J. Phys. Chem. 1996, 100, 11681-11687
11681
Electroacoustic Study of Adsorption of Ions on Anatase and Zirconia from Very Concentrated Electrolytes Marek Kosmulski† and Jarl B. Rosenholm* Department of Physical Chemistry, Åbo Akademi UniVersity, Porthangsgatan 3-5, 20500 Åbo, Finland ReceiVed: January 31, 1996; In Final Form: April 11, 1996X
The reduction of absolute values of negative ζ potentials (basic branch) of anatase with the concentration of 1-1 salts (alkali halides, nitrates, and perchlorates) is much more pronounced than that of the positive ζ potentials (acidic branch). The extent of this effect increases in the series Cs < K < Na < Li for a given anion and CH3COO < Cl < NO3 < ClO4 < Br < I for a given cation. For sufficiently high concentrations of most lithium and sodium salts, e.g., 0.53 mol dm-3 NaI, there is no isoelectric point (iep) and the ζ potentials are positive over the entire available pH range. For potassium and cesium salts, an iep is always observed, even at very high concentrations, but it is substantially shifted toward the higher pH values. Small cations show a differentiating effect: the course of ζ(pH) curves for particular lithium and sodium salts at a given high ionic strength is very sensitive to the nature of the anion, but the effect of the nature of the anion is relatively insignificant when different potassium salts are considered. Large anions (iodide) show a differentiating effect, while smaller anions (chloride) do not.
Introduction Very concentrated electrolytes are seldom visited by colloid chemistry for obvious reasons: both hydrophilic and hydrophobic colloids become unstable at high ionic strengths. Thus, applicability of methods and availability of data are limited. In spite of theoretical and experimental difficulties, the practical meaning of colloidal systems involving electrolyte solutions of high ionic strength is enormous. The different surface phenomena that occur in ocean water may serve as one example. The ζ potential, which is often used to characterize the electric field around a colloidal particle, could hardly be measured at electrolyte concentrations above 0.1 mol dm-3 using traditional equipment. The electroacoustic method, which has been recently developed, makes it possible to measure ζ potentials at very high ionic strengths. Once data are available, theories and equations using the ζ potential as a parameter may be applied and tested over a wider range of ionic strengths. At low ionic strengths the 1-1 electrolytes (e.g., alkali halides) show many common properties. An example is the fact that the isoelectric point (iep) of oxides and many other solids does not depend on the nature and concentration of these salts.1 Therefore, they are called indifferent electrolytes, in contrast with electrolytes involving multivalent ions, which cause a shift of the iep. The direction and magnitude of the shift depends on the nature and concentration of the electrolyte.2 At low ionic strengths, there are many water molecules per ion available, and the ions in aqueous solutions of indifferent electrolytes under these conditions are completely hydrated. The surface does not “see” the ions but only their hydration shells, and the individual properties of ions are screened by the surrounding water. Considering not only the primary but also secondary hydration sheaths, dozens of water molecules are engaged by one ion. At higher ionic strengths, far fewer water molecules per ion are available, and due to this water deficit, some ions experience partial dehydration. This dehydration allows the ions to approach closer to the surface. Moreover, † On leave of absence from Polish Academy of Sciences, Institute of Catalysis and Surface Chemistry of Interfaces, Laboratory of Adsorption and Surface Chemistry of Interfaces, Lublin, Poland. X Abstract published in AdVance ACS Abstracts, June 15, 1996.
