Home
Search
Collections
Journals
About
Contact us
My IOPscience
The influence of concentration on specific ion effects at the silica/water interface
This content has been downloaded from IOPscience. Please scroll down to see the full text. 2014 J. Phys.: Condens. Matter 26 244107 (http://iopscience.iop.org/0953-8984/26/24/244107) View the table of contents for this issue, or go to the journal homepage for more Download details: IP Address: 132.239.1.231 This content was downloaded on 10/06/2017 at 07:06 Please note that terms and conditions apply.
You may also be interested in: The electric double layer at a rutile TiO2 water interface modelled using density functional theory based molecular dynamics simulation J Cheng and M Sprik Electrostatic interaction forces in aqueous salt solutions of variable concentration andvalency Daniel Ebeling, Dirk van den Ende and Frieder Mugele Density functional theory simulation of hydrogen-bonding structure and vibrational densities of states at the quartz (101)-water interface and its relation to dissolution as a function of solution pH and ionic strength M J DelloStritto, J Kubicki and J O Sofo Ions at hydrophobic interfaces Yan Levin and Alexandre P dos Santos Molecular dynamics studies of aqueous silica nanoparticle dispersions: salt effects on the double layer formation Lucas S de Lara, Vagner A Rigo, Mateus F Michelon et al. Temperature responsive properties of poly(N-vinylcaprolactam) multilayer hydrogels in the presence of Hofmeister anions Oleksandra Zavgorodnya, Veronika Kozlovskaya, Xing Liang et al. Oxo-acidity in ionic melts Viktor L Cherginets Ion specific correlations in bulk and at biointerfaces I Kalcher, D Horinek, R R Netz et al.
Journal of Physics: Condensed Matter J. Phys.: Condens. Matter 26 (2014) 244107 (10pp)
doi:10.1088/0953-8984/26/24/244107
The influence of concentration on specific ion effects at the silica/water interface Md Shafiul Azam, Akemi Darlington and Julianne M Gibbs-Davis Department of Chemistry, University of Alberta, Edmonton, Alberta T6G 2G2, Canada E-mail: [email protected] Received 13 November 2013, revised 2 February 2014 Accepted for publication 4 February 2014 Published 27 May 2014 Abstract
Second harmonic generation spectroscopy is a useful tool for monitoring changes in interfacial potential at buried insulator/liquid interfaces. Here we apply this technique to the silica/ aqueous interface and monitor the changes in interfacial potential while varying the pH in the presence of different alkali halides at 0.1M concentration. Within the pH range explored, the bimodal distribution of acidic sites on planar silica is clearly observed, corresponding to two types of acidic SiOH groups. Comparing these data with previous work at 0.5M sheds light on whether the presence of the ions stabilizes the charged or neutral state of the surface sites. For the alkali chlorides, with the exception of NaCl, we observe that the presence of the alkali chlorides stabilize the less acidic site in the protonated (SiOH) rather than deprotonated (SiO−) form. This unusual influence of the cation is attributed to the combination of interactions at the interface between water, surface sites and the electrolyte. Overall, we observe that the influence of the alkali ion on the ratio of the two types of sites and their effective acid dissociation constants is minor at 0.1M, unlike that observed at 0.5M. In contrast, the influence of the anion on the cooperative dissociation of surface sites and their relative distribution is little affected upon decreasing the concentration, which indicates that these specific anion effects are prevalent in nature. Keywords: specific ion effects, buried interfaces, nonlinear optical spectroscopy (Some figures may appear in colour only in the online journal)
1. Introduction
represents environmentally relevant conditions using colloidal morphologies. One strategy to avoid the instability of colloids is to monitor planar silica surfaces. Such macroscopic substrates are amenable to different spectroscopic methods such as evanescent wave cavity ring down spectroscopy (EW-CRDS [10, 11]) and nonlinear optical methods like second harmonic generation (SHG) spectroscopy [12, 13]. In one approach, pH-sensitive dyes have been used in conjunction with EW-CRDS to probe the change in resonant frequency of the dye upon interacting with the surface [10, 11]. However, many have questioned whether the dye’s acid dissociation equilibrium constant Ka is different at the interface, knowing which is critical to accurately quantify the surface acidity using this approach [14, 15]. Another technique that represents a label-free method for monitoring the acid–base chemistry of the silica interface is based on the dependence of SHG on the interfacial potential. The first report by Eisenthal and co-workers utilizing
The silica/water interface is abundant in the natural environment, and consequently the properties of silica have been the subject of research for the last several decades [1–3]. In particular, a variety of techniques has been applied to the study of deprotonation and dissolution processes at the silica/water and functionalized silica/water interface [4–6]. Both bulk [7] and surface specific [8, 9] techniques have been employed, providing complementary pieces of information, which aid in forming a complete picture of these interfacial processes. Many of these studies have focused on colloidal silica and have included the effect of the aqueous electrolyte on surface deprotonation and dissolution events [1]. One challenge, however, in utilizing silica colloid dispersions is that they are not stable over a large pH range [1]. Consequently, it is difficult to map out the entire acid–base equilibria of silica over the broad pH range that 0953-8984/14/244107+10$33.00
1
© 2014 IOP Publishing Ltd Printed in the UK
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
this non-resonant SHG method to monitor the silica/aqueous interface found that planar silica exhibited two distinct deprotonation processes between pH 2 and 13 [12, 13]. Both of these acidic sites have been attributed to SiOH groups that are in two different hydrogen bonding environments resulting in the drastic difference in surface acidities. The exact nature of these sites has been the subject of computational investigations and varying origins for the surface acidity have been proposed [16, 17]. Recently we have utilized the same non-resonant SHG method referred to as the χ(3) technique to probe the influence of different alkali halides on the acid–base properties of the silica/aqueous interface over a broad pH range [18, 19]. We observed that the acidity of the silica surface depended greatly on the composition of the aqueous electrolyte. Indeed, in the presence of various alkali chlorides, the effective acid equilibrium constants (Ka) ranged from 10−8 to 10−11 for the less acidic silanol groups and decreased with increasing alkali ion polarizability [18]. The trend in Ka and cation polarizability was also observed for the more acidic sites deprotonated below pH 7, but the range of constants was smaller. Additionally, the relative distribution of the two types of silanol sites also varied widely. Specifically our work confirmed the previous results of Eisenthal, conducted with 0.5M NaCl, which showed that 20% of the silica sites were the more acidic sites that were deprotonated below pH 7. However, substituting NaCl with LiCl led to a shift in the distribution of these sites such that 60% of the surface was now composed of the more acidic sites, which were deprotonated below pH 7. Our observations suggested that the surface charge density of silica could vary dramatically at the same pH, depending on the composition of the aqueous phase, consistent with previous reports examining the acid–base behavior of the silica/water interface [20–23]. More recently, we examined the role of the anion on deprotonation processes at this interface [19]. Despite the fact that the silica/water interface is negatively charged above pH 2, the anion has an even larger influence on the apparent Ka of the surface sites as well as their relative distribution. For example, utilizing NaI in place of NaCl led to very sharp deprotonation curves in the surface titration experiments, which are indicative of positive cooperativity whereby the dissociation of one silanol makes subsequent deprotonation events more favorable for a certain number of neighboring sites. Additionally, the ratio of more acidic to less acidic silanol sites increased from 1 : 4 with 0.5M NaCl to ~4 : 1 with 0.5M NaI. For the latter, this ratio corresponded to 85% of the surface becoming deprotonated at neutral pH, very different from the 20% deprotonated sites in the presence of NaCl. Such specific anion effects had not been explored at a negative interface, and this was the first demonstration of cooperativity in these systems. Here we gain further insight into the origins and mechanism of these specific ion effects at the silica/water interface by performing these SHG experiments at a lower salt concentration (0.1M). By reducing the salt concentration and monitoring whether the surface becomes more or less acidic, we can probe the role of the ions in inhibiting or promoting surface deprotonation. Finally, monitoring the interface at this lower salt concentration sheds light on how robust these
specific ion interactions are at more relevant environmental and biological concentrations. 2. Methods: SHG experiments A detailed description of the laser system used in these experiments can be found in our previously published work [18, 19, 24]. Briefly, using the output of a regeneratively amplified laser system (Spitfire Pro, Spectra Physics, 1 kHz, 100 fs, 3.3 W) to pump an optical parametric amplifier (Spectra Physics OPA-800CF) the incident light was tuned to 550 ± 2 nm and directed through a half-wave plate and polarizer for s-polarization selection. This wavelength was selected as it leads to a second harmonic wavelength that is off resonance with respect to the electrolyte and is the same wavelength that was used for our experiments at 0.5M, which represents an important comparison for this work [18, 19]. The polarized light was focused onto the fused silica/water interface at an angle of 62° from surface normal near total internal reflection. The reflected second harmonic light generated at the interface was then passed through a colored glass filter (Thorlabs) to remove the reflected fundamental light and focused onto a monochromator (Optometrics Corp., MiniChrom MC1-02) tuned to the second harmonic wavelength (275 nm). SHG was detected using gated photon counting with a PMT assembly. Before performing each experiment, the quadratic power dependence and SHG wavelength dependence were verified. For all experiments, a freshly cleaned fused silica hemisphere (ISP optics, 1 inch diameter, QU-HS-25, UV-grade SiO2) was placed on a custom-built Teflon cell so that the flat surface of the hemisphere was in contact with the aqueous phase [18]. The silica hemisphere was cleaned prior to use by sonicating it in Millipore water, then methanol and water. Next, Nochromix (Godax Laboratories, 5% w/v solution in H2SO4) was placed on the flat side for 1 h followed by copious rinsing in Millipore water and further sonication in water, methanol then water. The hemisphere was allowed to dry in an oven at 100 °C for 10 min, then cooled down to room temperature and plasma cleaned (Plasma cleaner, PDC-32G, Harrick Plasma) in air for 2–3 min. Each sample was exposed to Milli-Q water until the SHG signal was optimized before adding the electrolyte solution. Separate experiments were conducted at low pH and high pH on different fresh samples to avoid hysteresis [25]. More details can be found in our previous publications [18, 19]. 3. Theoretical background: the χ (3) technique Second harmonic light is generated when light of frequency ω interacts with aligned dipoles at an interface causing induced polarization that oscillates at twice the incident frequency 2ω (P2ω) resulting in a new electric field at 2ω (E2ω). The intensity of the SHG response from the interface (I2ω) is proportional to the induced second-order polarization as follows [8, 26]: (1) I2ω = E2ω ∝ P2ω = χ (2) EωEω . 2
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
In this expression, Eω is the incident electric field and χ is the nonlinear susceptibility of the interface. When surfaces are charged the presence of a static electric field (E0) leads to an additional third-order contribution to the second harmonic field [12]: (2)
∫
E2ω ∝ χ (2) EωEω + χ (3) EωEω E0 dx = χ (2) EωEω + χ (3) EωEωΦ0 (2) where χ(3) is the third-order susceptibility of the interface. Integrating this static electric field that originates from the surface charges out into the bulk yields the interfacial potential Φ0. For the silica/aqueous interface, the surface is negatively charged above pH 2, which means that it is charged under most environmental conditions. Consequently, this χ(3) technique is well suited to probe processes at the silica surface that are dictated by electrostatic interactions, which modulate the interfacial potential [27]. For monitoring deprotonation processes, we wish to relate the change in E2ω directly to changes in the surface charge density of silica. Therefore, it is useful to perform the experiments at high salt concentrations (≥100 mM). At these high salt concentrations, we assume that the changes to χ(2) and χ(3) are negligible upon varying pH compared with the changes to Φ0 [12, 13, 18, 19, 25, 27, 29]. This assumption is based in part on the dependence of these susceptibilities on the amount of ordered water [30], which should vary minimally at high salt concentration as the Debye length is short and the static electric field decays rapidly into the bulk [31, 32]. In other words, the water molecules that experience the static electric field at the interface are primarily in the compact Stern layer of cations at these high salt concentrations. We reason that the net orientation and number of these water molecules may change slightly upon varying the pH, but not as significantly as the change in the static electric field or the corresponding interfacial potential upon changing the pH and deprotonating surface sites [19]. Furthermore, at such high salt concentrations the Stern model of the interface can be approximated by the constant capacitance model as the diffuse layer of ions has collapsed (figure 1) [28]. As a result, the surface charge density becomes directly proportional to the interfacial potential. From these assumptions, we can write the general equation:
Figure 1. SHG at the silica/electrolyte interface. The second harmonic field is proportional to the interfacial potential (Φ0) set up by charged surface sites. This charged interface can be represented by the basic Stern model, where the potential determining surface sites make up the 0-plane. According to this model a layer of cations forms at the interface, shown at the d-plane, which make up the Stern layer. Any remaining negative charges that are not neutralized by the Stern layer are neutralized by cations in the diffuse layer (not shown). At high salt concentrations (≥ 0.1M) the Stern model reduces to the constant capacitance model, which assumes that the surface is completely neutralized by the cations in the Stern layer leading to the disappearance of the diffuse layer and Φd = 0 [28].
