Thermodynamics of Proton Binding of Halloysite ... - ACS Publications

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Thermodynamics of Proton Binding of Halloysite Nanotubes Clemente Bretti,† Salvatore Cataldo,‡ Antonio Gianguzza,‡ Gabriele Lando,† Giuseppe Lazzara,‡ Alberto Pettignano,*,‡ and Silvio Sammartano† †

Dipartimento di Scienze Chimiche, Biologiche, Farmaceutiche ed Ambientali, Università degli Studi di Messina, Viale Ferdinando Stagno d’Alcontres 31, Villaggio Sant’Agata, I-98166 Messina, Italy ‡ Dipartimento di Fisica e Chimica, Università di Palermo, Viale delle Scienze, I-90128 Palermo, Italy S Supporting Information *

ABSTRACT: In this paper, new information on physical and chemical properties of the widely used nanostructured Halloysite mineral are reported. Given that the Halloysite has a tubular structure formed by a variable number of wrapped layers containing Si−OH and Al−OH groups, their proton binding affinity was measured at different ionic strengths and ionic media by means of potentiometric measurements in heterogeneous phase. One protonation constant for the Si−OH groups and two for the Al−OH groups were determined. The protonation constant values increase with increasing of the ionic strength in all the ionic media. This suggests that the presence of a background electrolyte stabilizes the protonated species through the formation of weak complexes between ions of the supporting electrolytes and the protonated species. Ten weak species were determined with different stoichiometry. It was shown that the interactions do not depend on the nature of the supporting electrolytes but on the charge. The surface charge of Halloysite was estimated by ζ potential measurements as a function of pH, and the values obtained are consistent with the nanotubes ionization predicted by using the protonation constants for the Si−OH and Al−OH groups. The total solubility of the Halloysite nanotubes, was also determined in NaCl aqueous solution. These measurements showed that the solubility slightly increases with increasing ionic strength and contact time between Halloysite and NaCl solution. Goodness-of-fit (GOF) criteria were used to test the application of these models with good results. The obtained results confirm that the behavior of Halloysite in water is strictly correlated to the experimental conditions of the aqueous suspension (e.g., pH, ionic strength, and ionic media). The thermodynamic data here reported are of main importance in the several applications where is exploited the charge separation between the inner and outer surfaces of this nanotubular material. a length of 1 ± 0.5 μm and inner and outer diameters of 10−15 and 50−70 nm, respectively.1,5 In each layer, the SiOH and the AlOH groups are disposed on the external and the internal surfaces, respectively. As a consequence, the chemistry of the lumen and of the outer surfaces of HNTs is completely different. In particular, in each nanotube, the inner surface is positively charged because of protonation of AlOH groups at low pH. On the contrary, at pH higher than 2, the gradual deprotonation of SiOH groups causes an excess of negative charges on the outer surface. Different studies6−9 carried out on the HNTs cytotoxicity show that the HNTs are nontoxic for humans and animals and are biocompatible and environmental friendly. The nontoxicity together with the charge separation and the particular hollow tubular structure of HNTs make this clay mineral very useful for different purposes. For example, HNTs

1. INTRODUCTION In the past decade, nanomaterials have assumed an increasing importance because of their particular characteristics mainly because of the nanometer-scale dimensions that confer to them a large surface/volume ratio. Among nanostructured compounds, natural clay minerals are attracting a growing scientific interest because they are cheap, easily available, and ecofriendly.1 One of the most promising nanostructured and naturally occurring clay minerals is Halloysite.2 It is an abundant and cheap clay, present in large deposits worldwide such as those in New Zealand, France, Belgium, China, and the United States.3,4 Halloysite (Al2Si2O5(OH)4·2H2O) is similar to kaolin, and its morphology depends on the extraction site. The principal structures are spheroidal, tubular, or platy and are due to particular crystallization conditions whose consequence is a deformation caused by a lattice mismatch.2,3 Among the different morphologies, the tubular is the most common and abundant. Typically, Halloysite nanotubes (HNTs) are formed by 15−20 aluminosilicate layers, having © 2016 American Chemical Society

Received: February 2, 2016 Revised: March 22, 2016 Published: March 22, 2016 7849

DOI: 10.1021/acs.jpcc.6b01127 J. Phys. Chem. C 2016, 120, 7849−7859

Article

The Journal of Physical Chemistry C

changing in the range 1−10 was also evaluated by ζ potential measurements of HNTs suspensions in NaCl aqueous solution at I = 0.1 mol dm−3, and correlations between the ζ potential trend and the distribution of protonated and deprotonated species of HNTs were found.