S0022-3654(96)00315-2 CCC: $12.00
once the hydration shell is distorted, the individual properties of particular ions and salts are displayed in ion-ion interactions and in ion-surface interactions that are unimportant under lower ionic strength conditions. Therefore, in studies at high ionic strengths, substantial differences between particular “indifferent” electrolytes may be expected. In the present study, the ζ potentials of anatase were measured as a function of the pH and ionic strength (up to 1.7 mol dm-3) for various 1-1 salts. Although the ζ potentials characterize mainly the interfacial region, the results provide information about mutual interactions of ions and water molecules in the bulk concentrated aqueous solutions of these salts. Experimental Section Apparatus. The Acoustosizer from MATEC Instruments with version 1.04T software was used according to the procedure described in the user’s manual. The factory calibration, which accounts for the nonlinearity of the electric field in the cell (e.g., due to the presence of the probes), is valid for alkali metal salt solutions of concentrations up to about 0.1 mol dm-3. At higher electrolyte concentrations the relative conductivity of the solution on the one hand and the cell walls and probes on the other hand is different, so the results obtained with the factorycalibrated instrument may be quantitatively incorrect. However, due to the present instrumental (software) restrictions, the reprocessing of the data files can only be done when the appropriate conductivity correction procedure becomes available. The present software version uses viscosity η and density F of pure water in all calculations. Therefore the apparent ζ potentials provided by the instrument had to be recalculated for the high ionic strength systems, where the viscosity and density were considerably different from those of pure water. To calculate the real ζ potential, the apparent value must be multiplied by (η∆F′F)/(η′∆FF′), where ∆F is the difference between solvent and particle density and the prime denotes the values used by the Acoustosizer. To our best knowledge, the Acoustosizer also ignores the electrolyte effect on the dielectric constant of water, but we have not introduced any corrections related to these effects. Materials. Anatase from Aldrich (99.9%) was washed according to a procedure described elsewhere.3 The washed © 1996 American Chemical Society
11682 J. Phys. Chem., Vol. 100, No. 28, 1996 material shows an iep at pH ) 5.85, in good agreement with the literature.4 This result confirms the absence of substantial amounts of multivalent ions in titania. In contrast, the same material before washing showed an iep at pH < 3.4, probably due to anionic impurities. The BET surface area of anatase was 9 m2/g, and the mean particle diameter from TEM was 300 nm. Acid-base titrations of anatase (20% w/w dispersion) were carried out in different electrolytes at 20 °C. A limited number of titrations (NaNO3 and NaBr) have also been carried out with zirconia from Merck, washed in the same way as the anatase. Most salts were of very high purity, e.g., 99.995% NaNO3 from Aldrich. However, additional experiments with regular AR reagents show that purity of reagents is not a crucial factor in the observed phenomena. The sample was open to the atmosphere during the titration, and no special efforts were made to remove carbon dioxide. The Titration Procedure. First, the electrolyte background data were collected for different electrolyte concentrations. Then the cell was partially filled with water, and titania was dispersed at a very high stirrer speed (1200 rpm). The stirrer speed was then reduced to 400 rpm, water and dry salt were supplied to get 0.37 dm3 (the operational volume of the cell) of solution of the desired ionic strength, and the dispersion was brought to adsorption and thermal equilibrium at pH ≈ 4 (pH adjusted using 1 mol dm-3 solutions of acid or base with the same anion/ cation as the supporting electrolyte) within 1 h. The sample was then manually titrated with 1 mol dm-3 base to get an experimental point every 0.2-0.4 pH unit. The preliminary experiments show that the equilibrium is rapidly established: fast (5 pH units per hour) and slow (2 pH units per hour) titrations give essentially the same results. In the vicinity of the iep, the mixer speed was increased to prevent sedimentation. When pH ≈ 9 was reached, the suspension was mixed for 30 min and then titrated back with the acid to reach pH ≈ 4. More salt was added to increase the ionic strength, water was added if necessary (evaporation), and the titration procedure was repeated for gradually increasing ionic strengths. With sodium acetate, the lower pH limit was set at pH ) 6 (strong buffer effect at lower pH values), and the same lower pH limit was set with iodides to avoid their oxidation in more acidic media. Base and acid titrations gave essentially the same ζ potentials (no hysteresis). This titration procedure substantially reduced the consumption of anatase as compared with using fresh titania for each ionic strength. For NaNO3, some titrations have been repeated starting at high ionic strengths instead of gradually increasing the ionic strength, and practically the same ζ potentials