and that of the less acidic sites rσLA is E2ω − E2ω (pH 7) , rσLA = (6) E2ω (pH 12) − E2ω (pH 7)
based on the observation that the more acidic sites are deprotonated between pH 2 and 7, while the less acidic sites are deprotonated between pH 7 and 12. We note that the validity of the constant capacitance model at 0.1M in the presence of NaCl is doubtful, as experimental data suggest that the diffuse layer is still present for this electrolyte at this concentration [28, 33]. However, for simplicity we utilize this model for all of the alkali chlorides investigated. 4. Results and discussion 4.1. Influence of the alkali chloride concentration
⎛σ⎞ (3) E2ω ∝ A + BΦ0 = A + B ⎜ ⎟ ⎝C ⎠
In our first report on specific ion effects we observed that the bulk pH value that corresponded to deprotonation of half of a given type of site, pH0.5 bulk, followed the trend NaCl < LiCl < KCl < CsCl for both the more and less acidic sites [18]. These data indicated that the presence of NaCl made both surface sites more acidic such that they were deprotonated at a lower pH, while CsCl made both surface sites less acidic. We proposed a simple model based on the affinity of the alkali ion for the siloxide sites generated after deprotonation, which is similar to commonly used surface complexation models for quantifying the effect of the electrolyte on surface acid–base equilibria [34]. In this proposed mechanism the acid dissociation of the surface silanol groups first occurred, leading to the siloxide conjugate base:
where A is equal to χ (2) EωEω and B is equal to χ (3) EωEω, which are constant, C is the capacitance of the Stern layer and σ is the surface charge density. At low pH near pH 2 the surface becomes net neutral and A = E2ω (pH 2). (4)
Therefore the relative surface charge density can be found by normalizing E2ω from zero to one. Specifically, the relative surface charge density stemming from the more acidic sites rσMA is E2ω − E2ω (pH 2) rσMA = (5) E2ω (pH 7) − E2ω (pH 2)
K
a (7) SiOH + H2 O ⇄ SiO− + H3O+ .
3
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107 +
The cation, M , then coordinated to the siloxide through electrostatic interactions. KSC (8) SiO− + M+ ⇄ SiO− ⋅ M+ .
Although not explicitly shown, this electrostatic interaction was presumed to be mediated through water, resulting in a screening interaction rather than neutralization of the negative charges [21]. Finally, the overall equilibrium expression with the effective acid dissociation constant (Kaeff ) was equal to the sum of the two separate equilibrium expressions: eff
Ka (9) SiOH + H2O + M+ ⇄ SiO− ⋅ M+ + H3O+ .
A useful parameter for exploring the acidity of surfaces that is model independent is the acidity quotient Qaeff , which is related to the equilibrium constant through a potential dependent term that accounts for the surface concentration of all species [31]. The expression related to equation (9) that utilizes the bulk concentrations is [SiO− ⋅ M+ ][H3O+ ]bulk (10) Qaeff = QaQSC = [SiOH][ M+ ]bulk
where Qa, QSC and Qaeff are quotients that can be related to the actual equilibrium constants Ka, KSC and Kaeff , respectively, by converting the bulk cation and hydronium concentrations to the surface concentrations (see bolow). The value of pQaeff (–log Qaeff ) can then be determined from pH0.5 bulk and the + value: –log[M+]bulk or pMbulk + . (11) pQaeff = pH 0.5 bulk − pMbulk
According to the proposed coupled equilibria, decreasing the cation concentration should shift the equilibrium shown in equation (9) to the left, resulting in a less acidic surface and higher pH0.5 bulk values. To test whether the presence of the cation stabilized the siloxide sites, we decreased the concentration of the cation to 0.1M for NaCl and monitored the resulting acid–base behavior (figure 2). As shown in fi gure 2, decreasing the Na+ cation concentration indeed led to an increase in the pH0.5 bulk value, which is consistent with the + value at lower salt and the rearranged form of higher pMbulk equation (11):
Figure 2. Representative plots of normalized E2ω as a function of pH for the silica/electrolyte interface with aqueous NaCl of different concentrations at (a) low pH and (b) high pH. The solid lines (for 0.5M) and dashed lines (for 0.1M) represent the fit of a sigmoidal function to the data.
the less acidic surface sites were deprotonated, contrary to equation (12) (figures 3(d)–(f )). These results suggested that the presence of alkali chlorides other than sodium chloride actually stabilized the neutral (SiOH) rather than the negative form (SiO−) of these less acidic silanol groups, thereby requiring more basic environments to achieve deprotonation. To rationalize the observed concentration-dependent behavior for these salts, we neglect association of the cation to the siloxide and instead consider that the first step is cation dissociation from the neutral silanol,
+ . (12) pH 0.5 bulk = pQaeff + pMbulk
Using the pH0.5 bulk values of 9.6 ± 0.1 and 8.6 ± 0.1 and the corresponding p[Na+] values in equation (12) yielded pQaeff values of 8.6 ± 0.1 and 8.3 ± 0.1 for 0.1 and 0.5M NaCl concentrations, respectively (table 1). The decent agreement for the two pQaeff values suggests that our simple model whereby sodium stabilizes the siloxide provides a reasonable interpretation of its influence on the surface equilibria. Very similar salt-dependent behavior was also observed for the nanoporous silica/NaCl electrolyte interface in a recent SHG experiment by Borguet and co-workers, indicating this is general for different silica morphologies [21]. However, repeating the concentration variation measurement for the other cations led to very different results (figure 3). Specifically, decreasing the salt concentration of Li+, K+ and Cs+ led to a decrease in the pH at which half
K
d (13) SiOH ⋅ M+ ⇄ SiOH + M+,
which leads to the following overall equilibrium expression: KdKa
(14) SiOH ⋅ M+ + H2 O ⇄ SiO− + M+ + H3O+ . The salt-dependent equation that results is [SiO− ][ M+ ]bulk [H3O+ ]bulk Qaeff ′ = QdQa = . (15) [SiOH] 4
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Table 1. Thermodynamic parameters from the alkali chloride variation experiments . a
Salt
[Salt] (M)
pH0.5 bulk
pQaeff
pQaeff ′
pH0.5 surface
pKaeff
pKaeff ′
LiCl
0.1 0.5 0.1 0.5 0.1 0.5 0.1 0.5
8.9 ± 0.1 9.9 ± 0.2 9.6 ± 0.1 8.6 ± 0.1 8.9 ± 0.3 10.2 ± 0.2 9.5 ± 0.2 11.1 ± 0.1
— — 8.6 ± 0.1 8.3 ± 0.1 — — — —
9.9 ± 0.2 10.2 ± 0.2 — — 9.9 ± 0.3 10.5 ± 0.2 10.5 ± 0.2 11.4 ± 0.1
7.4 ± 0.1 8.0 ± 0.2 7.8 ± 0.1 7.3 ± 0.1 7.4 ± 0.3 8.5 ± 0.2 7.9 ± 0.2 9.2 ± 0.1
— — 6.8 ± 0.1 7.0 ± 0.1 — — — —
8.4 ± 0.2 8.3 ± 0.2 — — 8.4 ± 0.3 8.8 ± 0.2 8.9 ± 0.2 9.5 ± 0.1
NaCl KCl CsCl
a All the given values are the average of at least two independent measurements and the error represents the range of measured values.
