Article pubs.acs.org/est
Aggregation Kinetics of Manganese Dioxide Colloids in Aqueous Solution: Influence of Humic Substances and Biomacromolecules Xiaoliu Huangfu, Jin Jiang,* Jun Ma,* Yongze Liu, and Jing Yang State Key Laboratory of Urban Water Resource and Environment, School of Municipal and Environmental Engineering, Harbin Institute of Technology, Harbin 150090, China S Supporting Information *
ABSTRACT: In this work, the early stage aggregation kinetics of manganese dioxide (MnO2) colloids in aqueous solution and the effects of constituents of natural organic matter (i.e., Suwannee River fulvic acid (SRFA), Suwannee River humic acid (SRHA), alginate, and bovine serum albumin (BSA)) were investigated by time-resolved dynamic light scattering. MnO2 colloids were significantly aggregated in the presence of monovalent and divalent cations. The critical coagulation concentrations were 28, 0.8, and 0.45 mM for NaNO3, Mg(NO3)2, and Ca(NO3)2, respectively. The Hamaker constant of MnO2 colloids in aqueous solution was 7.84 × 10−20 J. All the macromolecules tested slowed MnO2 colloidal aggregation rates greatly. The steric repulsive forces, originated from organic layers adsorbed on MnO2 colloidal surfaces, may be mainly responsible for their stabilizing effects. However, the complexes formed by alginate and Ca2+ (>5 mM) might play a bridging role and thus enhanced MnO2 colloidal aggregation instead. These results may be important for assessing the fate and transport of MnO2 colloids and associated contaminants.
■
INTRODUCTION The occurrence of manganese dioxide (MnO2) colloids in natural and engineered aquatic systems is well documented.1−3 In natural waters, MnO2 colloids are commonly formed by various processes (e.g., weathering of minerals and biological catalysis of dissolved manganese(II)).4,5 In engineered systems, MnO2 colloids are normal solid products of permanganate (Mn(VII)) reduction or dissolved manganese(II) oxidation.6 Once MnO2 colloids are formed, their surface reactions, for instance, adsorption/desorption and redox, etc., may affect the fate and transport of natural organic matter and synthetic contaminants in some aquatic environments.3,7,8 Also, these surface reactions might play an important role in the oxidative process involving Mn(VII).3,9 However, MnO2 colloids can undergo aggregation in aqueous solution, so their surface reactivity such as adsorption, oxidation, and catalysis may be changed.10−12 Thus, data on the MnO2 colloidal stability and aggregation kinetics in aquatic environments are critical for understanding the fate and transport of themselves as well as associated contaminants in natural waters and engineered systems (e.g., water/wastewater treatment plants). Nevertheless, to date, the fundamental aggregation of MnO2 colloids has not yet been studied. Numerous studies have been conducted to examine the aggregation behavior of various nanomaterials, such as silica nanoparticles (NPs), carbon nanotubes, and TiO2 NPs in synthetic and/or natural waters.13−15 According to classical Derjaguin−Landau−Verwey−Overbeek (DLVO) theory, the stability of nanoparticles was determined by their energy © 2013 American Chemical Society
barrier, which was strongly dependent upon solution chemistry (e.g., pH, ionic strength, and electrolyte ion valence).15 The interactions between colloids were more complicated when macromolecular organic matter (e.g., humic substances (HS), polysaccharides, and proteins) was present. The extent to which these macromolecules could stabilize colloidal particles was determined by the combining effects of the electrostatic, steric, and bridging forces induced by their adsorption on particle surfaces. For instance, fulvic acid (FA) and humic acid (HA) were widely reported to be able to stabilize colloids (e.g., heamatite NPs, single-walled carbon nanotubes (SWNTs), ZnO NPs, etc.) by increasing electrostatic16,17 and/or steric repulsive energies.14,18,19 Similarly, alginate (a widely used model compound for polysaccharides) was reported to stabilize NPs (e.g., SWNTs and hematite NPs) in the presence of Na+ or low concentrations of Ca2+,14,20 while with high concentrations of Ca2+, alginate could increase their aggregation rates instead due to the bridging effect of alginate-Ca2+ complexes.14,20 Protein, an important biomacromolecule, could also stabilize NPs primarily by the increase of steric repulsion. For instance, Saleh et al. found that bovine serum albumin (BSA) could enhance the stability of SWNTs colloids remarkably.14 Similarly, the significant steric stabilization of TiO2 colloids was Received: Revised: Accepted: Published: 10285
January 24, 2013 August 13, 2013 August 15, 2013 August 15, 2013 dx.doi.org/10.1021/es4003247 | Environ. Sci. Technol. 2013, 47, 10285−10292
Environmental Science & Technology
Article