were obtained, so it seems that the sample history does not influence the results. Some data points, especially those close to the iep, are missing due to instabilities in the Acoustosizer software: at random times, the software would fall into infinite loops or produce various error messages during either the data collection or reanalysis (background subtraction) portions of the experiment. In the vicinity of the iep, these types of software errors were commonplace and reproducible. However, even when the experiment and correction were executed without disturbances, the reliability of the magnitude and sign of very low ζ potentials is limited, because they are obtained from a difference of two large and approximately equal quantities: the overall signal and the background. The instrument also provides particle size data assuming a unimodal, lognormal distribution. This is calculated using viscosity and density values for pure water, so the validity of those data is limited. Far from the iep and at low ionic strengths, the apparent particle size of anatase was close to that from TEM, and it
Kosmulski and Rosenholm increased slightly with the ionic strength. Usually, sharp maxima of the particle size marked the iep. Results A background correction is crucial to obtain realistic values of ζ potentials from electroacoustic data at high ionic strengths since the signal from the electrolyte is likely to be comparable with or even higher than that from the particles. One might expect that a small relative difference in electrolyte concentration between the solution used to measure the background and the suspension would lead to a large absolute error in the calculated ζ potentials. To our surprise, in the systems of our interest, the background correction can be handled relatively easily even at very high ionic strengths because the exact match between the concentration of the electrolyte solution used to measure the background and that in the suspension is not essential, and deviations of a few percent are acceptable. This assertion is based on the following result. Solutions of an electrolyte of concentrations 0.9x (natural pH), x (pH ) 9), x (pH ) 4), and x (natural pH) are prepared (x is on the order of 1 mol dm-3), and the background signal from each of them is measured. Then, anatase is dispersed in the last solution and the suspension is titrated according to the procedure described in the previous subsection. The raw electroacoustic signal is corrected using different background data. Despite a 10% difference in the ionic strength between the first and the last background data, the resulting corrected ζ potentials differ only by a fraction of 1 mV. Also the effect of substantial pH difference between the samples used to measure background on the corrected ζ potentials is negligible. The small difference in electric conductance between suspension and the electrolyte used to prepare it and direct measurements of adsorption of ions (radiotracer method, ASA) clearly show that the adsorption on anatase does not significantly change the concentration of ions of the supporting electrolyte in the solution at the ionic strengths and solid to liquid ratios of interest, especially in the vicinity of the point of zero charge (pzc). Thus, one can safely assume that the effect of the difference in the electrolyte concentration caused by adsorption of ions on anatase on the magnitude of the background signal is negligible. Therefore, an electrolyte of the same concentration as the electrolyte that was used to prepare the suspension may be used to measure the background. The isoelectric point at low ionic strengths was found at pH0 ) 5.85 for anatase and at pH ) 7.6 for zirconia. The iep for anatase coincides with the point of zero charge found for a material prepared according to the same recipe, but it is at a considerably lower pH as compared with the iep reported in the same paper.3 The iep of anatase is insensitive to temperature over the range 18-35 °C and only slightly (by 0.15 pH unit) decreases when the temperature is raised to 45 °C. This result may be interpreted as nearly zero enthalpy of proton adsorption on anatase over the temperature range 18-35 °C. The calculation of enthalpy of proton adsorption on oxides from the temperature dependence of their pH0 is discussed in detail in ref 5. The iep of zirconia found in the current work is 1 pH unit above most pzc and iep values reported in the literature.6,7 However, recent measurements with carefully washed zirconia8 gave the pzc of zirconia very close to the iep found in the present study. A typical set of ζ(pH) curves for anatase at different ionic strengths is presented in Figure 1. At very low ionic strengths, the iep of anatase is independent of the ionic strength. This behavior is usually observed for the oxides.1 The low ionic
Adsorption of Ions on Anatase and Zirconia
J. Phys. Chem., Vol. 100, No. 28, 1996 11683 TABLE 1: Values of Ccrit2 (mol dm-3) of Anatase for Different 1-1 Salts (Values for Zirconia in Parentheses, NCC ) no Ccrit2, LS ) Low Solubility, Empty Cells ) Not Determined) Li CH3COO Cl NO3 ClO4 Br I a
1.02a
0.67
Na
K
Rb
NCCa >1.5 0.94 (0.76) 0.73 0.62 (0.48) 0.53
>1.5 >1.7 LS >1.5 NCC
>1.4 LS
Cs
NCC LS NCC
Two iep values (see text).