At pH0.5 bulk when half of the surface sites are deprotonated (i.e. when [SiO−] = [SiOH])
cesium and water. The net neutral silica surface at pH 2 may facilitate the dehydration of cesium, which in turn stabilizes the neutral form of the silanol groups by displacing interfacial water molecules. To test this hypothesis, complementary measurements and molecular dynamic simulations are needed to understand the interfacial water structure in the presence of polarizable cations like cesium at different concentrations [21, 30, 33, 35–40]. In addition to the effective acidity of each site, the relative distribution of the more acidic and less acidic sites was also very ion specific in the presence of 0.5M electrolyte. However, in the presence of 0.1M alkali chloride, the differences among the various alkali ions were much less significant (figure 4(a)), which mirrors the trend in pH0.5 bulk values that varied little at 0.1M alkali chloride (table 1). Once again, the influence of concentration on the distribution of sites varied for the different cations. Specifically, upon decreasing the salt concentration from 0.5M to 0.1M, the fraction deprotonated at pH 7 (equivalent to the fraction of more acidic sites) decreased from 0.60 ± 0.03 to 0.35 ± 0.06 for LiCl, from 0.40 ± 0.07 to 0.27 ± 0.04 for KCl and from 0.52 ± 0.06 to 0.40 ± 0.06 for CsCl, indicating that the presence of these salts shifted the site population distribution to the more acidic sites (figure 3(b)). Similar observations have been reported for the colloidal silica/LiCl(aq) interface that noted a 1.5-fold increase in density of charged sites when the salt concentration increased from 0.1M to 0.5M at pH 7 [6], consistent with our measured value (a 1.7-fold increase). On the other hand, decreasing the NaCl concentration from 0.5M to 0.1M increased rather than decreased the number of more acidic sites, leading to an increase in the fraction deprotonated at pH 7 from 0.20 ± 0.02 to 0.44 ± 0.04. The anomalous behavior of sodium is difficult to explain, but could stem from the influence of sodium on the interfacial water structure and/or specific interactions with the siloxide surface that minimize the presence of the more acidic sites. For example, a recent study by Sulpizi et al proposed that the more acidic sites acted as hydrogen bond donors with water, while the less acidic sites accepted hydrogen bonds from water molecules and donated their hydrogen to a neighboring Si–OH group [16]. It is possible that Na+ stabilized the hydrogen bonding between two neighboring Si–OH groups in contrast to the other alkali ions. Finally, the unique behavior of sodium is also consistent with atomic force microscopy studies at the silica/water interface monitoring the interaction between a silica sphere and
+ (16) pQaeff ′ = pH 0.5 bulk + pMbulk
and + . (17) pH 0.5 bulk = pQaeff ′− pMbulk
According to this revised mechanism for the lithium, potassium and cesium chloride salts, the pH at which half the surface is deprotonated should decrease with decreasing salt concentration (corresponding to an increase in pM+), which was observed in our experiments at high pH (figures 2(d)–(f )). As shown in table 1, when the salt concentration was decreased to 0.1M from 0.5M, the pH0.5 bulk value decreased to 8.9 ± 0.1 from 9.9 ± 0.2 for LiCl, to 8.9 ± 0.3 from 10.2 ± 0.2 for KCl, and to 9.5 ± 0.2 from 11.1 ± 0.1 for CsCl. From these titration results and the corresponding salt concentrations, we measured interfacial pQaeff ′ values of 9.9 ± 0.2 and 10.2 ± 0.1 for silica/LiCl(aq), 9.9 ± 0.3 and 10.5 ± 0.1 for silica/KCl(aq), and 10.5 ± 0.2 and 11.4 ± 0.1 for silica/CsCl(aq) at 0.1M and 0.5M electrolyte, respectively, using equation (17) (table 1). We observed that the differences between the calculated values of pQaeff ′ for these two salt concentrations increased with increasing cation size or polarizability (LiCl < KCl < CsCl) suggesting that other equilibria besides those considered in equation (14) contributed to the salt-dependent behavior for KCl and CsCl. Nevertheless, the trend in salt concentration suggests that all alkali ions but Na+ inhibited deprotonation of the less acidic sites, effectively stabilizing the less charged form of the surface. For the deprotonation of the less acidic sites that occurred above pH 7, it is important to note that the surface at pH 7 stabilized in the presence of Li+, K+ and Cs+ was still negatively charged. Thus, electrostatic interactions with these cations and the surface were possible at this pH although they were expected to increase at higher pH. The deprotonation of the more acidic sites did not exhibit significant changes in the pH0.5 bulk values for LiCl and KCl at the two different concentrations studied (figures 3(a) and (b)). Yet there was a slight decrease in the pH0.5 bulk value for the silica/CsCl(aq) electrolyte interface at low pH upon decreasing the CsCl concentration indicating that the CsCl stabilized the neutral surface found at pH 2 more than that at neutral bulk pH. It is highly unusual for a cation to stabilize a neutral site over a negatively charged site. We speculate that this counterintuitive result rises from the weak interactions between 5
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 3. Representative plots of normalized E2ω as a function of pH for silica/electrolyte interfaces of various salt concentrations at (a)–(c)
low pH and (d)–(f) high pH. The solid lines (for 0.5M) and dashed lines (for 0.1M) represent the fit of a sigmoidal function to the data.
the surface concentrations, then equation (10) (for the sodium ion) and equation (15) (for the other alkali ions) can be rewritten in terms of the effective equilibrium constants Kaeff :
the native oxide coating on a Si(1 0 0) wafer [33]. In these studies conducted at pH 5.5, Sivan and co-workers found that the interactions between the silica surfaces were attractive in the presence of 0.1M KCl and CsCl. In contrast, the interactions were strongly repulsive in the presence of 0.1M NaCl. However, at 0.5M the situation was markedly different; the surfaces were attractive in the presence of NaCl and the magnitude of the force was consistent with van der Waals interactions between effectively neutralized surfaces. At this pH and salt concentration, Cs+ began to exhibit repulsive interactions at very short range, which the authors attributed to charge reversal of the interface by excess Cs+. K+ led to slightly less attractive interactions between the surfaces at this salt concentration and pH compared with the data at 0.1M. In light of these results and the proposed model of surface acidity by Sulpizi and co-workers, we suggest that the ability of sodium to become dehydrated and form close interactions with the negative surface sites occurs at concentrations greater than 0.1M. Once a high enough salt concentration is reached, the sodium is able to interact strongly with the negative surface sites and promote the formation of a hydrogen-bonded network of silanols leading to a majority of less acidic sites on the surface. In contrast, the presence of all of the other salts in the compact Stern layer disturbs this hydrogen-bonding network, decreasing the number of less acidic sites. Thus far we have utilized the bulk hydronium and cation concentrations to determine the pQaeff values. However, most surface models consider that the surface concentrations of the hydronium ion and cations are not equal to their corresponding bulk concentrations but can be related to the bulk concentration by a potential dependent term. If we consider
[SiO− ⋅ M+ ][H3O+ ]surface . [SiOH][ M+ ]surface
(18)
[SiO− ⋅ M+ ][ M+ ]surface [H3O+ ]surface . [SiOH]surface
(19)
Kaeff = Kaeff ′ =
According to the Stern model, the hydronium ions can partition to the 0-plane, while the larger cations are present at the d-plane (the 0-plane and d-plane, sometimes referred to as the inner and outer Helmholtz planes, respectively) (figure 1). The magnitude of the potential at their respective planes determines the extent to which the hydronium or cations accumulate at the interface according to ⎛ eΦ ⎞ (20) [H3O+ ]surface = [H3O+ ]bulk exp ⎜ − 0 ⎟ ⎝ kT ⎠
and
⎛ eΦ ⎞ (21) [ M+ ]surface = [ M+ ]bulk exp ⎜ − d ⎟ ≈ [ M+ ]bulk ⎝ kT ⎠
where e is the elementary charge, Φ is the potential at the corresponding plane, k is the Boltzmann constant and T is the temperature in kelvin. At high salt concentrations, the layer of cations in the compact Stern layer can be considered to screen all of the surface charges leading to a Φd of zero, which is effectively the constant capacitance model (figure 1). Consequently, we only consider the potential dependence of the surface hydronium concentration. 6
AQ2
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 4. (a) The fraction of deprotonated surface sites as a function of solution pH for the silica/electrolyte interface at 0.1M salt concentration. (b) The fraction deprotonated of the silica/electrolyte interface at pH 7 in the presence of 0.1M and 0.5M alkali chloride electrolyte.
Figure 5. The interfacial potential Φ0 as a function of surface pH for the silica/electrolyte interface at (a) 0.1M and (b) 0.5M concentration. The data were compiled from multiple experiments at low and high pH and smoothed using a box filter.
To determine the interfacial potential at a given bulk pH, we used the relationship
Table 1 represents the values obtained for the pKaeff and pKaeff ′ of the less acidic sites for all the alkali chlorides at 0.1M and 0.5M salts determined from these pH0.5 surface values. We observed that the trend in effective pKa values obtained for both the 0.5M and 0.1M concentrations was pKaeff (NaCl) < pKaeff ′(LiCl) ≈ pKaeff ′(KCl) < pKaeff ′(CsCl). The difference in the highest and lowest effective pKa values was ~2 log units when the surface concentration of hydronium was taken into account. This range is smaller than that of the pQaeff , which indicates that the second pKa at higher pH is effected by the relative distribution of more and less acidic sites as this changes the relationship between the bulk and surface pH. Despite the influence of the relative population of sites, the span of pKaeff values also reveals significant variations in the intrinsic acidity of the silanol sites depending on the electrolyte present, which is consistent with our proposed mechanism of ion binding and its influence on the overall acid dissociation equilibria. Finally, we note that others have observed that the relative affinity for the silica/aqueous surface followed the order Cs+ > K+ > Na+ below pH 7 [1] and the reverse order above pH 7 [7]. This pH-dependent
σ max (22) Φ0 = fΦ0max = f 0 C
where f is the fraction of deprotonation determined by combining the experimental data for the low pH and high pH titrations as described in our previous work [18] and Φ0max is the maximum interfacial potential at high pH that depends on the maximum surface charge density of planar silica (σ0max = 0.139 C m−2) and C the capacitance. The capacitance is expected to vary with electrolyte composition, but for simplicity we used a common value of 1.0 F m−2 for all of the electrolytes as the exact composition of the compact Stern layer at these high salt concentrations is not known. With the resulting Φ0max value of 0.14 V, the surface pH was determined and plotted against the normalized interfacial potential, which is proportional to the fraction deprotonated (see below) (figure 5). Thus, the pH0.5 surface values were determined from the separate fit to the low and high pH ranges as this normalized interfacial potential is equivalent to the fraction of deprotonated (SiO−) sites. 7
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 6. Representative traces of the normalized SHG field as a function of bulk pH at the silica/NaI(aq) electrolyte interface of various salt concentrations at (a) low and (b) high pH. (c) The combined low and high pH data representing the normalized interfacial potential versus bulk pH for the silica/NaI(aq) electrolyte interface. (d), (e) Representative traces from the silica/KI(aq) electrolyte interface at (d) low and (e) high pH. (f) The normalized interfacial potential versus bulk pH for the silica/KI(aq) electrolyte interface.
reversal of trends could stem from the stronger interactions of the more polarizable Cs+ and K+ with the neutral surface sites that dominate at lower pH.
ΔG interaction ⎛⎜ 1⎞ (24) pKacoop = pKaint + 1− ⎟ . 2.303RT ⎝ n⎠
This pKaint is related to the cation specific pQaeff discussed in the previous section as the pH0.5 bulk values are utilized to determine pKaint [19] and, therefore, it is expected to vary with the cation identity. (Owing to the sharpness of the titration curves for the NaI and KI systems, calculating the surface pH led to multiple E2ω values for the same surface pH. Consequently the bulk pH was used in determining the equilibrium constants for the cooperative model.) Here ΔG interaction is the free energy of interaction between neighboring sites that behave cooperatively [41]. This value was found to be negative and therefore attractive for the more acidic sites corresponding to a ΔG interaction of –4RT [19]. In contrast, the less acidic sites exhibited repulsive interactions between neighboring sites, which corresponded to a ΔG interaction of 4RT [19]. To test the influence of concentration on cooperativity, we performed acid–base titrations of the silica/electrolyte interface in the presence of 0.1M chloride and iodide electrolytes and compared the observed thermodynamic parameters with those from the 0.5M experiments (figure 6). Interestingly, the positions and steepness of the curves and the corresponding pKacoop and n values remained essentially unchanged for the NaI electrolyte system at 0.5M and 0.1M (figures 6(a) and (b) and table 2). Moreover, the fraction deprotonated at pH 7 equivalent to the fraction of more acidic sites remained very large even at the lower NaI concentration (figure 6(c)).