are used as support for different polymerization reactions, for the immobilization of catalysts (enzymes, metal ions, or metal ion complexes), as adsorbent of inorganic and organic contaminants and pollutants in the environmental remediation procedures, as support for drug delivery, as nanocontainers for anticorrosion coating, and so on.1,4,7,10−21 Several reviews have been published on the different applications of HNTs and all the authors agree that although the tubular structure and the high surface/volume ratio are fundamental characteristics for the versatility of this clay mineral the most important role is played by the charge separation between the inner and the outer surfaces.2,4,5,12 This peculiar difference in the chemistry of the inner and outer surface is strategic to tailor the adsorption of charged species in a selective way.22,23 In spite of the great number of investigations reported in the literature, there is a lack of reports on the acid−base equilibria that bring such a peculiar charge separation between HNTs surfaces. The negative charges on the outer surface as well as the positive charges in the lumen are strictly dependent on the acid−base properties of HNTs. Although the protonation of HNTs is of fundamental importance for the comprehension of their behavior in aqueous solution, to our knowledge no quantitative data are reported in literature. This is probably due to the difficulties in carrying out the experiments and processing experimental data. In fact, HNTs have a great number of functional groups with a charge separation, in a wide pH range, between the inner and the outer surfaces. The classical approach used for the study of protonation of low-molecular-weight molecules cannot be used, and a model that takes into account all these features must be applied. The present study is devoted to the study of HNTs (Dragonite) behavior in aqueous solution in different ionic media and ionic strengths (0.025 ≤ I ≤ 0.750, mol dm−3) and in a wide pH range (2−11). Dragonite is a type of Halloysite extracted in North America, and its nano particles are thin and short tubes (length = 50−1500 nm; inner diameter = 5−30 nm; outer diameter = 20−150 nm; wall thickness = 5−50 nm).3 The protonation/deprotonation of SiOH and AlOH groups was evaluated by ISE-H+ potentiometric titrations at T = 298.15 K. The experimental data were analyzed by assuming the existence of two HNTs units, namely, HNT1 and HNT2. The HNT1 units contain AlOH groups present in the inner surface, whereas HNT2 units contain the SiOH groups of the outer surface of the nanotubes. Their protonation equilibria falls in two different pH ranges, and the ISE-H+ potentiometric data were processed by using the same models successfully used in previous works24−26 for carboxylic and phenol groups of humic and fulvic substances. The effect of the ions of the medium on the behavior of hydrated Halloysite was already studied by different authors, and some papers have been published on salts intercalation and on the formation of HNTs−salt complexes in different ionic media.27 Taking into account the literature findings, here the dependence of protonation constants on ionic strength and on ionic medium was also studied, and important evaluations were made on the influence of the ions of the medium on the acid− base properties of inner and outer groups, respectively. The solubility and the kinetics of HNTs dissolution in NaCl aqueous solution, in the ionic strength range 0.05−0.50 mol dm−3, was also evaluated on the basis of Si and Al concentrations in solution after different contact times. Moreover, the change of superficial charge with the pH

2. MATERIALS AND METHODS 2.1. Chemicals. HNTs were a commercial product (Sigma, lot MKBQ8631 V) and were used after washing with ultrapure water and drying in oven at T = 383.15 K. Thermogravimetric analysis, microscopy, and diffusion in water studies of HNTs are reported by Cavallaro et al.28 Scanning electron microscopy (SEM) image of HNTs obtained by using an ESEM FEI QUANTA 200F after gold coating in argon is reported in Figure 1.

Figure 1. SEM image of Halloysite; bar is 1 μm.

Sodium chloride, potassium nitrate, and tetraethylammonium chloride (Et4NCl) were supplied by Fluka, purity ≥98%, and were used without further purification. Hydrochloric acid and sodium or tetraethylammonium hydroxide solutions were prepared by diluting concentrated Fluka ampules and standardized against sodium carbonate and potassium hydrogen phthalate, respectively, previously dried in an oven at T = 383.15 K. Aluminum and silicon standard solutions for ICP measurements were purchased from Sigma-Aldrich. All the solutions were prepared using freshly prepared and CO2-free ultra pure water (ρ ≈ 18 MΩ cm), and grade A glassware was always employed. 2.2. Apparatus and Procedure for Potentiometric Measurements. The investigations on the acid−base properties of HNTs were carried out by potentiometry at T = (298.15 ± 0.1) K. Potentiometric titrations were carried out using an 809 Metrohm Titrando apparatus equipped with a combined Orion glass electrode Ross type 8102. The apparatus was connected to a PC, and automatic titrations were carried out using the Metrohm TiAMO 1.2 software to check for emf stability and to control titrant delivery and data acquisition. The estimated accuracy was ±0.20 mV and ±0.005 cm3 for emf and titrant volume readings, respectively. All titrations were carried out under magnetic stirring, and presaturated N2 was bubbled through the purified solution in order to exclude O2 and CO2 inside. The HNTs aqueous suspensions were prepared by 7850


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The Journal of Physical Chemistry C weighing 0.3−1.2 g of HNTs in 100 cm3 of aqueous solution containing the right amount of ionic medium used (NaCl, KNO3, or Et4NCl) to adjust the ionic strength at the preestablished value in the range 0.025 ≤ I ≤ 0.75 (mol dm−3). Hydrochloric or nitric acid were added to each solution in order to set the pH of suspensions at ∼2. A volume of 25 cm3 of each suspension was titrated with a standard CO2-free hydroxide solution containing in turn the same cation present in the ionic medium. For each experiment, independent titrations of strong acid with standard hydroxide solutions were carried out under the same experimental conditions of the system investigated to determine the formal electrode potential E°. 2.3. Apparatus and Procedure for Solubility Measurements. Solubility measurements were carried out as reported elsewhere.29−31 Briefly, saturated solutions of HNTs were prepared in water and in NaCl(aq) at different concentrations. Mixtures were shaken for different interval times to evaluate kinetic effects. The solid was removed by filtration over 0.45 μm Whatman filters. The total solubility of HNTs was determined by ICP-OES (PerkinElmer Model Optima 2100, equipped with an autosampler model AS-90) in terms of aluminum and silicon concentrations in solution. The measurements were repeated three times at two different λ, and the mean values were collected. 2.4. Apparatus and Procedure for ζ Potential Measurements. ζ potential measurements were carried out in a ZETASIZER NANO ZS90 (Malvern Instruments) equipped with a He−Ne laser (λ = 632.8 nm). Disposable folded capillary cell was used; temperature was controlled within ±0.1 °C. The electrophoretic mobility is obtained by measuring the velocity of the particles in an electric field by using laser Doppler velocimetry and phase analysis light scattering. The scattering angle of 12.8 was set. The Smoluchowski approximation was used to calculate the ζ potential because it is considered a good approximation for particles larger than 0.2 μm in water containing at least 0.001 mol dm−3 of electrolytes. All experiments were carried out in triplicate.

