type semiconductor, and also highly magnetic) in the q+ region was also found to be inconsistent with the general behavior of the other oxides. Hence the semiconductor and magnetic oxides require (23a) NOTEADDEDI N PROOF. Recent investigations of Al(OH2)a3+ complexes by nmr spectroscopy have shown that this complex is stable in C1-, NOa-, or Clod- solutions with no detectable re placement of HzO by C1- and that proton exchange reactions are predominant between the complexes and the acid solutions (D. W. Fong and E. Grunwald, J. Amer. Chem. Soc., 91, 2413 (1969)).
3551
SYMPOSIUM ON INTERFACIAL PHENOMENA
KCI
5
4
6
PHS
Figure 3. Variation in charge density ( q + ) on ThOz with final pH. in NaClO, and KCI solutions.
70 -
8 2a
60-
5
50-
3 IA
lL \
%
> t
40-
2
t
0
B
8
-60
-
-70
-
30-
-1
?
z
w a 20w
-80
LL
3
I
I
I
1
4
I
6
7
;. 8
P LL
PHI AND
IO / I
I
I
I
4
5
6
7
-.-
"Z03
I 0
I 9
,
(
IO
1: IIpHi
POTENTIAL IN mV RELATIVE TO Z.RC,
Figure 5. Variation in (y - YO) at the Al~O~-solution interface with final pH, and potential difference relative to the epc.
PHS
Figure 4. Variation in the differential capacity, C ( i ) , of the double layer on AlzOs, SnOa, and Ti02 in KNOs with final pHs
a more detailed study, particularly with reference to to the space-charge region inside these solids. The qf values of ZrOz and Tho2 in Clod-, C1-
(Figures 2 and 3) and in NO3- solutions, and the q+ values of FezOa6and Sn027 in C1- solutions have been found to increase not only with a H t but also markedly with the particular anion concentration. Thus the slopes, (bq+/bpm&t, as given by eq 8, were found Volume 76,Number 11 November 1069
S Y ~ DM.AHMRD
3552 to be much higher than expected from diffuse doublelayer theory for the nonspecific adsorption of ions on solid surfaces. Hence, in all the above cases the increase in q+ with increase in the salt concentration is due to the specific adsorption of the given anions on the metal atoms of the surface through basic dissociation of the surface hydroxyl groups. The type of anionic adsorption (specific or nonspecific) on different oxides is summarized in Table 11. As an
Table 11: Nature of Anion Adsorption on Oxides Anions nonspecifioally Oxide
Si02 quart^)^ zroz3
Th023 TiO? A1203
sno27
Fe2036
adsorbed
N03-,
c1-, clod... ...
Anions speoifically adsorbed
... NO*-, Cl-, ClodNos-, C1-, C10,-
*.. ...
Nos-, Cl-, clodNOa-, el-, clodNOa-, ClodNOa-, ClOi-
c1c1-
extreme case, however, if the specific interaction of anions with the surface-metal atoms is very strong, then a saturation limit in anionic adsorption may be reached at a low anionic concentration, so that a further increase in the anion concentration would show little increase in the Q+ values, e.q., C1- on Sn027. T h e Negative Surface Charge Densities. The negative surface charge on oxides, as discussed earlier, originates from the acidic dissociation of the surface hydroxyl groups and leads, subsequently, to the adsorption of E(+ on oxides. The g- values of A1209 (Figure 1) and hence the adsorption of K + on this oxide (and on other oxidesa~~-'in general) increase with the increasing pH, as well as with increase in the E(+ concentration in solution. The relative surface excess, I'K+ij due to the specific adsorption (as defined in the Introduction) of E(+ on oxides was obtained for various oxides by subtracting the values of r K + d (the cation excess in the diffuse double layer) from the Qvalues. These rgci values are shown plotted against the activity U I ( N O ~ in Figures 6 and 7. The rI(+dvalues were calculated18 from the theory of the diffuse double layerla,14knowing the oxide-electrode potentials relative t o the zpc and the ionic strength of the solution. This procedure may not be precise but the error involved is less than the experimental errors usually encountered in this type of work. These adsorption isotherms for K + on oxides (Figures 6 and 7) are comparable in shape and magnitude to the isotherms for the specific adsorption of and of T1+ 28 on Hg. AG' for th,e Specific Adsorption of K + on Oxides. As the pH, of the solution is increased relative to the zero point of charge of the oxide (cathodic side), the OHThe Journal of Physical Chemistry
Figure 6. Variation in the specific adsorption of K + on FezOs, FesO,, SnOz, and A1203 with the log of the mean activity of KNOs, a t different pH..