4.2. Influence of the sodium halide concentration
Unlike the alkali ions, which affected the intrinsic acidity of the two different surface sites, our previous work established that the presence of the halide ions influenced cooperativity among neighboring sites, more so than their intrinsic acidity [19]. The effect of cooperativity was quantified using a cooperative form of the Henderson Hasselbach equation [41]: [SiO− ] log = n (pHbulk ) − n (pKacoop ) (23) [SiOH]
where n is the cooperative unit and pKacoop is the overall equilibrium constant of the coupled silanol deprotonation events that we have proposed stem from the presence of the halide ion [19]. When n > 1 the deprotonation of one site facilitates deprotonation of neighboring sites (positive cooperativity). Conversely, when n < 1, the deprotonation of one site increases the energy cost of deprotonating neighboring sites (negative cooperativity). The [SiO−] : [SiOH] ratio can be determined from the relative surface charge density of a site according to equations (5) and (6). Thus, from the slope of log(σrel/(1– σrel)) versus pH, the cooperative number can be extracted. In this model, pKacoop is equal to pH0.5 bulk and can be used to quantify the intrinsic acid dissociation constant of a given silanol site pKaint according to 8
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Table 2. Parameters for the more acidic and less acidic sites at different concentrations of chloride and iodide electrolyte based on the cooperative equilibrium modela.
More acidic sites coop a
Salt
Conc. (M)
pK
NaCl
0.1 0.5 0.1 0.5 0.1 0.5 0.1 0.5
4.5(3) 4.06(4) 3.36(3) 3.2(1) 4.9(1) 4.9(3) 3.6(4) 3.4(2)
NaI KCl KI
Less acidic sites
N
pKaint
pK
0.75(3) 0.94(3) 3.2(6) 2.9(2) 1.06(2) 0.99(3) 3.0(5) 2.9(3)
3.9 4.0 4.5 4.3 4.8 4.9 4.8 4.5
9.6(1) 8.6(1) 11.2(1) 11.3(1) 8.9(3) 10.2(2) 10.3(4) 11.5(2)
coop a
n
pKaint
0.76(5) 0.55(1) 1.11(8) 1.17(4) 0.44(2) 0.50(1) 0.66(4) 1.10(6)
10.1 10.0 11.0 11.0 11.1 11.9 11.2 11.3
a The values in parentheses represent the standard deviation or range of measured values such that 3.4(2) is equivalent to 3.4 ± 0.2 and 1.10(6) is 1.10 ± 0.06.
The consistency in the pKaint values for NaCl at different salt concentrations suggests that the cooperative model might be sufficient for understanding the effects of alkali chloride concentration on the acid–base behavior of the interface. However, the pKaint values for the KCl system at the two different concentrations varied more significantly than the corresponding ΔpQaeff ′ values from the salt-dependent model described earlier ( ΔpKaint = 0.8 and ΔpQaeff ′ = 0.6, respectively), which indicates that the cooperative model is not as well suited for understanding the effects of electrolyte concentration for the alkali chlorides. Future work in conjunction with molecular dynamic simulations will be aimed at devising a unifying model that considers both the effect of cooperativity and cation specific interactions with each site.
These observations indicate that the iodide had a strong interaction with the interface even at lower concentrations. More importantly, these results support that cooperativity and anion effects on the population distribution of sites persist at lower salt concentrations, with important consequences for the accurate modeling of geochemical systems. For the KI electrolyte, little variation in the traces at low and high salt concentration as well as the associated pKacoop and n values was observed for the more acidic sites (figure 6(c), table 2). However, in contrast to the NaI system, for the less acidic silanol groups both the pKacoopand n values varied upon changing the KI concentration (figure 6(d)). Indeed, lowering the electrolyte concentration from 0.5M to 0.1M decreased the pKacoop value from 11.5 ± 0.2 to 10.3 ± 0.4, and n value from 1.10 ± 0.06 to 0.66 ± 0.04 for the less acidic sites at the silica/KI(aq) interface (table 2). We attributed this decrease in n to the changes in the extent of surface charging at neutral pH. According to our cooperative model, as the density of the less acidic sites increased, these sites exhibited more negative cooperativity owing to repulsive interactions between neighboring sites. As the fraction of deprotonated sites at neutral pH decreased from 0.91 ± 0.02 to 0.77 ± 0.07 for the 0.5M and 0.1M KI interfaces, respectively, the density of less acidic groups increased. Consequently, we observed a decrease in the cooperative number owing to increased repulsive interactions. Interestingly, varying the potassium iodide concentration changed the cooperative number, n, in such a way that the value of the intrinsic pKaint remained unchanged within ±0.1 (table 2), which provides further evidence that the cooperative model accurately describes the effect of iodide. Regarding the chloride data, like the case of KI described above we observed that the cooperative model could completely account for the changes in the pH0.5 bulk values observed for 0.1M and 0.5M NaCl based on the corresponding change in the steepness of the titration curve and the n values, which resulted in very similar Kaint values at these two NaCl concentrations. Additionally, the cooperative number for the less acidic sites at the NaCl(aq)/silica interface decreased as the fraction deprotonated at pH 7 decreased, which for NaCl uniquely occurred upon increasing the salt concentration, as the resulting higher density of less acidic sites facilitated more repulsive nearest neighbor interactions and a lower n value.
5. Conclusion We have explored the influence of electrolyte concentration for various alkali halides at the fused silica/aqueous interface. The influence of the ion can be quantified from the effective acid dissociation constants pKaeff and the change in distribution of more and less acidic sites manifested in the change in the fraction of the surface that is deprotonated at pH 7. For the alkali chlorides explored, we observed that the pKaeff of both types of sites increased with increasing cation size/polarizability. We attribute the lowest pKaeff of NaCl to the strong interactions between the sodium and the siloxide sites, suggesting that the SiO− sites have similar water affinities to that of Na+, which is reasonable given their similarity in polarizability and ionic radius. The order of LiCl < KCl < CsCl with respect to the remaining pKaeff values is attributed to the decreasing value of Kd corresponding to the dissociation of the cation from the neutral surface sites. From this we infer that Cs+ associates most strongly with the neutral sites followed by K+ and Li+. To fully understand the origin of these interactions, the hydrogen-bonding structure of the water must be ascertained and is the subject of current sum frequency generation studies in our group. We note that these specific cation effects are relatively weak at −0.1 M, which indicates that they may not be as relevant over the range of salinities found in the environment. In contrast, the specific anion effects were very insensitive to ion concentration particularly for the NaI 9
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
[15] Fisk J D, O’Reilly J P and Shaw A M 2006 J. Phys. Chem. B 110 15039 [16] Sulpizi M, Gaigeot M-P and Sprik M 2012 J. Chem. Theory Comput. 8 1037 [17] Leung K, Nielsen I M B and Criscenti L J 2009 J. Am. Chem. Soc. 131 18358 [18] Azam M S, Weeraman C N and Gibbs-Davis J M 2012 J. Phys. Chem. Lett. 3 1269 [19] Azam M S, Weeraman C N and Gibbs-Davis J M 2013 J. Phys. Chem. C 117 8840 [20] Abendroth R P 1970 J. Colloid Interface Sci. 34 591 [21] Campen R K, Pymer A K, Nihonyanagi S and Borguet E 2010 J. Phys. Chem. C 114 8465 [22] Eftekhari-Bafrooei A and Borguet E 2009 J. Am. Chem. Soc. 131 12034 [23] Kitamura A, Fujiwara K, Yamamoto T, Nishikawa S and Moriyama H 1999 J. Nucl. Sci. Technol. 36 1167 [24] Azam M S and Gibbs-Davis J M 2013 Anal. Chem. 85 8031 [25] Gibbs-Davis J M, Kruk J J, Konek C T, Scheidt K A and Geiger F M 2008 J. Am. Chem. Soc. 130 15444 [26] Shen Y R 1984 The Principles of Nonlinear Optics (New York: Wiley-Interscience) [27] Hayes P L, Malin J N, Jordan D S and Geiger F M 2010 Chem. Phys. Lett. 499 183 [28] Lützenkirchen J 1999 J. Colloid Interface Sci. 217 8 [29] Yan E C Y, Liu Y and Eisenthal K B 1998 J. Phys. Chem. B 102 6331 [30] Jena K C, Covert P A and Hore D K 2011 J. Phys. Chem. Lett. 2 1056 [31] Davis J A, James R O and Leckie J O 1978 J. Colloid Interface Sci. 63 480 [32] Behrens S H and Grier D G 2001 J. Chem. Phys. 115 6716 [33] Dishon M, Zohar O and Sivan U 2009 Langmuir 25 2831 [34] Morel F M M and Hering J G 1993 Principles and Applications of Aquatic Chemistry (New York: Wiley) [35] Du Q, Freysz E and Shen Y R 1994 Phys. Rev. Lett. 72 238 [36] Ostroverkhov V, Waychunas G and Shen Y 2005 Phys. Rev. Lett. 94 046102 [37] Yang Z, Li Q and Chou K C 2009 J. Phys. Chem. C 113 8201 [38] Jena K C and Hore D K 2009 J. Phys. Chem. C 113 15364 [39] Dewan S, Yeganeh M S and Borguet E 2013 J. Phys. Chem. Lett. 4 1977 [40] Brown M A, Huthwelker T, Redondo A B, Janousch M, Faubel M, Arrell C A, Scarongella M, Chergui M and van Bokhoven J A 2012 J. Phys. Chem. Lett. 3 231 [41] Urry D W 2006 What Sustains Life? Consilient Mechanisms for Protein-Based Machines and Materials (New York: Springer)
electrolyte. As cooperativity is the predominant factor governing the surface acidity of these sites in the presence of iodide, we infer that this cooperative structure is so stable that it is able to form at 0.1 M concentrations. Moreover, the relative distribution of sites still favors the presence of the more acidic sites even at these lower iodide electrolyte concentrations, which supports that the iodide ion has a stronger affinity for the silica/aqueous interface than the chloride ion in the presence of sodium or potassium. Acknowledgments We acknowledge the Natural Sciences and Engineering Research Council of Canada and the Canada Foundation for Innovation for funding. References [1] Iler R K 1979 Chemistry of Silica—Solubility, Polymerization, Colloid and Surface Properties and Biochemistry (New York: Wiley) [2] Dove P M and Rimstidt J D 1994 Reviews in Mineralogy Series: The Silica Polymorphs vol 29 ed P Heaney et al (Chantilly, VA: Mineralogical Society of America) p 259 [3] Sahai N 2002 Environ. Sci. Technol 36 445 [4] Icenhower J P and Dove P M 2000 Geochim. Cosmochim. Acta 64 4193 [5] Dove P M 1999 Geochim. Cosmochim. Acta 63 3715 [6] Milonjic S K 1987 Colloids Surf. 23 301 [7] Dove P and Craven C 2005 Geochim. Cosmochim. Acta 69 4963 [8] Eisenthal K B 1996 Chem. Rev. 96 1343 [9] de Beer A, Campen R and Roke S 2010 Phys. Rev. B 82 235431 [10] Fisk J D, Batten R, Jones G, O’Reill J P and Shaw A M 2005 J. Phys. Chem. B 109 14475 [11] O’Reilly J P, Butts C P, I’Anson I A and Shaw A M 2005 J. Am. Chem. Soc. 127 1632 [12] Ong S, Zhao X and Eisenthal K B 1992 Chem. Phys. Lett. 191 327 [13] Zhao X, Ong S, Wang H and Eisenthal K B 1993 Chem. Phys. Lett. 214 203 [14] Petersen P B and Saykally R J 2006 Annu. Rev. Phys. Chem. 57 333
10
Search
Collections
Journals
About
Contact us
My IOPscience
The influence of concentration on specific ion effects at the silica/water interface
This content has been downloaded from IOPscience. Please scroll down to see the full text. 2014 J. Phys.: Condens. Matter 26 244107 (http://iopscience.iop.org/0953-8984/26/24/244107) View the table of contents for this issue, or go to the journal homepage for more Download details: IP Address: 132.239.1.231 This content was downloaded on 10/06/2017 at 07:06 Please note that terms and conditions apply.