make very difficult the study of acid−base properties of this material. Moreover, the solubility of HNTs in water is very low, and the potentiometric study must be carried out on a suspension continuously stirred in order to create a colloidallike solution. For these reasons, the analysis of HNTs protonation equilibria cannot be carried out on a clear solution and by following the classical approach used for low-molecularweight molecules with a limited number of functional groups. Another element to be evaluated for the calculation of protonation constants of HNTs is the presence of two different hydroxyl groups, SiOH and AlOH, whose acid−base equilibria fall in different pH regions. A first approximation consists in the attribution of the AlOH and the SiOH groups at two different units named HNT1 (inner surface of HNTs) and HNT2 (outer surface of HNTs), respectively. The approach used for the calculation of their protonation constants was different. The protonation constants of the AlOH groups of HNT1 unit K = aH(HNT) /(aHa HNT1) 1

(1)

referred to the equilibria (z + 1) H+ + HNT1z ⇄ H(HNT) 1

(1a)

were calculated by using different models for the dependence of K on the dissociation degree (α): α = [HNT]/([H(HNT)] + [HNT]) 1 1 1 = [HNT]/[HNT] 1 1 T

(1b)

In fact, by increasing α the charge of HNT1 changes, and consequently, the result is a variation of K value. These models were successfully employed in a previous paper for the study of the acid−base properties of several natural and synthetic polyelectrolytes that have in common with HNTs the numerous functional groups on their surface.34 The models used were as follows: (i) the modified Henderson−Hasselbalch equation log K = log K n − (n − 1) log[(1 − α)/α]

(2)

(ii) the three-parameters equation, based on the Guggenheim 0th approximation,35 proposed by Högfeldt

3. DATA ANALYSIS 3.1. Calculations. The nonlinear least-squares computer program ESAB2M was used for the refinement of all the parameters of the acid−base titration, such as the formal electrode potential (E°), the ionic product of water (Kw), the acidic junction potential (Ej = ja[H+]), and the analytical concentration of reagents. The BSTAC and STACO computer programs were used for the refinement of protonation constants. The LIANA computer program was used to fit different linear and nonlinear functions, in particular, to calculate the protonation constants of the HNTs with the different proposed models for the dependence of protonation constants on the dissociation degree (α). The ES2WC program32 was used to study the dependence of log K on ionic medium and ionic strength and to calculate the formation constants of weak complexes formed by HNTs with the ions of ionic media. The ES4ECI program was used to draw speciation diagrams and to calculate species formation percentages. Details on the computer programs used are given in De Stefano et al.33 3.2. Models Used for the Calculation of Protonation Constants of Halloysite Clay Nanotubes. The numerous hydroxyl groups of Halloysite together with their particular disposition in the inner and outer surfaces of the nanotubes

log K = α 2 log K1 + (1 − α)2 log K 0 + 2α(1 − α) log K m (3)

(iii) the linear model log K = α log K1 + (1 − α) log K 0

(4)

and (iv) the diprotic-like model, which considers the HNT1 unit (AlOH groups) of HNTs a difunctional unit whose acid− base properties can be defined by two protonation constants (K1 and β2) independent of α. The diprotic-like model can be used, with a negligible loss of precision, when in the range 0.1 ≤ α ≤ 0.9 the variation of protonation constant of the functional groups considered is not very large (∼2.5 log units). Moreover, the charge of the HNT1 unit of Halloysite is formally equal to 0, but the effective charge of the Halloysite clay is given by the contribution of the groups of HNT1 and HNT2 units placed in the inner and outer surfaces, respectively. In the case of diprotic-like model, the protonation constants of HNT1 are given by K i = [HiHNT1z + i]/([Hi − 1HNT1z + i − 1][H+]) 7851


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Table 1. Solubility of HNTs in NaCl Aqueous Solution at Different Ionic Strengths and Residence Times at T = 298.15 K Together with the [Si]/[Al] Ratio time (h)

[NaCl] (mol dm−3)

pH

[Al] (mmol dm−3, ± 0.08%)

log ST

[Si] (mmol dm−3, ± 0.05%)

log ST

[Si]/[Al]

24 24 24 24 24 168 168 168 168 168 282 282 282 282 282 619 619 619 619 619 1800 1800 1800 1800 1800

0.0000 0.0525 0.1036 0.2546 0.5003 0.0000 0.0502 0.1000 0.2508 0.5013 0.0000 0.0502 0.1000 0.2508 0.5013 0.0000 0.0502 0.1000 0.2508 0.5013 0.0000 0.0502 0.1000 0.2508 0.5013

6.78 6.05 5.22 5.92 5.18 6.70 6.12 6.27 6.50 6.41 6.83 5.81 6.36 6.00 6.05 7.80 5.83 6.07 5.40 6.55 7.50 6.26 6.08 5.66 5.56

0.0215 0.0193 0.0222 0.0300 0.0419 0.0208 0.0211 0.0193 0.0230 0.0334 0.0271 0.0178 0.0219 0.0245 0.0300 0.0315 0.0222 0.0167 0.0237 0.0363 0.0259 0.0163 0.0219 0.0300 0.0612

−4.668 −4.715 −4.653 −4.523 −4.378 −4.683 −4.675 −4.715 −4.639 −4.477 −4.568 −4.750 −4.660 −4.611 −4.523 −4.502 −4.653 −4.778 −4.625 −4.440 −4.586 −4.788 −4.660 −4.523 −4.214

0.080 0.072 0.077 0.073 0.064 0.067 0.058 0.070 0.070 0.129 0.069 0.059 0.073 0.094 0.082 0.067 0.061 0.064 0.061 0.069 0.090 0.089 0.107 0.098 0.192

−4.094 −4.145 −4.116 −4.139 −4.193 −4.172 −4.236 −4.156 −4.156 −3.891 −4.161 −4.226 −4.135 −4.027 −4.087 −4.174 −4.218 −4.196 −4.213 −4.159 −4.047 −4.051 −3.970 −4.011 −3.717

3.74 3.71 3.44 2.42 1.53 3.24 2.75 3.62 3.04 3.85 2.55 3.34 3.35 3.84 2.73 2.12 2.72 3.82 2.58 1.91 3.46 5.46 4.90 3.25 3.14

log K = log K 0 − z*(I 0.5/(2 + 3I 0.5) − 0.1I1.5) + CI

and consequently, the two protonation constants are expressed as K1 = aH(HNT)1 /(aHa HNT1)

(5)

β2 = a H2(HNT)1 /((aH)2 a HNT1)

(6)

(8)

where K is the protonation constant, K0 is the relative value at infinite dilution, and C is an empirical parameter for the dependence on I.