-
2
-50-
0
a +^-4OY
-3
-2
-I log a + KN03
0
-
Figure 7. Variation in the specific adsorption of K + on Ti02 with the log of the mean activity of KNOa, at different pH..
ions from the solution may be considered, as a first step, to neutralize the surface H + ions to form water molecules which are initially adsorbed on the oxide surface. The Kf from solution subsequently replace the water molecules and are specifically adsorbed on (24) D.C.Grahame and R. Parsons, J . Amer. Chem. SOC.,83, 1291 (1961). (25) D.C.Grahame, ibid., 80,4201 (1958). (26) J. Lawrence, R. Parsons, and R. Payne, J . Electroanal. Chem., 16, 193 (1968). (27) H.Wroblowa, Z. Kovac, and J. O'M. Bockris, Trans. Faraday SOC.,61,1623 (1965). (28) G. G.Susbielles, P. Delahay, and E. Solon, J. Phys.Chem., 70, 2601 (1966).
SYMPOSIUM ON INTERFACIAL PHENOMENA
I
0
I .05
I 0.I
3553
I
I
O*lS
042
I 0.25
I
I
0.3
0.35
I 0.4
I
0.45 eKt
SURFACE COVERAGE Figure 8. Variation in the AGO^^ for the specific adsorption of K f on the cathodic surfaces of oxides: (A) with surface coverage OKt ; (B) with potential (relative to zpc) and with KNOs concentration, for SnOl.
the oxide. The apparent A G O for the K + replacing the water molecules, at a ,given pH, and U K N O ~ , is given21by
where e is the fractional surface coverage due to the specific adsorption of K + on oxide and U K N O ~is the mean activity of the salt in the bulk solution. For calculating 8, a surface charge of 65 pC/cm2 is assumed to correspond to an effective saturation coverage. The apparent free energy values, as obtained from eq 10, for the adsorption of K + on oxides vary systematically with pH, (or potential relative to zpc) as well as with U K N O ~ ?as shown in Figures 8A and 8B. These variations will be discussed shortly. One interesting observation from Figure 8A is that the AGO,, values (for a given agNos)of different oxides fall on the same line. This indicates that the differences in the free energy values of different oxides are not revealed from a physical parameter such as surface coverage. Instead, the differences in the free energies of different oxides appear in their egt values being different for different oxides for a given composition of the solution (Figures 6 and 7). The increase in - A G O , , with increasing pH, for cassiterite (Figure 8B) is typical of the oxides investigated and is obviously due to the ionization of the surface hydroxyl groups which increases the po-
tential difference between the oxide surface and the bulk solution. However, this pH, dependence of the AGO,, may be incorporated in the adsorption isotherm by taking into account the equivalent change in the electrochemical free energy, of the oxide-solution half-cell, that occurs during the neutralization of the surfacebound H+ ions by OH- (first step, see above). This change in the free energy of the half cell is given by - A G O H +
=
RT In
4'=
(&I
p f
(UH
+>
- pHi)
2.303RT(pHt
(11)
where i and f refer to the initial and final values of a H + or pH, during adsorption. By combining eq 10 and 11 -AG°Kt
=
2.303RT X
(Alternatively, A G O in terms of U K N O ~ and a H s O + may be obtained by considering replacement of HaO+ from the surface by E(+ and by equating their electrochemical potentials in the adsorbed layer and in the bulk solution.21 Then, by substituting CH~Oa H t for U H ~ O+ and including further changes in U H + during adsorption, the isotherm obtained is identical with eq 12.) In Figure 9A, the A G o g + for hematite and cassiterite, as obtained from eq 12, for the three U K N O ~