You may also be interested in: The electric double layer at a rutile TiO2 water interface modelled using density functional theory based molecular dynamics simulation J Cheng and M Sprik Electrostatic interaction forces in aqueous salt solutions of variable concentration andvalency Daniel Ebeling, Dirk van den Ende and Frieder Mugele Density functional theory simulation of hydrogen-bonding structure and vibrational densities of states at the quartz (101)-water interface and its relation to dissolution as a function of solution pH and ionic strength M J DelloStritto, J Kubicki and J O Sofo Ions at hydrophobic interfaces Yan Levin and Alexandre P dos Santos Molecular dynamics studies of aqueous silica nanoparticle dispersions: salt effects on the double layer formation Lucas S de Lara, Vagner A Rigo, Mateus F Michelon et al. Temperature responsive properties of poly(N-vinylcaprolactam) multilayer hydrogels in the presence of Hofmeister anions Oleksandra Zavgorodnya, Veronika Kozlovskaya, Xing Liang et al. Oxo-acidity in ionic melts Viktor L Cherginets Ion specific correlations in bulk and at biointerfaces I Kalcher, D Horinek, R R Netz et al.
Journal of Physics: Condensed Matter J. Phys.: Condens. Matter 26 (2014) 244107 (10pp)
doi:10.1088/0953-8984/26/24/244107
The influence of concentration on specific ion effects at the silica/water interface Md Shafiul Azam, Akemi Darlington and Julianne M Gibbs-Davis Department of Chemistry, University of Alberta, Edmonton, Alberta T6G 2G2, Canada E-mail: [email protected] Received 13 November 2013, revised 2 February 2014 Accepted for publication 4 February 2014 Published 27 May 2014 Abstract
Second harmonic generation spectroscopy is a useful tool for monitoring changes in interfacial potential at buried insulator/liquid interfaces. Here we apply this technique to the silica/ aqueous interface and monitor the changes in interfacial potential while varying the pH in the presence of different alkali halides at 0.1M concentration. Within the pH range explored, the bimodal distribution of acidic sites on planar silica is clearly observed, corresponding to two types of acidic SiOH groups. Comparing these data with previous work at 0.5M sheds light on whether the presence of the ions stabilizes the charged or neutral state of the surface sites. For the alkali chlorides, with the exception of NaCl, we observe that the presence of the alkali chlorides stabilize the less acidic site in the protonated (SiOH) rather than deprotonated (SiO−) form. This unusual influence of the cation is attributed to the combination of interactions at the interface between water, surface sites and the electrolyte. Overall, we observe that the influence of the alkali ion on the ratio of the two types of sites and their effective acid dissociation constants is minor at 0.1M, unlike that observed at 0.5M. In contrast, the influence of the anion on the cooperative dissociation of surface sites and their relative distribution is little affected upon decreasing the concentration, which indicates that these specific anion effects are prevalent in nature. Keywords: specific ion effects, buried interfaces, nonlinear optical spectroscopy (Some figures may appear in colour only in the online journal)
1. Introduction
represents environmentally relevant conditions using colloidal morphologies. One strategy to avoid the instability of colloids is to monitor planar silica surfaces. Such macroscopic substrates are amenable to different spectroscopic methods such as evanescent wave cavity ring down spectroscopy (EW-CRDS [10, 11]) and nonlinear optical methods like second harmonic generation (SHG) spectroscopy [12, 13]. In one approach, pH-sensitive dyes have been used in conjunction with EW-CRDS to probe the change in resonant frequency of the dye upon interacting with the surface [10, 11]. However, many have questioned whether the dye’s acid dissociation equilibrium constant Ka is different at the interface, knowing which is critical to accurately quantify the surface acidity using this approach [14, 15]. Another technique that represents a label-free method for monitoring the acid–base chemistry of the silica interface is based on the dependence of SHG on the interfacial potential. The first report by Eisenthal and co-workers utilizing
The silica/water interface is abundant in the natural environment, and consequently the properties of silica have been the subject of research for the last several decades [1–3]. In particular, a variety of techniques has been applied to the study of deprotonation and dissolution processes at the silica/water and functionalized silica/water interface [4–6]. Both bulk [7] and surface specific [8, 9] techniques have been employed, providing complementary pieces of information, which aid in forming a complete picture of these interfacial processes. Many of these studies have focused on colloidal silica and have included the effect of the aqueous electrolyte on surface deprotonation and dissolution events [1]. One challenge, however, in utilizing silica colloid dispersions is that they are not stable over a large pH range [1]. Consequently, it is difficult to map out the entire acid–base equilibria of silica over the broad pH range that 0953-8984/14/244107+10$33.00
1
© 2014 IOP Publishing Ltd Printed in the UK
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
this non-resonant SHG method to monitor the silica/aqueous interface found that planar silica exhibited two distinct deprotonation processes between pH 2 and 13 [12, 13]. Both of these acidic sites have been attributed to SiOH groups that are in two different hydrogen bonding environments resulting in the drastic difference in surface acidities. The exact nature of these sites has been the subject of computational investigations and varying origins for the surface acidity have been proposed [16, 17]. Recently we have utilized the same non-resonant SHG method referred to as the χ(3) technique to probe the influence of different alkali halides on the acid–base properties of the silica/aqueous interface over a broad pH range [18, 19]. We observed that the acidity of the silica surface depended greatly on the composition of the aqueous electrolyte. Indeed, in the presence of various alkali chlorides, the effective acid equilibrium constants (Ka) ranged from 10−8 to 10−11 for the less acidic silanol groups and decreased with increasing alkali ion polarizability [18]. The trend in Ka and cation polarizability was also observed for the more acidic sites deprotonated below pH 7, but the range of constants was smaller. Additionally, the relative distribution of the two types of silanol sites also varied widely. Specifically our work confirmed the previous results of Eisenthal, conducted with 0.5M NaCl, which showed that 20% of the silica sites were the more acidic sites that were deprotonated below pH 7. However, substituting NaCl with LiCl led to a shift in the distribution of these sites such that 60% of the surface was now composed of the more acidic sites, which were deprotonated below pH 7. Our observations suggested that the surface charge density of silica could vary dramatically at the same pH, depending on the composition of the aqueous phase, consistent with previous reports examining the acid–base behavior of the silica/water interface [20–23]. More recently, we examined the role of the anion on deprotonation processes at this interface [19]. Despite the fact that the silica/water interface is negatively charged above pH 2, the anion has an even larger influence on the apparent Ka of the surface sites as well as their relative distribution. For example, utilizing NaI in place of NaCl led to very sharp deprotonation curves in the surface titration experiments, which are indicative of positive cooperativity whereby the dissociation of one silanol makes subsequent deprotonation events more favorable for a certain number of neighboring sites. Additionally, the ratio of more acidic to less acidic silanol sites increased from 1 : 4 with 0.5M NaCl to ~4 : 1 with 0.5M NaI. For the latter, this ratio corresponded to 85% of the surface becoming deprotonated at neutral pH, very different from the 20% deprotonated sites in the presence of NaCl. Such specific anion effects had not been explored at a negative interface, and this was the first demonstration of cooperativity in these systems. Here we gain further insight into the origins and mechanism of these specific ion effects at the silica/water interface by performing these SHG experiments at a lower salt concentration (0.1M). By reducing the salt concentration and monitoring whether the surface becomes more or less acidic, we can probe the role of the ions in inhibiting or promoting surface deprotonation. Finally, monitoring the interface at this lower salt concentration sheds light on how robust these
specific ion interactions are at more relevant environmental and biological concentrations. 2. Methods: SHG experiments A detailed description of the laser system used in these experiments can be found in our previously published work [18, 19, 24]. Briefly, using the output of a regeneratively amplified laser system (Spitfire Pro, Spectra Physics, 1 kHz, 100 fs, 3.3 W) to pump an optical parametric amplifier (Spectra Physics OPA-800CF) the incident light was tuned to 550 ± 2 nm and directed through a half-wave plate and polarizer for s-polarization selection. This wavelength was selected as it leads to a second harmonic wavelength that is off resonance with respect to the electrolyte and is the same wavelength that was used for our experiments at 0.5M, which represents an important comparison for this work [18, 19]. The polarized light was focused onto the fused silica/water interface at an angle of 62° from surface normal near total internal reflection. The reflected second harmonic light generated at the interface was then passed through a colored glass filter (Thorlabs) to remove the reflected fundamental light and focused onto a monochromator (Optometrics Corp., MiniChrom MC1-02) tuned to the second harmonic wavelength (275 nm). SHG was detected using gated photon counting with a PMT assembly. Before performing each experiment, the quadratic power dependence and SHG wavelength dependence were verified. For all experiments, a freshly cleaned fused silica hemisphere (ISP optics, 1 inch diameter, QU-HS-25, UV-grade SiO2) was placed on a custom-built Teflon cell so that the flat surface of the hemisphere was in contact with the aqueous phase [18]. The silica hemisphere was cleaned prior to use by sonicating it in Millipore water, then methanol and water. Next, Nochromix (Godax Laboratories, 5% w/v solution in H2SO4) was placed on the flat side for 1 h followed by copious rinsing in Millipore water and further sonication in water, methanol then water. The hemisphere was allowed to dry in an oven at 100 °C for 10 min, then cooled down to room temperature and plasma cleaned (Plasma cleaner, PDC-32G, Harrick Plasma) in air for 2–3 min. Each sample was exposed to Milli-Q water until the SHG signal was optimized before adding the electrolyte solution. Separate experiments were conducted at low pH and high pH on different fresh samples to avoid hysteresis [25]. More details can be found in our previous publications [18, 19]. 3. Theoretical background: the χ (3) technique Second harmonic light is generated when light of frequency ω interacts with aligned dipoles at an interface causing induced polarization that oscillates at twice the incident frequency 2ω (P2ω) resulting in a new electric field at 2ω (E2ω). The intensity of the SHG response from the interface (I2ω) is proportional to the induced second-order polarization as follows [8, 26]: (1) I2ω = E2ω ∝ P2ω = χ (2) EωEω . 2
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
In this expression, Eω is the incident electric field and χ is the nonlinear susceptibility of the interface. When surfaces are charged the presence of a static electric field (E0) leads to an additional third-order contribution to the second harmonic field [12]: (2)
∫
E2ω ∝ χ (2) EωEω + χ (3) EωEω E0 dx = χ (2) EωEω + χ (3) EωEωΦ0 (2) where χ(3) is the third-order susceptibility of the interface. Integrating this static electric field that originates from the surface charges out into the bulk yields the interfacial potential Φ0. For the silica/aqueous interface, the surface is negatively charged above pH 2, which means that it is charged under most environmental conditions. Consequently, this χ(3) technique is well suited to probe processes at the silica surface that are dictated by electrostatic interactions, which modulate the interfacial potential [27]. For monitoring deprotonation processes, we wish to relate the change in E2ω directly to changes in the surface charge density of silica. Therefore, it is useful to perform the experiments at high salt concentrations (≥100 mM). At these high salt concentrations, we assume that the changes to χ(2) and χ(3) are negligible upon varying pH compared with the changes to Φ0 [12, 13, 18, 19, 25, 27, 29]. This assumption is based in part on the dependence of these susceptibilities on the amount of ordered water [30], which should vary minimally at high salt concentration as the Debye length is short and the static electric field decays rapidly into the bulk [31, 32]. In other words, the water molecules that experience the static electric field at the interface are primarily in the compact Stern layer of cations at these high salt concentrations. We reason that the net orientation and number of these water molecules may change slightly upon varying the pH, but not as significantly as the change in the static electric field or the corresponding interfacial potential upon changing the pH and deprotonating surface sites [19]. Furthermore, at such high salt concentrations the Stern model of the interface can be approximated by the constant capacitance model as the diffuse layer of ions has collapsed (figure 1) [28]. As a result, the surface charge density becomes directly proportional to the interfacial potential. From these assumptions, we can write the general equation:
Figure 1. SHG at the silica/electrolyte interface. The second harmonic field is proportional to the interfacial potential (Φ0) set up by charged surface sites. This charged interface can be represented by the basic Stern model, where the potential determining surface sites make up the 0-plane. According to this model a layer of cations forms at the interface, shown at the d-plane, which make up the Stern layer. Any remaining negative charges that are not neutralized by the Stern layer are neutralized by cations in the diffuse layer (not shown). At high salt concentrations (≥ 0.1M) the Stern model reduces to the constant capacitance model, which assumes that the surface is completely neutralized by the cations in the Stern layer leading to the disappearance of the diffuse layer and Φd = 0 [28].