TABLE 2: The Most Negative Values of ζmin (mV) of Anatase in 0.4 mol dm-3 Solutions of Different 1-1 Salts (LS ) Low Solubility, Empty Cells ) Not Determined) Li CH3COO Cl NO3 ClO4 Br I a
Figure 1. ζ potential of anatase at different pH values and different concentrations of NaBr (a, top) and KNO3 (b, bottom).
strength ζ potentials are not shown in Figure 1; the iep is indicated by a vertical line (pH0). When the ionic strength exceeds a critical value (Ccrit1), which is below 0.1 mol dm-3 for common 1-1 salts, the iep starts to shift gradually toward higher pH values with increasing ionic strength. Finally, when the ionic strength exceeds a second critical value (Ccrit2), there is no iep and the ζ potentials are positive over the entire measured pH range. On the other hand, the acidic branch of the ζ(pH) curves shows a “normal” behavior even at electrolyte concentrations above Ccrit2, namely, that the positive ζ potentials at a given pH gradually decrease when the ionic strength increases. The absolute values of negative ζ potentials (basic branch) decrease much faster with increasing ionic strength. The curves shown in Figure 1 have a common intersection point corresponding to pH ) 5.5 and ζ ) 11 mV. The analogous curves for many other salts studied in this paper also show a common intersection point corresponding to ζ potentials between 4 and 10 mV. Sodium bromide gave a result unique among the salts studied in this paper: the apparent ζ potentials (before background correction) give an iep at pH ) 6, which was independent of the ionic strength. On the other hand, the potassium, rubidium, and cesium salts as well as NaCl always show an iep that is shifted toward higher pH values as compared
-11
-3b b
Na
K
Rb
-10 a -11 -8 -6 -2
-11 -12 LS -6 -7
-11
Error messages. 0.5 mol
LS
Cs
-16 LS -15
dm-3.
with the low ionic strength value. This type of behavior is illustrated in Figure 1b for KNO3. The values of Ccrit2 for several 1-1 electrolytes are summarized in Table 1. The difference between different sodium salts considerably levels out when critical activities are calculated as products of Ccrit2 from Table 2 and mean activity coefficients of the particular salts.9 Low activity coefficients of potassium and cesium salts for which no Ccrit2 was found are in line with this observation. Since a scatter of results and instrument failures for |ζ| < 1 mV make it difficult to measure ζ potentials near the iep, the value of Ccrit2 cannot be determined directly. An interpolation method has been applied to estimate Ccrit2. At high ionic strengths and high pH values, the ζ potential of titania is rather insensitive to the pH (Figure 1) and a plateau is established. The plateau values ζmin are plotted against the electrolyte concentration (Figure 2), and Ccrit2 is interpolated as the electrolyte concentration corresponding to ζmin ) 0. An unexpected effect is observed with sodium acetate at concentrations higher than 0.6 mol dm-3: at pH > 7, the ζ potential of anatase constantly increases (becomes less negative), and at 1.2 mol dm-3 it becomes positive at a pH as high as 10.2. The course of ζ(pH) curves with the ζ potential decreasing with the pH, reaching a minimum and then increasing, resembles that which is observed for rutile in solutions of alkaline earth metal nitrates.2 Using the terminology introduced in the present paper, the Ccrit2 values for alkaline earth metal nitrates found in ref 2 or, generally, for salts of multivalent metals are well below 10-3 mol dm-3 (with the exception of magnesium), so they are lower than those found in the present paper for alkali metal cations by 3 orders of magnitude. Many further analogies between specific adsorption of cations and the results obtained in the present study may be found; for example, the acidic branch of the ζ(pH) curves is rather insensitive to the nature of the cation in both cases. Various parameters may be used to characterize the effects of different salts on the basic branch of the ζ(pH) curves. Ccrit2 (Table 1) is one of these parameters, and it is appropriate to characterize the effects at very high pH values and very high salt concentrations. For lithium and sodium salts a value for Ccrit2 is actually observed. This is not the case for the cesium, rubidium, and potassium salts. There is clearly a large gap
11684 J. Phys. Chem., Vol. 100, No. 28, 1996
Kosmulski and Rosenholm
Figure 2. Minimum values of ζ potential of anatase observed at different concentrations of sodium salts (a, top left), iodides (b, top right), chlorides (c, bottom left) and potassium salts (d, bottom right).