-
Volume 73, Number 11 November 1969
SYEDM. AHMED
3554
To 5
-
"
c
0 0
0.001M KNO3
0
0
00
9102
I
I
I
1
B
3
-2
Log Figure 9. Variation of
AGOK+
for Fen01 and SnOl: (A) with pHa at different
values, are shown to be independent of pH, within =t2% error. This deviation was somewhat larger (.t.5%) at extreme values of pH, because of higher experimental error at these pH, values and because the solubility corrections in pH; and pHf were also neglected. I n Figure 9B, it is further seen that these - A G o ~ +values decrease linearly (;.e., adsorption of K + becomes increasingly difficult) with increasing values of log CLKNO~. The slope of these plots in Figure 9B and the extrapolated values of AG'K+ to unit salt activity appear to be characteristic of each oxide and represent their reactivity for ionic adsorption. Hence, the AGOK+ at unit salt activity (--2.9, -3.8, and -5.7 kcal/mol for hematite, cassiterite, and rutile, respectively) may be taken as a standard value for comparing different oxides. A comparative study of the oxides in this way is being carried out. The AG"K+ values for each oxide are seen (Figure 9A and 9B) to vary by about 2 kcal during adsorption. Thus, the variation in AG'K+ for hematite from 5 to 3 kcal/mol, corresponds to a variation from 200 to 100 mV on the potential scale. These values in mV, say $Kt, represent the effective potential (of the plane of K + adsorption, relative to bulk solution) that a K + ion would encounter as it approaches the surface for adsorption from bulk solution. The variation of AGO,, and A G o ~ +(or the $ K A values) (Figures 8 and 9) and of the potentials of o ~ i d e s with ' ~ ~ pH, ~ ~ and art (salt) are qualitatively similar. The # K + and the { potentials of oxides are also of the same order of The Journal of Physical Chemistry
UKNO~;
-1 Q
(B) with log
0
KNO~
UKNO~.
magnitude, which provides some evidence for the validity of assumptions made in deriving the adsorption isotherm. For a strict comparison between the { potentials of oxides and the present #K+ values, the potentials of the particular samples used in this investigation would have to be determined. However, the potentials of oxides in general (-- 120 and -200 mVmaxfor quartzlg and Sn02,29respectively) appear t o be somewhat lower than the $ ~ tvalues, for a given composition of the solution. This difference is in qualitative agreement with the general assumption that the slipping planeg in electrokinetic measurements lies somewhat farther from the compact layer. It is hoped t o examine this subject further by parallel measurements of the { potentials. The Diferential Capacity and the (y - yo) Curves. The differential capacities, I?(*), of the double layer on Al208, SnO2,' and Ti02 are shown plotted against the final pH, in Figure 4. The data in Figure 4 and the previous plots3j6 of C against pH, exhibit# three main features that may be interpreted on the basis of the mechanisms involved. First, the increase in e(+) values with decreasing pH, (Figure 4) on the anodic side of the zpc is attributed to the chemisorption of H+ on the surfaces. A similar increase in the differential capacity of Sn electrodes with decreasing pH has also been observed recently from direct (29) D. J. O'Connor and A. S. Buohannan, Aust. J . Chem., 6 , 278
(1963).