and that of the less acidic sites rσLA is E2ω − E2ω (pH 7) , rσLA = (6) E2ω (pH 12) − E2ω (pH 7)
based on the observation that the more acidic sites are deprotonated between pH 2 and 7, while the less acidic sites are deprotonated between pH 7 and 12. We note that the validity of the constant capacitance model at 0.1M in the presence of NaCl is doubtful, as experimental data suggest that the diffuse layer is still present for this electrolyte at this concentration [28, 33]. However, for simplicity we utilize this model for all of the alkali chlorides investigated. 4. Results and discussion 4.1. Influence of the alkali chloride concentration
⎛σ⎞ (3) E2ω ∝ A + BΦ0 = A + B ⎜ ⎟ ⎝C ⎠
In our first report on specific ion effects we observed that the bulk pH value that corresponded to deprotonation of half of a given type of site, pH0.5 bulk, followed the trend NaCl < LiCl < KCl < CsCl for both the more and less acidic sites [18]. These data indicated that the presence of NaCl made both surface sites more acidic such that they were deprotonated at a lower pH, while CsCl made both surface sites less acidic. We proposed a simple model based on the affinity of the alkali ion for the siloxide sites generated after deprotonation, which is similar to commonly used surface complexation models for quantifying the effect of the electrolyte on surface acid–base equilibria [34]. In this proposed mechanism the acid dissociation of the surface silanol groups first occurred, leading to the siloxide conjugate base:
where A is equal to χ (2) EωEω and B is equal to χ (3) EωEω, which are constant, C is the capacitance of the Stern layer and σ is the surface charge density. At low pH near pH 2 the surface becomes net neutral and A = E2ω (pH 2). (4)
Therefore the relative surface charge density can be found by normalizing E2ω from zero to one. Specifically, the relative surface charge density stemming from the more acidic sites rσMA is E2ω − E2ω (pH 2) rσMA = (5) E2ω (pH 7) − E2ω (pH 2)
K
a (7) SiOH + H2 O ⇄ SiO− + H3O+ .
3
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107 +
The cation, M , then coordinated to the siloxide through electrostatic interactions. KSC (8) SiO− + M+ ⇄ SiO− ⋅ M+ .
Although not explicitly shown, this electrostatic interaction was presumed to be mediated through water, resulting in a screening interaction rather than neutralization of the negative charges [21]. Finally, the overall equilibrium expression with the effective acid dissociation constant (Kaeff ) was equal to the sum of the two separate equilibrium expressions: eff
Ka (9) SiOH + H2O + M+ ⇄ SiO− ⋅ M+ + H3O+ .
A useful parameter for exploring the acidity of surfaces that is model independent is the acidity quotient Qaeff , which is related to the equilibrium constant through a potential dependent term that accounts for the surface concentration of all species [31]. The expression related to equation (9) that utilizes the bulk concentrations is [SiO− ⋅ M+ ][H3O+ ]bulk (10) Qaeff = QaQSC = [SiOH][ M+ ]bulk
where Qa, QSC and Qaeff are quotients that can be related to the actual equilibrium constants Ka, KSC and Kaeff , respectively, by converting the bulk cation and hydronium concentrations to the surface concentrations (see bolow). The value of pQaeff (–log Qaeff ) can then be determined from pH0.5 bulk and the + value: –log[M+]bulk or pMbulk + . (11) pQaeff = pH 0.5 bulk − pMbulk
According to the proposed coupled equilibria, decreasing the cation concentration should shift the equilibrium shown in equation (9) to the left, resulting in a less acidic surface and higher pH0.5 bulk values. To test whether the presence of the cation stabilized the siloxide sites, we decreased the concentration of the cation to 0.1M for NaCl and monitored the resulting acid–base behavior (figure 2). As shown in fi gure 2, decreasing the Na+ cation concentration indeed led to an increase in the pH0.5 bulk value, which is consistent with the + value at lower salt and the rearranged form of higher pMbulk equation (11):
Figure 2. Representative plots of normalized E2ω as a function of pH for the silica/electrolyte interface with aqueous NaCl of different concentrations at (a) low pH and (b) high pH. The solid lines (for 0.5M) and dashed lines (for 0.1M) represent the fit of a sigmoidal function to the data.
the less acidic surface sites were deprotonated, contrary to equation (12) (figures 3(d)–(f )). These results suggested that the presence of alkali chlorides other than sodium chloride actually stabilized the neutral (SiOH) rather than the negative form (SiO−) of these less acidic silanol groups, thereby requiring more basic environments to achieve deprotonation. To rationalize the observed concentration-dependent behavior for these salts, we neglect association of the cation to the siloxide and instead consider that the first step is cation dissociation from the neutral silanol,
+ . (12) pH 0.5 bulk = pQaeff + pMbulk
Using the pH0.5 bulk values of 9.6 ± 0.1 and 8.6 ± 0.1 and the corresponding p[Na+] values in equation (12) yielded pQaeff values of 8.6 ± 0.1 and 8.3 ± 0.1 for 0.1 and 0.5M NaCl concentrations, respectively (table 1). The decent agreement for the two pQaeff values suggests that our simple model whereby sodium stabilizes the siloxide provides a reasonable interpretation of its influence on the surface equilibria. Very similar salt-dependent behavior was also observed for the nanoporous silica/NaCl electrolyte interface in a recent SHG experiment by Borguet and co-workers, indicating this is general for different silica morphologies [21]. However, repeating the concentration variation measurement for the other cations led to very different results (figure 3). Specifically, decreasing the salt concentration of Li+, K+ and Cs+ led to a decrease in the pH at which half
K
d (13) SiOH ⋅ M+ ⇄ SiOH + M+,
which leads to the following overall equilibrium expression: KdKa
(14) SiOH ⋅ M+ + H2 O ⇄ SiO− + M+ + H3O+ . The salt-dependent equation that results is [SiO− ][ M+ ]bulk [H3O+ ]bulk Qaeff ′ = QdQa = . (15) [SiOH] 4
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Table 1. Thermodynamic parameters from the alkali chloride variation experiments . a
Salt
[Salt] (M)
pH0.5 bulk
pQaeff
pQaeff ′
pH0.5 surface
pKaeff
pKaeff ′
LiCl
0.1 0.5 0.1 0.5 0.1 0.5 0.1 0.5
8.9 ± 0.1 9.9 ± 0.2 9.6 ± 0.1 8.6 ± 0.1 8.9 ± 0.3 10.2 ± 0.2 9.5 ± 0.2 11.1 ± 0.1
— — 8.6 ± 0.1 8.3 ± 0.1 — — — —
9.9 ± 0.2 10.2 ± 0.2 — — 9.9 ± 0.3 10.5 ± 0.2 10.5 ± 0.2 11.4 ± 0.1
7.4 ± 0.1 8.0 ± 0.2 7.8 ± 0.1 7.3 ± 0.1 7.4 ± 0.3 8.5 ± 0.2 7.9 ± 0.2 9.2 ± 0.1
— — 6.8 ± 0.1 7.0 ± 0.1 — — — —
8.4 ± 0.2 8.3 ± 0.2 — — 8.4 ± 0.3 8.8 ± 0.2 8.9 ± 0.2 9.5 ± 0.1
NaCl KCl CsCl
a All the given values are the average of at least two independent measurements and the error represents the range of measured values.