between lithium and sodium on one hand and potassium, rubidium, and cesium on the other in terms of electrokinetic properties of anatase in concentrated solutions of their salts. The Ccrit2 values for zirconia reported in Table 1 are very close to those found for anatase. This suggests that the Ccrit2 values are rather insensitive to the nature of the oxide, and they depend mainly on mutual interactions of water and ions in the bulk solution. The effects of different salts at very high pH values can also be compared in terms of the values of ζmin at a given concentration. In contrast to Ccrit2, this parameter can also be used to characterize the salt effects at moderate concentrations. At 0.1 and 0.2 mol dm-3 almost the same ζmin values have been obtained for most salts with exception of those which show very low Ccrit2 values. At higher concentrations, substantial differences between particular salts appear. The ζmin values of anatase in 0.4 mol dm-3 solutions of different 1-1 salts are summarized in Table 2. For all salts, the absolute values of ζ potentials reported in Table 2 are considerably lower than those observed at pH ) 4 (ζ ≈ 25 mV for most salts) at the same concentration. The data in Table 2 show the same trend as in Table 1: the extent of the cation effect (lower Ccrit2, lower ζmin at a given ionic strength) increases in the series Cs < K < Na < Li for a given anion and CH3COO < Cl < NO3 < ClO4 < Br < I for
a given cation. There is one exception from this rule: the effect of NaI is more pronounced than that of LiI. These series are well correlated with direct measurements of adsorption of particular ions on anatase:10 lithium adsorbs much better than cesium, and chloride adsorbs better than iodide. Thus, the combination of weakly adsorbing anions and strongly adsorbing cations leads to an excess of cations between the surface and the slipping plane and thus to more positive ζ potentials. The anion series observed in this study is also identical with the Hofmeister lyotropic anion series given in ref 11: low Ccrit2 of sodium salt corresponds to low flocculating power (and thus low adsorption) of anions. However, there are many different anion series (for different versions of Hofmeister series see references in ref 12), and most of them are only partially correlated with our results. Certainly, the correlation with ion adsorption or Hofmeister series, which are only experimental results, does not explain the molecular mechanism of the phenomena observed in the present study. Discussion Most measurable properties of the alkali metal cations, e.g., the ionic radii,13 change monotonically from Li to Cs. Therefore, correlation of the results obtained in the present study with
Adsorption of Ions on Anatase and Zirconia
J. Phys. Chem., Vol. 100, No. 28, 1996 11685
TABLE 3: Structure Breaking and Making Scale15 cation
∆GHB
anion
∆GHB
Li+ Na+ K+ Rb+ Cs+
0.28 -0.03 -0.52 -0.56 -0.69
CH3COOClNO3BrClO4I-
0.12 -0.61 -0.68 -0.80 -1.01 -1.09
different anion rather than cation series is more likely to provide an answer as to which property of the ions is responsible for the difference in electrokinetic potentials of titania at high ionic strengths for the different salts. For example, there is no correlation between ionic radii of anions and Ccrit2 for the corresponding sodium salts. The strong effect of combinations of hard cations (Li, Na) and soft anions (I) as compared to other salts suggests that the HSAB rule14 may be applied here. Unfortunately, most available hardness-softness scales14 cover only some of the anions studied in the present paper. Properties of 1-1 electrolytes in aqueous solutions are influenced by water structure and Vice Versa. At low ionic strengths these effects may be neglected, but this is no longer true at high ionic strengths. In terms of many properties (viscosity, dielectric constant) water at high and low ionic strengths behaves as “different solvents”. The structure of water may be defined by an average number of hydrogen bonds in which a water molecule participates (1.55 in pure water at 25 °C), and