An exhaustive explanation of the models used is reported in Cigala et al.34 The HNT2 unit contains the SiOH groups present in the outer surface of the Halloysite nanotubes, and with a good approximation in the acidic pH range (2 ≤ pH ≤ 6), only the acid−base equilibria of these groups are present. The negligible variation of protonation constants of the SiOH groups of HNT2 unit with α suggests that their acid−base properties are well-defined by only one protonation constant: K1 = aH(HNT2)/(aHa HNT2)

4. RESULTS AND DISCUSSION 4.1. Solubility of HNTs in NaCl Aqueous Solution. The solubility of HNTs in NaCl aqueous solution at different ionic strengths 0 ≤ I ≤ 0.5 (mol dm−3) at different contact times (24−1800 h) and T = 298.15 K was evaluated by ICP-OES measurements of Si and Al concentrations present in the solutions in contact with 0.3−0.5 g of HNTs. The results are reported in Table 1 together with the [Si]/[Al] ratio. The average [Si]/[Al] ratio of 3.2 ± 0.8 is in line with the average [SiOH]/[AlOH] ratio of 4.0 ± 0.3 calculated on the basis of concentrations refined by STACO and BSTAC computer programs during the potentiometric data processing in the different ionic media and ionic strengths. Moreover, considering the literature data regarding the mean length and the mean internal and external radii,1,3,5 an average external area/internal area ∼4 can be calculated. It can be supposed that there exists a correlation between the values of the three parameters. Moreover, for obvious reasons, the external layer is more available for participating in solution equilibria; therefore, it is reasonable to obtain a greater concentration of these sites than of the Al ones in solution. Dixon’s test on the abovementioned ratio confirmed that no outliers are present in the set, and statistical analysis of the residues was carried out with some usual goodness-of-fit tests (GOF, see ref 41), which confirmed that is acceptable to determine the average value of the whole data set instead of considering various data sets

(7)

3.3. Dependence on Ionic Strength of Protonation Constants. The dependence of protonation constants of the AlOH and SiOH groups on ionic strength was studied considering the formation of weak complexes of these groups present in the different units (HNT1 and HNT2, respectively) with the ions of the supporting electrolytes. This model is one of the chemical approaches to the variation of the activity coefficients with ionic strength. The physical and the hybrid approaches are widely discussed by Pytkowicz.36,37 The theoretical basis of this approach are discussed in some reference works (e.g., refs 38−40). For convenience, the formal charge of the HNT1 groups was set to z = −1, whereas that of the HNT2 to z = 0. The equation used for the modeling of the ionic strength dependence of the protonation and weak complex formation constants is 7852


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The Journal of Physical Chemistry C Table 2. Refined Parameters of eq 9 Analysing the Experimental Data Reported in Table 1

a

sites

P1

aluminum silicon

−4.69 ± 0.01 −4.18 ± 0.01

P2

P3

−0.0010 ± 0.0006 0.0013 ± 0.0004

b

0.36 ± 0.05 0.08 ± 0.07

b

m.d.a

P4 0.007 ± 0.002 0.007 ± 0.002

b

b

0.059 0.057

Statistical parameters of the fit. b±95% confidence interval (C.I.).

Table 3. Protonation Constants of AlOH groups of HNT1 Units of HNTs Calculated with Four Different Models in NaCl, KNO3, and Et4NCl Media at Different Ionic Strengths and at T = 298.15 K Together with Standard Deviation on the Fit protonation models

Högfeldt (eq 3)

modified Henderson−Hasselbalch (eq 2)

linear (eq 4)

diprotic-like

medium

I (mol dm−3)

log K1

log K0

log Km

σa

log K1

log K0

σa

log Kn

n

σa

log K1(HNT1)

log β2(HNT1)

σa

NaCl NaCl NaCl NaCl NaCl NaCl KNO3 KNO3 KNO3 KNO3 KNO3 Et4NCl Et4NCl Et4NCl Et4NCl Et4NCl Et4NCl

0.041 0.060 0.098 0.224 0.437 0.626 0.067 0.104 0.228 0.428 0.635 0.043 0.059 0.11 0.239 0.459 0.664

9.63 9.78 9.810 10.04 10.29 10.36 9.50 10.01 9.86 10.08 10.14 9.82 9.62 9.55 9.58 9.84 10.11

6.92 7.33 7.774 7.59 7.570 7.64 7.12 7.424 7.352 7.669 7.455 7.57 7.84 7.96 7.81 7.85 7.83

9.06 9.29 9.429 9.68 9.90 9.87 9.21 9.837 9.81 10.08 10.04 9.18 9.23 9.25 9.33 9.78 9.97

0.042 0.037 0.012 0.021 0.015 0.023 0.015 0.013 0.006 0.010 0.010 0.025 0.008 0.077 0.028 0.027 0.014

10.00 10.05 10.092 10.419 10.884 11.18 10.00 10.64 10.66 11.00 11.12 10.02 9.82 9.81 9.85 10.21 10.50

7.21 7.70 8.002 7.926 7.832 7.78 7.403 7.77 7.680 7.939 7.715 7.76 8.04 8.09 8.05 8.22 8.23