papers and the grid was air-dried for 10 min. For the aqueous samples of MnO2 colloids with alginate and Ca(NO3)2, the same procedure used in the following aggregation experiments were employed. About 2 min after finishing this aqueous sample preparation, the TEM grid preparation was then carried out. The particle sizes and electrophoretic mobility (EPM) of MnO2 colloids were evaluated by a Zetasizer (Nano ZS90, Malvern, UK) operating with a He−Ne laser at a wavelength of 633 nm. A series of freshly prepared MnO2 colloidal samples and more than 10 runs were used for the measurements, and the average data were presented. The MnO2 stock suspension was found to be stable for at least 3 months (Figure S1a and S1b). To determine the impact of pH on the EPM of suspensions at ionic strength of 10 mM (NaNO3), the pH in the range from 2 to 12 was adjusted by HNO3 and NaOH. The EPM of MnO2 suspensions were also measured over a wide range of salt concentrations in the absence and presence of the macromolecular organic matter (i.e., SRFA, SRHA, alginate, and BSA). To prepare these samples, organic matter and salt stock solutions were added immediately prior to the EPM measurements in an identical manner as that for the aggregation experiments below. HNO3 and NaOH were used to adjust pH to 6. Aggregation Experiments. Time-resolved dynamic light scattering (TR-DLS) was used to measure the increase of MnO2 colloidal intensity-weighted averaged hydrodynamic diameter (Dh) with time. It was conducted on Zetasizer (Nano ZS90, Malvern, UK) operating with a He−Ne laser at a wavelength of 633 nm. In these experiments, a certain amount of electrolyte stock solution was added into the DLS cuvette containing a predetermined volume of diluted MnO2 colloidal suspension and macromolecular organic matter so that the total volume of the final suspension was 2 mL. The concentration of MnO2 colloids in DLS vial was controlled at 30 μM. The quartz cuvette was then immediately vortexed for 1 s and placed into the light scattering unit. Immediately, the Dh measurement was started and monitored over a time period of 15−60 min. All DLS measurements were conducted at a scattering angle of 90° and each autocorrelation function was accumulated for 15 s. Determination of Aggregation Kinetics. The initial increase of the MnO2 colloidal Dh with time (dDh(t)/dt) was calculated in the time period from aggregation initiation (t0) to the time when Dh exceeded 1.50Dh,0. The initial aggregation rate constants (k) are proportional to (dDh(t)/dt):23
reported in the presence of fetal serum albumin and human serum albumin.21 In this context, we presented the first experimental work to determine MnO2 colloidal stability as a function of the ionic strength of NaNO3, Mg(NO3)2, and Ca(NO3)2 solutions using the dynamic light scattering (DLS) approach. The effects of model constituents of HS (i.e., Suwannee River fulvic acid (SRFA), Suwannee River humic acid (SRHA)) and biomacromolecules (i.e., alginate and BSA) on the aggregation kinetics were also investigated in the presence of monovalent (Na+) or divalent (Ca2+) cations. The critical coagulation concentrations (CCC) for the monovalent and divalent cations determined from the kinetic studies were used to assess the stability of MnO2 colloids under various solution conditions.
■
MATERIALS AND METHODS Materials. All chemicals were used as received. KMnO4, Na2S2O3, NaOH, HNO3, Mg(NO3)2, and Ca(NO3)2 were purchased from Sinopharm Chemical Reagent Co., Ltd. BSA was obtained from Beijing Solarbio Science & Technology Co., Ltd. Sodium alginate was purchased from Sigma-Aldrich Company. SRFA (standard II) and SRHA (standard I) were obtained from International Humic Substances Society. The stock solutions (500 mg/L, pH 6) of SRFA, SRHA, alginate, and BSA were prepared in deionized (DI) water (>18.2 MΩ/ cm) and stored at 4 °C. The work solutions of HS and biomacromolecules were maintained at 2 mg/L of TOC during the MnO2 NP aggregation and characterization experiments. Synthesis and Characterization of MnO2 Colloids. MnO2 colloids were synthesized following the method used in our previous study.9 Briefly, KMnO4 solution was rapidly stirred with a magnetic stirrer, and purged with N2. Then, the stoichiometric amount of Na2S2O3 solution was added dropwise. The brown MnO2 colloids were formed immediately and continuously stirred over 12 h. This stock solution was stored in the dark at 4 °C. The MnO2 colloidal concentrations in solutions were determined through Mn contents by ICP-MS (NexION 300Q, PerkinElmer Corp.) after reducing MnO2 to Mn2+ by hydroxylamine in excess. The average oxidation state of colloidal MnO2 suspension was determined by the iodimetric method 22 according to eq 1: OSav = 2 +
2A351 ε351l[Mn]T
(1)
where A351 is the absorbance of the solution in excess of I− at 351 nm, ε351 is the molar absorptivity of tri-iodide ion (I3−) at 351 nm, l is the optical path length of the cuvettes (1 cm), [Mn]T is the total manganese concentration. The XPS spectra were measured on a PHI 5700 ESCA System using Al Ka radiation (1486.6 eV) to determine the oxidation state of Mn on the surface of MnO2. Aliquots of freshly synthesized MnO2 colloidal suspension were freezedried for XPS analysis. UV−vis spectra from 200 to 800 nm for aqueous MnO2 colloids were obtained by using a Varian Cray 300 spectrometer. TEM was used to verify the size, shape, and morphology of MnO2 colloids and aggregates. Images were captured by a JEOL-1200 EX TEM (JEOL Inc.) operated at 120 kV. TEM samples were prepared by depositing several drops of the MnO2 suspension on a 200 mesh carbon coated copper grid. The excess water was then removed with filter