At pH0.5 bulk when half of the surface sites are deprotonated (i.e. when [SiO−] = [SiOH])
cesium and water. The net neutral silica surface at pH 2 may facilitate the dehydration of cesium, which in turn stabilizes the neutral form of the silanol groups by displacing interfacial water molecules. To test this hypothesis, complementary measurements and molecular dynamic simulations are needed to understand the interfacial water structure in the presence of polarizable cations like cesium at different concentrations [21, 30, 33, 35–40]. In addition to the effective acidity of each site, the relative distribution of the more acidic and less acidic sites was also very ion specific in the presence of 0.5M electrolyte. However, in the presence of 0.1M alkali chloride, the differences among the various alkali ions were much less significant (figure 4(a)), which mirrors the trend in pH0.5 bulk values that varied little at 0.1M alkali chloride (table 1). Once again, the influence of concentration on the distribution of sites varied for the different cations. Specifically, upon decreasing the salt concentration from 0.5M to 0.1M, the fraction deprotonated at pH 7 (equivalent to the fraction of more acidic sites) decreased from 0.60 ± 0.03 to 0.35 ± 0.06 for LiCl, from 0.40 ± 0.07 to 0.27 ± 0.04 for KCl and from 0.52 ± 0.06 to 0.40 ± 0.06 for CsCl, indicating that the presence of these salts shifted the site population distribution to the more acidic sites (figure 3(b)). Similar observations have been reported for the colloidal silica/LiCl(aq) interface that noted a 1.5-fold increase in density of charged sites when the salt concentration increased from 0.1M to 0.5M at pH 7 [6], consistent with our measured value (a 1.7-fold increase). On the other hand, decreasing the NaCl concentration from 0.5M to 0.1M increased rather than decreased the number of more acidic sites, leading to an increase in the fraction deprotonated at pH 7 from 0.20 ± 0.02 to 0.44 ± 0.04. The anomalous behavior of sodium is difficult to explain, but could stem from the influence of sodium on the interfacial water structure and/or specific interactions with the siloxide surface that minimize the presence of the more acidic sites. For example, a recent study by Sulpizi et al proposed that the more acidic sites acted as hydrogen bond donors with water, while the less acidic sites accepted hydrogen bonds from water molecules and donated their hydrogen to a neighboring Si–OH group [16]. It is possible that Na+ stabilized the hydrogen bonding between two neighboring Si–OH groups in contrast to the other alkali ions. Finally, the unique behavior of sodium is also consistent with atomic force microscopy studies at the silica/water interface monitoring the interaction between a silica sphere and
+ (16) pQaeff ′ = pH 0.5 bulk + pMbulk
and + . (17) pH 0.5 bulk = pQaeff ′− pMbulk
According to this revised mechanism for the lithium, potassium and cesium chloride salts, the pH at which half the surface is deprotonated should decrease with decreasing salt concentration (corresponding to an increase in pM+), which was observed in our experiments at high pH (figures 2(d)–(f )). As shown in table 1, when the salt concentration was decreased to 0.1M from 0.5M, the pH0.5 bulk value decreased to 8.9 ± 0.1 from 9.9 ± 0.2 for LiCl, to 8.9 ± 0.3 from 10.2 ± 0.2 for KCl, and to 9.5 ± 0.2 from 11.1 ± 0.1 for CsCl. From these titration results and the corresponding salt concentrations, we measured interfacial pQaeff ′ values of 9.9 ± 0.2 and 10.2 ± 0.1 for silica/LiCl(aq), 9.9 ± 0.3 and 10.5 ± 0.1 for silica/KCl(aq), and 10.5 ± 0.2 and 11.4 ± 0.1 for silica/CsCl(aq) at 0.1M and 0.5M electrolyte, respectively, using equation (17) (table 1). We observed that the differences between the calculated values of pQaeff ′ for these two salt concentrations increased with increasing cation size or polarizability (LiCl < KCl < CsCl) suggesting that other equilibria besides those considered in equation (14) contributed to the salt-dependent behavior for KCl and CsCl. Nevertheless, the trend in salt concentration suggests that all alkali ions but Na+ inhibited deprotonation of the less acidic sites, effectively stabilizing the less charged form of the surface. For the deprotonation of the less acidic sites that occurred above pH 7, it is important to note that the surface at pH 7 stabilized in the presence of Li+, K+ and Cs+ was still negatively charged. Thus, electrostatic interactions with these cations and the surface were possible at this pH although they were expected to increase at higher pH. The deprotonation of the more acidic sites did not exhibit significant changes in the pH0.5 bulk values for LiCl and KCl at the two different concentrations studied (figures 3(a) and (b)). Yet there was a slight decrease in the pH0.5 bulk value for the silica/CsCl(aq) electrolyte interface at low pH upon decreasing the CsCl concentration indicating that the CsCl stabilized the neutral surface found at pH 2 more than that at neutral bulk pH. It is highly unusual for a cation to stabilize a neutral site over a negatively charged site. We speculate that this counterintuitive result rises from the weak interactions between 5
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 3. Representative plots of normalized E2ω as a function of pH for silica/electrolyte interfaces of various salt concentrations at (a)–(c)
low pH and (d)–(f) high pH. The solid lines (for 0.5M) and dashed lines (for 0.1M) represent the fit of a sigmoidal function to the data.
the surface concentrations, then equation (10) (for the sodium ion) and equation (15) (for the other alkali ions) can be rewritten in terms of the effective equilibrium constants Kaeff :
the native oxide coating on a Si(1 0 0) wafer [33]. In these studies conducted at pH 5.5, Sivan and co-workers found that the interactions between the silica surfaces were attractive in the presence of 0.1M KCl and CsCl. In contrast, the interactions were strongly repulsive in the presence of 0.1M NaCl. However, at 0.5M the situation was markedly different; the surfaces were attractive in the presence of NaCl and the magnitude of the force was consistent with van der Waals interactions between effectively neutralized surfaces. At this pH and salt concentration, Cs+ began to exhibit repulsive interactions at very short range, which the authors attributed to charge reversal of the interface by excess Cs+. K+ led to slightly less attractive interactions between the surfaces at this salt concentration and pH compared with the data at 0.1M. In light of these results and the proposed model of surface acidity by Sulpizi and co-workers, we suggest that the ability of sodium to become dehydrated and form close interactions with the negative surface sites occurs at concentrations greater than 0.1M. Once a high enough salt concentration is reached, the sodium is able to interact strongly with the negative surface sites and promote the formation of a hydrogen-bonded network of silanols leading to a majority of less acidic sites on the surface. In contrast, the presence of all of the other salts in the compact Stern layer disturbs this hydrogen-bonding network, decreasing the number of less acidic sites. Thus far we have utilized the bulk hydronium and cation concentrations to determine the pQaeff values. However, most surface models consider that the surface concentrations of the hydronium ion and cations are not equal to their corresponding bulk concentrations but can be related to the bulk concentration by a potential dependent term. If we consider
[SiO− ⋅ M+ ][H3O+ ]surface . [SiOH][ M+ ]surface
(18)
[SiO− ⋅ M+ ][ M+ ]surface [H3O+ ]surface . [SiOH]surface
(19)
Kaeff = Kaeff ′ =
According to the Stern model, the hydronium ions can partition to the 0-plane, while the larger cations are present at the d-plane (the 0-plane and d-plane, sometimes referred to as the inner and outer Helmholtz planes, respectively) (figure 1). The magnitude of the potential at their respective planes determines the extent to which the hydronium or cations accumulate at the interface according to ⎛ eΦ ⎞ (20) [H3O+ ]surface = [H3O+ ]bulk exp ⎜ − 0 ⎟ ⎝ kT ⎠
and
⎛ eΦ ⎞ (21) [ M+ ]surface = [ M+ ]bulk exp ⎜ − d ⎟ ≈ [ M+ ]bulk ⎝ kT ⎠
where e is the elementary charge, Φ is the potential at the corresponding plane, k is the Boltzmann constant and T is the temperature in kelvin. At high salt concentrations, the layer of cations in the compact Stern layer can be considered to screen all of the surface charges leading to a Φd of zero, which is effectively the constant capacitance model (figure 1). Consequently, we only consider the potential dependence of the surface hydronium concentration. 6
AQ2
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 4. (a) The fraction of deprotonated surface sites as a function of solution pH for the silica/electrolyte interface at 0.1M salt concentration. (b) The fraction deprotonated of the silica/electrolyte interface at pH 7 in the presence of 0.1M and 0.5M alkali chloride electrolyte.
Figure 5. The interfacial potential Φ0 as a function of surface pH for the silica/electrolyte interface at (a) 0.1M and (b) 0.5M concentration. The data were compiled from multiple experiments at low and high pH and smoothed using a box filter.
To determine the interfacial potential at a given bulk pH, we used the relationship
Table 1 represents the values obtained for the pKaeff and pKaeff ′ of the less acidic sites for all the alkali chlorides at 0.1M and 0.5M salts determined from these pH0.5 surface values. We observed that the trend in effective pKa values obtained for both the 0.5M and 0.1M concentrations was pKaeff (NaCl) < pKaeff ′(LiCl) ≈ pKaeff ′(KCl) < pKaeff ′(CsCl). The difference in the highest and lowest effective pKa values was ~2 log units when the surface concentration of hydronium was taken into account. This range is smaller than that of the pQaeff , which indicates that the second pKa at higher pH is effected by the relative distribution of more and less acidic sites as this changes the relationship between the bulk and surface pH. Despite the influence of the relative population of sites, the span of pKaeff values also reveals significant variations in the intrinsic acidity of the silanol sites depending on the electrolyte present, which is consistent with our proposed mechanism of ion binding and its influence on the overall acid dissociation equilibria. Finally, we note that others have observed that the relative affinity for the silica/aqueous surface followed the order Cs+ > K+ > Na+ below pH 7 [1] and the reverse order above pH 7 [7]. This pH-dependent
σ max (22) Φ0 = fΦ0max = f 0 C
where f is the fraction of deprotonation determined by combining the experimental data for the low pH and high pH titrations as described in our previous work [18] and Φ0max is the maximum interfacial potential at high pH that depends on the maximum surface charge density of planar silica (σ0max = 0.139 C m−2) and C the capacitance. The capacitance is expected to vary with electrolyte composition, but for simplicity we used a common value of 1.0 F m−2 for all of the electrolytes as the exact composition of the compact Stern layer at these high salt concentrations is not known. With the resulting Φ0max value of 0.14 V, the surface pH was determined and plotted against the normalized interfacial potential, which is proportional to the fraction deprotonated (see below) (figure 5). Thus, the pH0.5 surface values were determined from the separate fit to the low and high pH ranges as this normalized interfacial potential is equivalent to the fraction of deprotonated (SiO−) sites. 7
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Figure 6. Representative traces of the normalized SHG field as a function of bulk pH at the silica/NaI(aq) electrolyte interface of various salt concentrations at (a) low and (b) high pH. (c) The combined low and high pH data representing the normalized interfacial potential versus bulk pH for the silica/NaI(aq) electrolyte interface. (d), (e) Representative traces from the silica/KI(aq) electrolyte interface at (d) low and (e) high pH. (f) The normalized interfacial potential versus bulk pH for the silica/KI(aq) electrolyte interface.
reversal of trends could stem from the stronger interactions of the more polarizable Cs+ and K+ with the neutral surface sites that dominate at lower pH.
ΔG interaction ⎛⎜ 1⎞ (24) pKacoop = pKaint + 1− ⎟ . 2.303RT ⎝ n⎠
This pKaint is related to the cation specific pQaeff discussed in the previous section as the pH0.5 bulk values are utilized to determine pKaint [19] and, therefore, it is expected to vary with the cation identity. (Owing to the sharpness of the titration curves for the NaI and KI systems, calculating the surface pH led to multiple E2ω values for the same surface pH. Consequently the bulk pH was used in determining the equilibrium constants for the cooperative model.) Here ΔG interaction is the free energy of interaction between neighboring sites that behave cooperatively [41]. This value was found to be negative and therefore attractive for the more acidic sites corresponding to a ΔG interaction of –4RT [19]. In contrast, the less acidic sites exhibited repulsive interactions between neighboring sites, which corresponded to a ΔG interaction of 4RT [19]. To test the influence of concentration on cooperativity, we performed acid–base titrations of the silica/electrolyte interface in the presence of 0.1M chloride and iodide electrolytes and compared the observed thermodynamic parameters with those from the 0.5M experiments (figure 6). Interestingly, the positions and steepness of the curves and the corresponding pKacoop and n values remained essentially unchanged for the NaI electrolyte system at 0.5M and 0.1M (figures 6(a) and (b) and table 2). Moreover, the fraction deprotonated at pH 7 equivalent to the fraction of more acidic sites remained very large even at the lower NaI concentration (figure 6(c)).