0.034 0.052 0.030 0.035 0.035 0.019 0.032 0.034 0.022 0.024 0.027 0.027 0.029 0.054 0.033 0.042 0.032

8.617 8.875 9.058 9.185 9.375 9.411 8.715 9.220 9.178 9.461 9.403 8.89 8.93 8.97 8.96 9.23 9.38

2.38 2.23 2.023 2.251 2.487 2.52 2.31 2.45 2.50 2.495 2.63 2.09 1.86 1.82 1.89 1.99 2.13

0.047 0.058 0.026 0.039 0.032 0.018 0.034 0.030 0.014 0.016 0.005 0.042 0.038 0.038 0.010 0.050 0.037

9.78 9.89 9.99 10.23 10.38 10.30 9.67 9.92 10.16 10.38 10.33 9.81 9.83 9.82 9.87 10.12 10.32

17.24 17.60 18.00 18.19 18.44 18.39 17.23 17.69 18.05 18.56 18.38 17.64 17.8 17.91 17.84 18.3 18.57

0.037 0.030 0.032 0.030 0.040 0.032 0.034 0.036 0.033 0.037 0.037 0.039 0.028 0.033 0.029 0.035 0.034

a

Standard deviation on the fit.

σ2(e) = 0.006, and no outliers were found with the Dixon’s test.41 4.2. Protonation Constants of HNTs in NaCl, KNO3, and Et4NCl Aqueous Media. The protonation of Halloysite nanotubes in different ionic media and ionic strengths was investigated by ISE-H+ potentiometry. The concentration of AlOH and SiOH groups (mmol g−1) of HNTs was checked by direct potentiometric titrations of HNTs suspensions. The measurements were carried out with different amounts of HNTs in the range 5−12 g dm−3 and after different contact times. Independent of the amount of HNTs, the concentration of AlOH and SiOH groups gradually changes and stabilizes after 3 days from the preparation of each suspension (CSiOH = 0.80 ± 0.05 mmol g−1; CAlOH = 0.20 ± 0.05 mmol g−1). As pointed out before, two different units named HNT1 and HNT2 containing the AlOH and SiOH groups, respectively, were considered in the study of acid−base properties of HNTs. The SiOH groups of HNT2 units are protonated at pH 2 (AlOH2+), and their deprotonation occurs in the pH range 2− 6. The acid−base equilibria of AlOH groups of HNT1 units are in a different pH range (6−11). For this reason, the potentiometric data of the acidic and of neutral−alkaline pH ranges were processed in two different steps with STACO, BSTAC, and LIANA computer programs. During the data processing, the AlOH and SiOH concentrations were always refined together with the protonation constants calculated using different models (section 3.2). The same data processing was successfully used

divided as a function of the contact time or the ionic strength. Looking at the table, we note that the measurement of the pH is only indicative in the case of those measurements without NaCl because of the absence of any background electrolyte. The solubility data listed in the table were divided into two sets (Al and Si) and were fitted against ionic strength (I) and time (h/24) using eq 9: log ST = P1 + P2 h/24 + P3 I + P4 Ih/24

(9)

The whole fit has a standard deviation of σ = 0.039 and a mean deviation of m.d. = 0.058; therefore, it must be regarded as a tentative analysis. However, some interesting information may be gathered: although the total solubility in pure water of the Al sites slightly decreases with increasing contact time, that of the Si sites increases. No information about the possible reasons of this behavior are available, though one may suppose a time-dependent organization of HNTs going in this direction. On the contrary, the solubility of the Al sites increases more quickly than that of the Si ones with increasing ionic strength. This can be correlated to the formation of the weak species between the variously protonated species of HNT1 and HNT2 and the ions of the supporting electrolyte (see below). The mixed parameter P4 is very similar for both the Al and Si sites. The whole sets of the refined parameters of eq 9 are reported in Table 2, together with the m.d. of the fits. GOF criteria41 are skewness γ1(e) = 0.6, kurtosis γ2(e) = 0.7, arithmetic mean of residuals e ̅ = 0.0018, absolute mean of residuals |e|̅ = 0.058, standard deviation of residuals σ1(e) = 0.011, and variance 7853


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degree α, and the values calculated at the different experimental conditions are reported in Table 5 together with the standard deviation on the fits. As an example, the distribution diagrams of HNT1 and HNT2 species in NaCl at two ionic strengths (0.025 and 0.75 mol dm−3) are reported in Figure 3. The diagrams were drawn considering a concentration of HNT1 (dashed lines) and HNT2 (continuous line) units of 1 and 4 mmol dm−3, respectively. The concentrations of the two units used to simulate the distribution of the HNTs species in the pH range 2−11 were established on the basis of the following considerations: (i) The mean concentration ratio of HNT2 and HNT1 units obtained by refining their concentrations during the processing of potentiometric titration data was CHNT2/CHNT1 = 4.0 ± 0.3. (ii) The mean outer surface/inner surface ratio of HNTs calculated on the basis of the mean length and of the mean diameter of HNTs reported in literature1,5 is ∼4. (iii) The mean value [Si]/ [Al] = 3.2 ± 0.8 obtained by solubility measurements is in line with the data reported in the previous two points. Figure 3 shows that at pH 2 the two units are fully protonated. Starting from this pH, the SiOH groups of HNT2 units (continuous lines) gradually deprotonate, and at pH ∼6.5, they are totally deprotonated with the consequence that the outer surface of nanotubes reaches the maximum negative charge. At the same time, the inner surface of HNTs has the highest positive charge at pH 2 when the AlOH groups of HNT1 units are in the protonated form (AlOH2+). Starting from pH ∼6, the H2(HNT1) species gradually deprotonates to H(HNT1) and HNT1 species. The effect of ionic strength on the acid−base properties of Halloysite is evident looking at the distribution of the HNTs species at I = 0.025 mol dm−3 (black lines) and at I = 0.75 mol dm−3 (gray lines) reported in Figure 3. For example, at pH 8, the formation percentage of H2(HNT1) species is 22.1 and 55.0% at I = 0.025 and 0.75 mol dm−3, respectively. The same diagrams can be drawn at every ionic strength inside the considered range and for each ionic medium investigated in order to correlate the charge values of the internal and external surfaces of HNTs (calculable at each pH value as the sum of the products of the charge of each species for its formation percentage) with the experimental conditions of the aqueous suspension. 4.3. Formation of Weak Complexes. Analyzing the ionic strength dependence of the protonation constants, it seems evident that the ions of the supporting electrolytes play an important role in these reactions. In fact, when both the silicate and aluminate groups undergo protonation or deprotonation reactions, the formation of a charged species is involved, and its interaction with inorganic ions present in solution, such as Na+ or Cl−, may stabilize the species by the formation of weak complexes. For example, the Si−O− species may form weak complexes with Na+, K+, and even Et4N+ ions. The formation of such interaction has been proved for many ligand classes, and relevant reviews have been published.39,42 In this case, because of the tubular nature of the Halloysite, probably the formation of the weak complexes allows for structure stabilization. This is emphasized by the fact that increasing the ionic strength, there is an increase in the values of either the protonation constants (Table 4) or total solubility (Table 1). Generally, the strength of the weak species can be estimated by (i) comparing the protonation constants determined in two different media, one of which considered as not interacting (e.g., NaCl vs Et4NI,via