k∝
1 ⎛ dDh(t ) ⎞ ⎜ ⎟ N0 ⎝ dt ⎠t → 0
(2)
where N0 is the initial particle concentration, Dh(t) is the intensity-weighted average hydrodynamic diameter at time t. The value of dDh(t)/dt was obtained through linear leastsquares regression analysis. This relationship has been widely used in the aggregation kinetics studies of various NPs.8,24−26 The particle attachment efficiency (α), otherwise known as the inverse stability ratio (1/W), is used to quantify the aggregation kinetics of MnO2 colloids. It is defined as normalizing the measured k by the aggregation rate constant (kfast) under diffusion-limited conditions, where the aggregation is independent of the electrolyte concentration. The MnO2 colloidal concentration in all aggregation experiments is kept constant, thus the α value is calculated by the initial slope of the aggregation profile at a given electrolyte concentration (i.e., 10286
dx.doi.org/10.1021/es4003247 | Environ. Sci. Technol. 2013, 47, 10285−10292
Environmental Science & Technology
Article
absence of organics. In the presence of NaNO3 (Figure 1a), the increase of the electrolyte concentration led to a corresponding
dDh(t)/dt) normalized by the initial slope under favorable (fast) aggregation conditions (i.e., (dDh(t)/dt)fast):23,24,27 1 k α= = = W k fast
( (
dD h (t ) dt
dD h (t ) dt
)
t→0
)
t → 0,fast
(3)
To calculate α in the presence of HS, alginate, and BSA, (dDh(t)/dt)t→0,fast is obtained from the average values of (dDh(t)/dt)fast in the diffusion-limited regime in the presence of either NaNO3 or Ca(NO3)2. Ohshima’s Soft Particle Theory. Theoretical analysis of EPM data by Ohshima’s soft particle theory could define the characteristics of the adsorbed organic layers and further our understanding on the interactions between colloids. Because both organic matter (i.e., SRFA, SRHA, Alginate, and BSA) and MnO2 colloids carry charges at various aqueous conditions, the EPM data can be interpreted according to Ohshima’s soft theory for a charged bare particle with a charged layer. Detailed analysis can be found in the SI (Text S1, Table S1, and Figure S2). A MATLAB code employing iterative least-squares minimization was used for the model fitting. The Ohshima’s soft particle theory was widely applied for characterizing the adsorbed layers on various NPs.18,28,29
■
RESULTS AND DISCUSSION Characteristics of Aqueous MnO2 Colloids. The average Mn oxidation state of MnO2 colloids was 4.03 ± 0.04 (n = 5) from the iodimetric method, consistent with the results obtained from XPS analysis (Figure S3). The UV−vis spectra (Figure S1a) of colloidal suspensions showed maximum absorption peaks at the wavelengths of 215 and 365 nm. The MnO2 colloidal sizes obtained from DLS measurements (Figure S4c) exhibited a narrow distribution with the peak width at half-maximum was 11.12 nm. These particles had the diameters of 24−105 nm with an average value of 55.86 ± 0.26 nm (n = 30). However, Joaquin and co-workers reported that the size range of MnO2 colloids as prepared by this method was 89−193 nm, with a peak at 120 nm which was greater than the value we obtained.22 The exact reasons for the difference are not clear. The diameters of freshly prepared MnO2 colloids determined by TEM (Figure S4a and b) were slightly greater than the one obtained by in situ DLS measurements (Figure S4c). The difference may be attributed to the aggregation during TEM samples preparation. Such aggregations were also reported for other NPs such as Ag NPs.30 MnO2 aggregates showed an irregular structure (Figure S4a and S4b). The absolute zeta potential (ζ potential) of these colloids were lower than previously reported (−41.96 ± 1.88 mV vs −47 ± 3 mV).22 When the pH increased from 2 to 12, the EPMs (Figure S5) in 10 mM NaNO3 varied in the range of (−2.0 to −3.25) × 10−8 m2 V−1 s−1. It should be noted that MnO2 NP properties were largely variable (e.g., size, structure, morphology, and surface charge). Thus, their aggregation and mobility may correspondingly be different in aquatic environments. Nevertheless, it seems reasonable to prepare this MnO2 NP as model colloids to gain insight into their aggregation kinetics in relevant water chemistry. MnO2 Colloidal Aggregation Kinetics with Monovalent and Divalent Cations. The aggregation behaviors of MnO2 colloids clearly exhibited DLVO-type interactions in the
Figure 1. Aggregation profiles of MnO2 colloids in various electrolyte solutions: (a) NaNO3; (b) Mg(NO3)2; (c) Ca(NO3)2. Aggregation experiments and EPM measurements were conducted at pH 6 and 25 °C.
increase in the aggregation rate at relatively low concentration regime (i.e.,