the solute effect on water structure is defined as a change of this number, ∆GHB.15 Structure-making ions have positive ∆GHB values and structure-breaking ions have negative ∆GHB values, respectively. The ∆GHB values for different ions have been recently summarized, and their correlation with other parameters defining the effects of ions on water structure has been discussed.15 Selected ∆GHB values shown in Table 3 explain not only the series of cations and anions found in this paper but also the large gap between sodium and potassium. The salts for which ∆∆GHB )∆GHB(cation) - ∆GHB(anion) > 0.6 show Ccrit2, while the other salts do not. This is strictly an empirical rule without any theoretical justification, but it works well for all the salts studied in this paper. It may serve as a semiquantitative prediction of the numerical value of Ccrit2: high values of ∆∆GHB (>1) correspond to low Ccrit2 values. In addition, specific adsorption of cations (very low Ccrit2) can be treated in terms of ∆∆GHB: most multivalent cations are structure making, and ∆∆GHB > 1 for their halides and nitrates. Thus, combinations of structure-making cations and structurebreaking anions may lead to positive ζ potentials of anatase despite high negative surface charge3 over the entire pH range. This means that at high ionic strengths, structure-making cations accumulate in the interfacial region even against the electrostatic force, while an excess of structure-breaking anions remain in the bulk solution. A more careful analysis of the effects of various salts shows that despite clear anion and cation series discussed above, the effect of a given salt cannot be considered as a simple combination of the anion and cation contributions. Small cations show a differentiating effect: the course of ζ(pH) curves of particular lithium and sodium (Figure 2a) salts at a given high ionic strength is very sensitive to the nature of the anion. In contrast, the effect of the nature of the anion is relatively insignificant when different potassium salts are considered (Figure 2d). On the other hand, large anions (I) show a differentiating effect (Figure 2b), while smaller anions (Cl) do not (Figure 2c). The differentiating effect observed in this study is not unique. Similar differentiation effects are observed when the activity
coefficients of different 1-1 salts at a given concentration are compared. The mean activity coefficients in 1 mol dm-3 solutions of LiI (0.910), CsI (0.533), LiCl (0.774), and CsCl (0.544)9 clearly illustrate a differentiating effect of lithium (large difference between iodide and chloride) and lack of such an effect for cesium (negligible difference between iodide and chloride). The same examples may be used to show the strong differentiating effect of iodide and much weaker differentiating effect of chloride. The position of the iep at a given salt concentration can be used to characterize the effects of different salts on the basic branch of the ζ(pH) curves in the region of moderate pH values (6-9) at different concentrations below Ccrit2 and also for these salts that do not show Ccrit2. This parameter can only be roughly estimated: near the iep, the measured ζ potentials show a significant scatter and the instrument often failed to give any results at all. Typical dependence of the iep position on the concentration of different salts is presented in Figure 3. The comparison of particular salts at concentrations above 0.7 mol dm-3 shows the same trends as those found for Ccrit2 and ζmin values. The same anion and cation scales are obtained, and the same differentiating effect of sodium and lithium on one hand and of iodide on the other is observed. However, at lower ionic strengths, 0.1 < I < 0.7 mol dm-3, only the anion scale found for Ccrit2 is valid, while the cation scale is not. In terms of the iep shift, the effect of potassium salts (KCl, KI) is more pronounced than that of corresponding lithium salts (LiCl, LiI) at electrolyte concentrations below 0.6 mol dm-3 (Figure 3b,c). This result is in line with the Hofmeister series of