in a previous work with potentiometric titrations of carboxylic and phenolic groups of humic and fulvic acids.24−26 The protonation constants of AlOH groups of HNT1 unit in the different ionic media, and ionic strengths together with the standard deviation on the fit of the four above-discussed models are reported in Table 3. As can be seen, the four models previously used for the study of acid−base properties of polyelectrolytes also give good results when adapted to the HNTs with low standard deviation on the fits of experimental data. In fact, the four models proposed have comparable modeling ability and give satisfactory results for the interpretation of the experimental titration curves. In other words, considering the results obtained for the measurements in NaCl medium, the mean values of the standard deviation of the fits (in log K units) are 0.025 (Högfeldt, eq 3), 0.034 (linear, eq 4), 0.037 (modified Henderson−Hasselbalch eq 2), and 0.034 (diprotic-like). Similar considerations can also be done for the results obtained in KNO3 and in Et4NCl. The dependence of the log K1(HNT1) (according to the diprotic-like model) of the HNT1 sites as a function of ionic strength is shown in Figure 2, where it can be observed that, in

Figure 2. Dependence of the first protonation constant of the HNT1 sites (according to the diprotic-like model, Table 3) on ionic strength in NaCl (□), and Et4NCl (○) at T = 298.15 K.

the considered ionic media, the presence of a background electrolyte stabilizes the protonated species. For this reason, the ionic strength dependence of the protonation constants was analyzed in terms of formation of weak complexes. In this graph, the data points are connected by segmented lines, and the complete data analysis is given in section 4.3, which is devoted to the formation of weak complexes. The average values of protonation constants according to the diprotic-like model (log β1/2), to the Högfeldt three-parameter equation [(log K1 + log K0 + log Km)/3], and to the linear model [(log K1 + log K0)/2] are reported in Table 4. The calculation of these parameters is very useful because they are directly comparable and gives an idea of the accordance between the results obtained with the four models. Considering the experimental errors, the mean protonation constants are in good accordance with each other and with the log Kn values calculated with the modified Henderson−Hasselbalch model. It confirms that even though each model is based on different assumptions and approximations each of them can be used to study the protonation of HNTs. The acid−base properties of silanol groups are well-defined by a single protonation constant independent of dissociation 7854


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Table 4. Mean Log K Values of AlOH Groups of HNT1 Units of HNTs Calculated with Four Different Models in NaCl, KNO3, and Et4NCl Media at Different Ionic Strengths and at T = 298.15 K Together with Standard Deviation on the Fit protonation models Högfeldt

a

modified Henderson− Hasselbalch

linear

I (mol dm−3)

log Ka

σb

log Kc

0.041 0.060 0.098 0.224 0.437 0.626

8.53 8.8 9.0 9.1 9.25 9.14

0.042 0.037 0.012 0.021 0.015 0.023

8.61 8.88 9.047 9.173 9.36 9.48

0.067 0.104 0.228 0.428 0.635

8.61 9.09 9.0 9.28 9.21

0.015 0.013 0.006 0.010 0.010

8.70 9.20 9.17 9.47 9.42

0.043 0.059 0.11 0.239 0.459 0.664

8.86 8.9 8.92 8.91 9.16 9.3

0.025 0.008 0.077 0.028 0.027 0.014

8.89 8.93 8.95 8.95 9.22 9.37

σb NaCl 0.034 0.052 0.030 0.035 0.035 0.019 KNO3 0.032 0.034 0.022 0.024 0.027 Et4NCl 0.027 0.029 0.054 0.033 0.042 0.032

diprotic-like

log Kn

σb

log β1/2

σb

8.617 8.875 9.058 9.185 9.375 9.411

0.047 0.058 0.026 0.039 0.032 0.018

8.62 8.8 9.0 9.1 9.22 9.19

0.037 0.030 0.032 0.030 0.040 0.032

8.715 9.220 9.178 9.461 9.403

0.034 0.030 0.014 0.016 0.005

8.62 8.85 9.03 9.28 9.19

0.034 0.036 0.033 0.037 0.037

8.89 8.93 8.97 8.96 9.23 9.38

0.042 0.038 0.038 0.010 0.050 0.037

8.82 8.9 8.95 8.92 9.15 9.29

0.039 0.028 0.033 0.029 0.035 0.034

(log K1 + log K0 + log Km)/3. bStandard deviation on the fit. c(log K1 + log K0)/2.