SYMPOSIUM ON INTERFACIAL PHENOMENA capacitance measurements. 30 Second, in the cathodic region, a plateau-like break occurs in the C(-> against pH, plots of many oxides a t C 30-40 MFlcrn2, and is followed by a steep rise in C( -) due t o the increaeing specific adsorption of K + on the oxide layer. The plateau probably corresponds to the smoothed-out “hump” that usually occurs in the double-layer capacitance of Hg, if the specific adsorption of counterions is not i n t e n ~ e . ’ ~ - ~These j plateaus were found to be particularly prominent in the capacity plots of quartz3 and Fez03 and t o disappear for Thoz and A1203 (Figure 4) which is attributed to the intense specific adsorption of Kfon the latter two oxides. The third noteworthy characteristic is the occurrence of broad and flat minima in the differential-capacity curves of some oxides quart^,^ Fe203,6and A1203 (Figure 4)) at the cathodic side of their zpc. The occurrence of such flat minima in the double-layer capacity of certain metals, notably ZnJ3lhas also been observed by direct capacitance measurements. In the present case of A1203,for example, the broad minima in the C(-) values (Figure 4) have resulted from the flat portions in the Q- against pH, plots of Figure 1. This indicates that a considerable potential drop occurs across the hydration layer of the oxide (-130 mV for Al2O3)before proton dissociation occurs from the surface hydroxyl groups. I n electrocapillarity, such a potential drop across the hydration layer of solid surfaces is known as the x potential. This effect is most pronounced for oxide surfaces that are strongly hy,~ and A1203 (Figure 4))and drated, e.q., q ~ a r t z Fe203,‘j does not arise in the anodic region of Be203 or A1203, because the surface charge (+) on oxides in this region originates from proton adsorption instead of from proton dissociation. The (y - yo) against pH, plots of A1203(Figure 5 ) and other oxides3*jare similar in shape to the electrocapillary curves but have a different meaning. These ( y - y o ) values represent the integral values of the electrochemical work done in transferring the adsorbed ions from the bulk of the solution to the interface. Except for a small correction for the contribution of the diffuse double layer, in general the (y - yo) values of the oxides examined are of the
-
3555
same order of magnitude and have the same meaning as the “equivalent, effective surface pressure” calculated by Parsons and ~ t h e r P -for ~ ~the specific adsorption of anions on Hg. The single anodic branch of the (y - 70) against pH, plots represents nonspecific adsorption of anions on the metal atoms of the oxide surfaces and in this respect resembles the cathodic branch of the electrocapillary curves of Hg, where the cations remain nonspecifically adsorbed.
Conclusions As a result of some inherent uncertainties arising from factors such as pH,, the absolute values of surface areas and the solubility effects, this work is not sufficiently precise to investigate the properties of the diffuse double layer on oxides. However, it is evident from the present investigation that several major effects and the general trends in the properties of the oxide-solution interface can be studied using the present techniques, provided that the solubility effects are prevented from interfering with the surface effects. The usefulness of this work could be further enhanced be performing simultaneous adsorption studies using radioactive tracers and streaming potential studies on the same material. Some of the properties of the oxide-solution. interface considered in this paper are expected to be helpful in the future theoretical development of the subject. The applications of this work in practical fields have already been discussed elsewhere, s6
Acknowledgments. The assistance of D. Maksimov (Visiting Scientist, Skochinski Mining Institute, Moscow, U.S.S.R.) in the work6 on A1203,Zr02, and ThOt, and Relpful discussions with Professor B. E. Conway, University of Ottawa, and Dr. H. P. Dibbs, Mineral Sciences Division, are gratefully acknowledged. (30) N. A. Hampson and D. Larkin, J. Electrochem. Soc., 115, 612 (1968). (31) D. S. Brown, J. P.G. Farr, N. A. Hampson, D. Larkin, and C. Lewis, J . EZectroanaE Chem., 17,421 (1968). (32) J. M. Parry and R. Parsons, Trans. Faraday Soc., 59, 241 (1963). (33) R.Payne, J . Chem. Phgs., 42,3371(1965). (34) R.Payne, J. Electrochem. Soc., 113,999 (1966).
Volume 79,Number 11 November 1969