4.2. Influence of the sodium halide concentration
Unlike the alkali ions, which affected the intrinsic acidity of the two different surface sites, our previous work established that the presence of the halide ions influenced cooperativity among neighboring sites, more so than their intrinsic acidity [19]. The effect of cooperativity was quantified using a cooperative form of the Henderson Hasselbach equation [41]: [SiO− ] log = n (pHbulk ) − n (pKacoop ) (23) [SiOH]
where n is the cooperative unit and pKacoop is the overall equilibrium constant of the coupled silanol deprotonation events that we have proposed stem from the presence of the halide ion [19]. When n > 1 the deprotonation of one site facilitates deprotonation of neighboring sites (positive cooperativity). Conversely, when n < 1, the deprotonation of one site increases the energy cost of deprotonating neighboring sites (negative cooperativity). The [SiO−] : [SiOH] ratio can be determined from the relative surface charge density of a site according to equations (5) and (6). Thus, from the slope of log(σrel/(1– σrel)) versus pH, the cooperative number can be extracted. In this model, pKacoop is equal to pH0.5 bulk and can be used to quantify the intrinsic acid dissociation constant of a given silanol site pKaint according to 8
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
Table 2. Parameters for the more acidic and less acidic sites at different concentrations of chloride and iodide electrolyte based on the cooperative equilibrium modela.
More acidic sites coop a
Salt
Conc. (M)
pK
NaCl
0.1 0.5 0.1 0.5 0.1 0.5 0.1 0.5
4.5(3) 4.06(4) 3.36(3) 3.2(1) 4.9(1) 4.9(3) 3.6(4) 3.4(2)
NaI KCl KI
Less acidic sites
N
pKaint
pK
0.75(3) 0.94(3) 3.2(6) 2.9(2) 1.06(2) 0.99(3) 3.0(5) 2.9(3)
3.9 4.0 4.5 4.3 4.8 4.9 4.8 4.5
9.6(1) 8.6(1) 11.2(1) 11.3(1) 8.9(3) 10.2(2) 10.3(4) 11.5(2)
coop a
n
pKaint
0.76(5) 0.55(1) 1.11(8) 1.17(4) 0.44(2) 0.50(1) 0.66(4) 1.10(6)
10.1 10.0 11.0 11.0 11.1 11.9 11.2 11.3
a The values in parentheses represent the standard deviation or range of measured values such that 3.4(2) is equivalent to 3.4 ± 0.2 and 1.10(6) is 1.10 ± 0.06.
The consistency in the pKaint values for NaCl at different salt concentrations suggests that the cooperative model might be sufficient for understanding the effects of alkali chloride concentration on the acid–base behavior of the interface. However, the pKaint values for the KCl system at the two different concentrations varied more significantly than the corresponding ΔpQaeff ′ values from the salt-dependent model described earlier ( ΔpKaint = 0.8 and ΔpQaeff ′ = 0.6, respectively), which indicates that the cooperative model is not as well suited for understanding the effects of electrolyte concentration for the alkali chlorides. Future work in conjunction with molecular dynamic simulations will be aimed at devising a unifying model that considers both the effect of cooperativity and cation specific interactions with each site.
These observations indicate that the iodide had a strong interaction with the interface even at lower concentrations. More importantly, these results support that cooperativity and anion effects on the population distribution of sites persist at lower salt concentrations, with important consequences for the accurate modeling of geochemical systems. For the KI electrolyte, little variation in the traces at low and high salt concentration as well as the associated pKacoop and n values was observed for the more acidic sites (figure 6(c), table 2). However, in contrast to the NaI system, for the less acidic silanol groups both the pKacoopand n values varied upon changing the KI concentration (figure 6(d)). Indeed, lowering the electrolyte concentration from 0.5M to 0.1M decreased the pKacoop value from 11.5 ± 0.2 to 10.3 ± 0.4, and n value from 1.10 ± 0.06 to 0.66 ± 0.04 for the less acidic sites at the silica/KI(aq) interface (table 2). We attributed this decrease in n to the changes in the extent of surface charging at neutral pH. According to our cooperative model, as the density of the less acidic sites increased, these sites exhibited more negative cooperativity owing to repulsive interactions between neighboring sites. As the fraction of deprotonated sites at neutral pH decreased from 0.91 ± 0.02 to 0.77 ± 0.07 for the 0.5M and 0.1M KI interfaces, respectively, the density of less acidic groups increased. Consequently, we observed a decrease in the cooperative number owing to increased repulsive interactions. Interestingly, varying the potassium iodide concentration changed the cooperative number, n, in such a way that the value of the intrinsic pKaint remained unchanged within ±0.1 (table 2), which provides further evidence that the cooperative model accurately describes the effect of iodide. Regarding the chloride data, like the case of KI described above we observed that the cooperative model could completely account for the changes in the pH0.5 bulk values observed for 0.1M and 0.5M NaCl based on the corresponding change in the steepness of the titration curve and the n values, which resulted in very similar Kaint values at these two NaCl concentrations. Additionally, the cooperative number for the less acidic sites at the NaCl(aq)/silica interface decreased as the fraction deprotonated at pH 7 decreased, which for NaCl uniquely occurred upon increasing the salt concentration, as the resulting higher density of less acidic sites facilitated more repulsive nearest neighbor interactions and a lower n value.
5. Conclusion We have explored the influence of electrolyte concentration for various alkali halides at the fused silica/aqueous interface. The influence of the ion can be quantified from the effective acid dissociation constants pKaeff and the change in distribution of more and less acidic sites manifested in the change in the fraction of the surface that is deprotonated at pH 7. For the alkali chlorides explored, we observed that the pKaeff of both types of sites increased with increasing cation size/polarizability. We attribute the lowest pKaeff of NaCl to the strong interactions between the sodium and the siloxide sites, suggesting that the SiO− sites have similar water affinities to that of Na+, which is reasonable given their similarity in polarizability and ionic radius. The order of LiCl < KCl < CsCl with respect to the remaining pKaeff values is attributed to the decreasing value of Kd corresponding to the dissociation of the cation from the neutral surface sites. From this we infer that Cs+ associates most strongly with the neutral sites followed by K+ and Li+. To fully understand the origin of these interactions, the hydrogen-bonding structure of the water must be ascertained and is the subject of current sum frequency generation studies in our group. We note that these specific cation effects are relatively weak at −0.1 M, which indicates that they may not be as relevant over the range of salinities found in the environment. In contrast, the specific anion effects were very insensitive to ion concentration particularly for the NaI 9
Md Shafiul Azam et al
J. Phys.: Condens. Matter 26 (2014) 244107
[15] Fisk J D, O’Reilly J P and Shaw A M 2006 J. Phys. Chem. B 110 15039 [16] Sulpizi M, Gaigeot M-P and Sprik M 2012 J. Chem. Theory Comput. 8 1037 [17] Leung K, Nielsen I M B and Criscenti L J 2009 J. Am. Chem. Soc. 131 18358 [18] Azam M S, Weeraman C N and Gibbs-Davis J M 2012 J. Phys. Chem. Lett. 3 1269 [19] Azam M S, Weeraman C N and Gibbs-Davis J M 2013 J. Phys. Chem. C 117 8840 [20] Abendroth R P 1970 J. Colloid Interface Sci. 34 591 [21] Campen R K, Pymer A K, Nihonyanagi S and Borguet E 2010 J. Phys. Chem. C 114 8465 [22] Eftekhari-Bafrooei A and Borguet E 2009 J. Am. Chem. Soc. 131 12034 [23] Kitamura A, Fujiwara K, Yamamoto T, Nishikawa S and Moriyama H 1999 J. Nucl. Sci. Technol. 36 1167 [24] Azam M S and Gibbs-Davis J M 2013 Anal. Chem. 85 8031 [25] Gibbs-Davis J M, Kruk J J, Konek C T, Scheidt K A and Geiger F M 2008 J. Am. Chem. Soc. 130 15444 [26] Shen Y R 1984 The Principles of Nonlinear Optics (New York: Wiley-Interscience) [27] Hayes P L, Malin J N, Jordan D S and Geiger F M 2010 Chem. Phys. Lett. 499 183 [28] Lützenkirchen J 1999 J. Colloid Interface Sci. 217 8 [29] Yan E C Y, Liu Y and Eisenthal K B 1998 J. Phys. Chem. B 102 6331 [30] Jena K C, Covert P A and Hore D K 2011 J. Phys. Chem. Lett. 2 1056 [31] Davis J A, James R O and Leckie J O 1978 J. Colloid Interface Sci. 63 480 [32] Behrens S H and Grier D G 2001 J. Chem. Phys. 115 6716 [33] Dishon M, Zohar O and Sivan U 2009 Langmuir 25 2831 [34] Morel F M M and Hering J G 1993 Principles and Applications of Aquatic Chemistry (New York: Wiley) [35] Du Q, Freysz E and Shen Y R 1994 Phys. Rev. Lett. 72 238 [36] Ostroverkhov V, Waychunas G and Shen Y 2005 Phys. Rev. Lett. 94 046102 [37] Yang Z, Li Q and Chou K C 2009 J. Phys. Chem. C 113 8201 [38] Jena K C and Hore D K 2009 J. Phys. Chem. C 113 15364 [39] Dewan S, Yeganeh M S and Borguet E 2013 J. Phys. Chem. Lett. 4 1977 [40] Brown M A, Huthwelker T, Redondo A B, Janousch M, Faubel M, Arrell C A, Scarongella M, Chergui M and van Bokhoven J A 2012 J. Phys. Chem. Lett. 3 231 [41] Urry D W 2006 What Sustains Life? Consilient Mechanisms for Protein-Based Machines and Materials (New York: Springer)
electrolyte. As cooperativity is the predominant factor governing the surface acidity of these sites in the presence of iodide, we infer that this cooperative structure is so stable that it is able to form at 0.1 M concentrations. Moreover, the relative distribution of sites still favors the presence of the more acidic sites even at these lower iodide electrolyte concentrations, which supports that the iodide ion has a stronger affinity for the silica/aqueous interface than the chloride ion in the presence of sodium or potassium. Acknowledgments We acknowledge the Natural Sciences and Engineering Research Council of Canada and the Canada Foundation for Innovation for funding. References [1] Iler R K 1979 Chemistry of Silica—Solubility, Polymerization, Colloid and Surface Properties and Biochemistry (New York: Wiley) [2] Dove P M and Rimstidt J D 1994 Reviews in Mineralogy Series: The Silica Polymorphs vol 29 ed P Heaney et al (Chantilly, VA: Mineralogical Society of America) p 259 [3] Sahai N 2002 Environ. Sci. Technol 36 445 [4] Icenhower J P and Dove P M 2000 Geochim. Cosmochim. Acta 64 4193 [5] Dove P M 1999 Geochim. Cosmochim. Acta 63 3715 [6] Milonjic S K 1987 Colloids Surf. 23 301 [7] Dove P and Craven C 2005 Geochim. Cosmochim. Acta 69 4963 [8] Eisenthal K B 1996 Chem. Rev. 96 1343 [9] de Beer A, Campen R and Roke S 2010 Phys. Rev. B 82 235431 [10] Fisk J D, Batten R, Jones G, O’Reill J P and Shaw A M 2005 J. Phys. Chem. B 109 14475 [11] O’Reilly J P, Butts C P, I’Anson I A and Shaw A M 2005 J. Am. Chem. Soc. 127 1632 [12] Ong S, Zhao X and Eisenthal K B 1992 Chem. Phys. Lett. 191 327 [13] Zhao X, Ong S, Wang H and Eisenthal K B 1993 Chem. Phys. Lett. 214 203 [14] Petersen P B and Saykally R J 2006 Annu. Rev. Phys. Chem. 57 333
10