Table 5. Protonation Constants of SiOH Groups of HNT2 Units of HNTs in NaCl, KNO3, and Et4NCl Media at Different Ionic Strengths and at T = 298.15 K Together with Standard Deviation on the Fit I (mol dm−3) 0.041 0.060 0.098 0.224 0.437 0.626 0.067 0.104 0.228 0.428 0.635 0.043 0.059 0.11 0.239 0.459 0.664 a

log K NaCl 4.066 4.187 4.195 4.274 4.325 4.46 KNO3 4.157 4.230 4.304 4.355 4.386 Et4NCl 4.147 4.17 4.2 4.28 4.3 4.29

σa 0.037 0.030 0.032 0.030 0.040 0.032 0.034 0.036 0.033 0.037 0.037

Figure 3. Distribution diagrams of the HNT1 (dotted lines) and HNT2 (continuous lines) species vs pH in NaCl, at I = 0.025 mol dm−3 (black) and I = 0.75 mol dm−3 (light gray). Experimental conditions: CHNT1 = 1 mmol dm−3, CHNT2 = 4 mmol dm−3, T = 298.15 K.

0.039 0.028 0.033 0.029 0.035 0.034

for the dependence of the activity coefficient on ionic strength. According to this model (See refs 40, 32, and 43−46 and references therein.), the variation of the equilibrium constants with ionic strength is given in eq 8. In total, the stability of ten weak species was determined; however, the stoichiometry of these species can be reduced to four different reactions, namely, M(HNT2), MH(HNT2)+, H(HNT1)L, and H2(HNT1)L2 (M = Na+, K+, or Et4N+ and L = Cl− and NO3−). The values of the protonation and complexes formation constants at infinite dilution together with the values of the ionic strength dependence parameters of eq 8 are listed in Table 6. It is interesting that the stability of the Na(HNT2) species is log K = 0.09 at infinite dilution, whereas that of the

Standard deviation on the fit.

ΔpK method), (ii) considering a general model for the activity coefficients in not interacting media, and (iii) experimental data obtained in aqueous NaCl solution using Na+−ISE electrode. In this paper, the stability of the weak complex species was determined by means of ES2WC computer program32 using the well-known ΔpK method and considering a general model 7855


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Table 6. Ionic Strength Dependence Parameters of Protonation Constants Calculated with Diprotic-Like Model and of Formation Constants of Weak Complexes with the Ions of Medium According to Eq 8 of the HNTs in Different Ionic Media and at T = 298.15 K

H+ + (HNT2)− ⇄ H(HNT2) Na+ + (HNT2)− ⇄ Na(HNT2) K+ + (HNT2)− ⇄ K(HNT2) Et4N+ + (HNT2)− ⇄ Et4N(HNT2) Na+ + H+ + (HNT2)− ⇄ NaH(HNT2)+ K+ + H+ + (HNT2)− ⇄ KH(HNT2)+ Et4N+ + H+ + (HNT2)− ⇄ (Et4N)H(HNT2)+

C ± 0.01b

log K0

equilibrium SiOH Sites, z = −1a

4.23 ± 0.02b 0.09 ± 0.06 0.09 ± 0.04 −0.01 ± 0.04 4.66 ± 0.09 4.54 ± 0.07 4.62 ± 0.07

0.56 0.56 0.56 0.56 0.66 0.66 0.66

9.75 ± 0.08 17.47 ± 0.08 10.43 ± 0.14 19.08 ± 0.14 10.55 ± 0.13 19.02 ± 0.17

0.10 −0.26 0.66 1.32 0.66 1.32

AlOH Sites, z = 0c H+ + (HNT1) = H(HNT1)+ 2 H+ + (HNT1) ⇄ H2(HNT1)2+ H+ + (HNT1) + Cl− ⇄ H(HNT1)Cl 2 H+ + (HNT1) + 2 Cl− ⇄ H2(HNT1)Cl2 H+ + (HNT1) + NO3− ⇄ H(HNT1)NO3 2 H+ + (HNT1) + 2 NO3− ⇄ H2(HNT1) (NO3)2 a

σ fit = 0.03; m.d. fit = 0.02. b±95% C.I. cσ fit = 0.08; m.d. fit = 0.06.

H(HNT1)Cl species (according to the equilibrium H(HNT1)+ + Cl− ⇄ H(HNT1)Cl) is log K = 0.7 under the same conditions. This important difference in the stability of the species may help with interpreting the fact that the total solubility of the HNT1 sites increased more quickly than that of the HNT2 ones with increasing ionic strength (section 4.1). It is interesting that the stability of the species with the same stoichiometry is very similar, indicating that these coloumbic interactions are probably not specifically dependent on the nature of the ions, but only on the charge, favoring the solvation of the molecules. For example, the average value for the generic equilibrium M+ + (HNT2)− ⇄ M(HNT2) is log K = 0.06 ± 0.06, which is not statistically different with respect to the values of the specific Na+, K+, or Et4N+ species. On the external surface of the HNTs, many deprotonated Si−O− groups are available, and the presence of positively charged ions of the background electrolyte may reduce the repulsion among charges of same sign that are close to each other, independent of the structure and size by means of the formation of weak complexes. In Figure 4, the distribution of the protonated and weak complex species vs pH is showed at I = 0.1 mol dm−3, with the same HNTs concentrations of Figure 3. As can be seen, the presence of the four complex species is found in the entire pH range, with a cumulative formation percentage of ∼25%. The NaH(HNT2) species is present between pH ∼2.0 and ∼5.0, whereas the formation of the Na(HNT2) species starts at pH ∼3.5 and its percentage (∼10%) remains constant up to pH ∼11.0. The H(HNT1)Cl species reaches a maximum of formation of ∼20% at pH ∼9.0, whereas the H2(HNT1)Cl2 species represents ∼10% of the total HNT1 species between pH ∼2.0 and ∼7.0, disappearing at pH ∼8.0. This diagram demonstrates that even at low ionic strength in a correct speciation scheme the presence of the mentioned weak species cannot be neglected because their total formation percentage is quite high (∼25%) in the considered pH range. 4.4. ζ Potential of HNTs in NaCl Aqueous Solution. ζ potential experiments were carried out at two different ionic strengths as a function of pH to highlight the surface charge variation of Halloysite nanotubes. The obtained results are provided in Table 7. The decrease of the ζ potential reflects the