cations. While the anion series from different sources are not very consistent, the same cation series: Ba2+ > Sr2+ > Ca2+ > Cs+ > K+ > Na+ > Li+ (decreasing flocculating power), can be found in most handbooks. More pronounced shifts of the iep correspond to higher flocculating power (and high adsorption) of cations when the ionic strength is not too high. Similar correlation is observed with specific adsorption of bivalent cations. There is no iep in the electrokinetic curves of titania in 10-3 mol dm-3 solutions of alkaline earth nitrates except for magnesium,2 but the ζmin characterizing the cation effect on the basic branch of these curves clearly increases from Mg to Ba, according to the Hofmeister series. At higher electrolyte concentrations (about 1 mol dm-3), sodium and lithium salts reach their Ccrit2 (no i.e.p), while for potassium and cesium salts (with exception of KNO3) the trend in iep shift is reversed; that is, it shifts back toward lower pH values. The reversal in the direction of the iep shift is accompanied by reversal or at least stabilization of the trend shown in Figure 2, namely, of the increase of ζmin with the ionic strength. This concentration dependence of the cation series clearly indicates the difference in adsorption properties between hydrated (low ionic strength) and partially dehydrated (high ionic strength) lithium and sodium ions: at lower ionic strengths, the effect of hydrated sodium and lithium cations on the iep of anatase is less pronounced as compared with potassium, but at higher ionic strengths, partially dehydrated sodium and lithium cations accumulate at the surface even against the electrostatic force. The Hofmeister series is based on the comparison of cations at ionic strengths well below 0.5 mol dm-3. It follows from the present results that at concentrations between 0.5 and 1 mol dm-3, depending on the nature of the anion, the adsorbability of sodium and lithium ions becomes higher than that of potassium, rubidium, and cesium. The “normal” order that is observed at lower ionic strengths and that corresponds to the Hofmeister series is reversed when the molar water to sodium
11686 J. Phys. Chem., Vol. 100, No. 28, 1996
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Figure 3. iep of anatase observed at different concentrations of sodium salts (a, top left), iodides (b, top right), chlorides (c, bottom left), and potassium salts (d, bottom right). The data points corresponding to the concentrations above Ccrit2 are plotted as pHiep ) 10.
(or lithium) ratio is greater than 50:1. This suggests that sodium and lithium ions adsorbed on anatase from low ionic strength media are surrounded by more than one layer of water molecules. This thick hydration sheath is responsible for the relatively low adsorbability of sodium and lithium. However, when the water activity is reduced, partially dehydrated sodium and lithium ions can adsorb specifically, i.e., against the electrostatic repulsion. The difference in adsorption behavior between sodium and potassium plays an important role in many physiological, geochemical, and industrial processes. This study shows that the trends observed at low ionic strengths are not necessarily valid at higher ionic strengths, and especially the concept of indifferent electrolytes is strictly limited to ionic strengths below 0.1 mol dm-3. The assertion16 that the iep of titania remains unchanged at NaCl concentration as high as 0.2 mol dm-3 may serve as an example of uncritical extrapolation of low ionic strength behavior to significantly higher ionic strengths, well beyond the range of this rule’s applicability. First, it is very surprising that any electroctrophoretic mobilities of titania could be measured at such a high value of the ionic strength using a traditional ζ-meter. Taking into account the large particle size, high specific density, and high Hamaker constant17 of titania, particles used in ref 16 should sediment out within a few seconds