Figure 4. Distribution diagrams of the HNT1 (dashed lines) and HNT2 (continuous lines) protonated and weak complex species vs pH in NaCl, at I = 0.100 mol dm−3. Experimental conditions: CHNT1 = 1 mmol dm−3, CHNT2 = 4 mmol dm−3, T = 298.15 K.

Table 7. ζ Potential Values of HNTs in NaCl at I = 0.05 and 0.10 mol dm−3 I (mol dm−3)

pH

0.05 0.05 0.05 0.05 0.05 0.10 0.10 0.10 0.10 0.10 0.10

1.2 2.0 5.0 7.0 10.0 1.1 2.1 4.0 6.0 8.0 10.0

ζ potential −5.0 −4.2 −19.7 −39.7 −49.4 −7 −5.7 −16.0 −32.3 −43 −48

± ± ± ± ± ± ± ± ± ± ±

0.7 0.7 0.4 0.8 0.7 1 0.4 0.5 0.2 1 1

deprotonation of HNTs surfaces, which progressively show a more significant net negative charge (Table 7). The ionic strength change from 0.05 to 0.1 mol dm−3 does not alter significantly the ζ potential values as expected by the small shift in the distribution diagram of the HNT1 and HNT2 species versus pH (Figure S1; data in Table 7). For a quantitative 7856


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The Journal of Physical Chemistry C comparison between ζ potential and potentiometric data, the surface charge (σ) in millimoles of elementary charge per surface unit was estimated from the ζ potential values as established by the Grahame:47 σ = (8RTεε0C0)0.5 sinh(Zeζ /(2k bT ))

complexes between the HNTs and the ions of the supporting electrolytes were determined with quite good precision. Their stability is in fairly good agreement with interactions of similar stoichiometry reported in the literature. In general, it was determined that the formation of the weak complexes produces a stabilization of the protonated species of both sites. The solubility of the Halloysite was studied in terms of total concentration of aluminum and silicon as a function of ionic strength and time. The results indicate that HNT2 units are more soluble than the HNT1 ones and that an increase of both ionic strength and contact time produces a slight increase of the solubility of the two units, according to the following equations:

(10)

where ε is the solvent relative dielectric constant, ε0 is the vacuum dielectric constant, C0 is the bulk concentration of electrolyte, e is the electron elemental charge, and Z is 1 for NaCl. The full charge per gram of Halloysite nanoparticle can be calculated from the distribution diagram obtained by potentiometric titration, considering the charge of each species and the millimole per gram of Halloysite for each site. In particular, each HNT2 site was considered with a −1 charge, whereas H2(HNT1) and H(HNT1) were assigned a +1 charge. Because the millimole of each site per mass unit of Halloysite was determined by fitting the potentiometric data, the full charge of Halloysite per gram of material can be computed as a function of pH. Results are reported in Figure 5 and show a

log ST(HNT) 1 ± 0.06 = −4.69 − 0.0010h/24 + 0.36I + 0.007Ih /24 log ST(HNT2) ± 0.06 = −4.18 + 0.0013h/24 + 0.08I + 0.007Ih /24

It was proven that the knowledge of the acid−base equilibrium features of Halloysite provides a powerful tool to predict the surface charge of the nanomaterial in agreement to ζ potential experiments.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.6b01127. Distribution diagrams. (PDF)

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AUTHOR INFORMATION

Corresponding Author

Figure 5. Surface charge values calculated by Grahame equation for Halloysite as a function of pH in NaCl aqueous solution at I = 0.10 mol dm−3 (●). The dotted line is calculated from potentiometric data.

*Tel.: +39-091-23897959. Fax: +39-091-590015. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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good agreement between the charges versus pH trends calculated by ζ potential and that predicted by acid−base equilibria. The comparison involving the two curves becomes quantitative if one assumes a specific surface of ca. 30 m2 g−1, which is in straightforward agreement with the experimental BET data, ranging between 20 and 80 m2 g−1 for Halloysite samples from different deposits.3 Such a satisfactory agreement supports the reliability of the acid−base equilibrium model and of it predicting the surface charge features of the nanotubes as a function of pH.

ACKNOWLEDGMENTS Italian MiUR is acknowledged for funding through the projects PON R&C 2007−2013 PON03PE_00214_1 “Nanotechnology and Nanomaterials for Cultural Heritage − TECLA” and FIRB 2012 prot. RBFR12ETL5 “Clay nanotubes for designing ecocompatible smart materials”.

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REFERENCES

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5. CONCLUSIONS The main purpose of this work is to furnish thermodynamic data and predictive equations on the protonation features of HNTs in variable pH, ionic strength, ionic medium, and concentration conditions. In particular, the quantitative data, here reported for the first time, provide the possibility to calculate the positive and negative charge values in the inner and outer surfaces of HNTs under the experimental conditions that characterize the numerous applications. A rigorous approach in the study of the acid−base properties of Halloysite has been carried out using four different models previously adopted for polyelectrolytes and other molecules. The results, in terms of goodness of fit and dependence of the equilibrium constants on ionic strength, are satisfying, indicating that this data treatment is applicable to this molecule. Four weak 7857


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