after removing the dispersion from the ultrasonic bath. This settling out is expected over a wide range of pH values. The stability of the dispersion used to measure the ζ potential of titania in ref 16 may be due to the presence of surface active agents in the titania, which were not completely removed during the washing procedure. Small amounts of such substances might also influence the values of the ζ potential and the position of the iep. On the other hand, the turbidity data as a function of pH at various ionic strengths reported in the same paper clearly show that the minimum of turbidity (and the real iep) shifts toward higher pH values as the ionic strength increases: they confirm the trends observed in this study. It is also noteworthy that the range of low turbidity reported in ref 16 is symmetrical with respect to the iep at low ionic strengths, but in 0.2 mol dm-3 NaCl the region of low turbidity on the basic side is much wider. This indicates that the absolute values of ζ potentials on the basic side are lower than those on the acidic side, according to the trend found in the present study. Also the increase of turbidity at pH ) 6.2 at very high ionic strengths found in ref 16 can be easily explained in terms of the present results. For low ionic strengths, the pH ) 6.2 is the iep, but the same pH value is not the iep at the NaCl concentrations above 0.1 mol dm-3. Therefore, the absolute value
Adsorption of Ions on Anatase and Zirconia of the ζ potential of titania at this pH value increases when the ionic strength increases and the system becomes more stable. In view of the above discussion, there is no reason to consider any “repulsive hydration force” introduced in ref 16 to explain the stability of titania at high ionic strengths at pH ) 6.2, and it is clear why this apparent force was not observed in direct force measurements.17 The shift of the iep toward higher pH values at concentrations of “indifferent” electrolytes above 0.1 mol dm-3 is probably a common phenomenon for the oxides. It has been observed for silica in 0.1 mol dm-3 RbCl and CsCl.18 Unfortunately, reliable electrophoretic data for silica at higher ionic strengths are not available. In mixed aqueous-organic solvents,19 Ccrit1 and Ccrit2 are much lower than those found in the present study for aqueous systems. The probable explanation is that the molecules of organic cosolvents (e.g., methanol) compete with the ions for water. Therefore, a greater degree of water deficit and incomplete hydration is observed at lower salt concentrations than in aqueous systems. Note Added in Proof: One of the referees has done some measurements on another oxide at high electrolyte concentration taking into account the conductivity effects and has found that the results presented in this paper are qualitatively consistent. References and Notes (1) Hunter, R. J. Zeta Potential in Colloid Science. Academic Press: New York, 1981.
J. Phys. Chem., Vol. 100, No. 28, 1996 11687 (2) Jang, H. M.; Fuerstenau, D. W. Colloids Surf. 1986, 21, 235. (3) Kosmulski, M.; Matijevic, E. Colloids Surf. 1992, 64, 57. (4) Parfitt, G. D. Prog. Surf. Membr. Sci. 1976, 11, 181. (5) Kosmulski, M.; Matysiak, J.; Szczypa, J. J. Colloid Interface Sci. 1994, 164, 280. (6) Blesa, M. A.; Maroto, A. J.; Regazzoni, A. E. J. Colloid Interface Sci. 1984, 99, 32. (7) Mandel, F. S.; Spencer, H. G. J. Colloid Interface Sci. 1980, 77, 577. (8) Janusz, W. Unpublished data. (9) CRC Handbook of Chemistry and Physics, 60th ed.; CRC: Boca Raton, 1979; p D-169. (10) Sprycha, R. J. Colloid Interface Sci. 1984, 102, 173. (11) Stauff, J. Kolloidchemie; Springer: Berlin, 1960; p 344. (12) Collins, K. M.; Washabaugh, M. W. Quatum ReV. Biophys. 1985, 18, 323. (13) Marcus, Y. Chem. ReV. 1988, 88, 1475. (14) Pearson, R. G. J. Am. Chem. Soc. 1963, 85, 3533; 1986, 108, 6109; 1988, 110, 7684. (15) Marcus, Y. J. Solution Chem. 1994, 23, 831. (16) Yotsumoto, H.; Yoon, R.-H. J. Colloid Interface Sci. 1993, 157, 426. (17) Larson, I.; Drummond, C. J.; Chan, D. Y.; Grieser, F. J. Am. Chem. Soc. 1993, 115, 11885. (18) Kosmulski, M.; Matijevic, E. Colloid Polym. Sci. 1992, 270, 1046. (19) Kosmulski, M. Colloids Surf. 1995, 95, 81.
JP960315U
