Colloid Systems and Interfaces Stability of Cerium Oxide Nanoparticles in Aqueous Environments: Effects of pH, Ionic Composition, and Suwanee River Humic and Fulvic Acids

Abstract

This study investigates the colloid systems and interfaces stability of cerium oxide nanoparticles in aqueous environments as a function of pH, monovalent cations (Na⁺) and divalent cations (Ca²⁺), and humic substances (humic acid (HA) and fulvic acid (FA)). Results show that the solution chemistry affected the colloidal stability and aggregation kinetics of CeO₂ NPs. The pH point of zero charge (pH_PZC) of CeO₂ NPs was measured at pH 10.2 with diameter of CeO₂ NPs aggregates of ∼1,700 nm. The effects of Na⁺ and Ca²⁺ and HA and FA on the magnitudes and rates of aggregation were pH-dependent. In addition, when salts were present in the aqueous systems, although the CeO₂ NPs were stable at pH < pH_PZC (expect for 1 mM of NaCl/CaCl₂) and pH > pH_PZC (except for 0.5 mM CaCl₂), the aggregation was enhanced at pH = pH_PZC, with the diameter of CeO₂ NPs in the ∼1,300–3,600 nm range. HA also stabilized CeO₂ NPs under pH > pH_PZC with an enhanced aggregation of pH = pH_PZC with the diameter of CeO₂ NPs in the ∼1,500–1,900 nm range, and in the presence of 0 and 1 mM of NaCl/CaCl₂ at pH < pH_PZC. At three pH levels (8.2, 10.2, and 12.2) and under all different electrolyte concentrations (0–1 mM of NaCl or CaCl₂), FA (0.14 mg/L) exhibited a greater degree of efficiency in stabilizing CeO₂ NPs than HA (5 mg/L), with CeO₂ NPs aggregates growing at low rates and resulting in diameter of ∼95–115 nm.

1. Introduction

Revolutionary advances in the field of nanoscience and nanotechnology have played a key role in the development of new materials and technologies for products and applications to all the sectors of industry—agriculture, manufacturing, and services—and have impacted our daily life. Among the nanoparticles engineered in nanoscience, the metal nanoparticles (NPs) CeO₂ NPs are a common nanomaterial used increasingly in products and applications [1, 2]. CeO₂ NPs have unique physicochemical properties, such as strong ultraviolet (UV) radiation absorbing capacity, high capacity for donating and storing oxygen, and low Ce^IV/Ce^III redox potential [3–6]. CeO₂ NPs are used in exhaust gas catalysts in diesel fuel additives [7–9], capacitors and semiconductors in electronic devices [10, 11], antioxidants in medical products [12], UV-blocking agents in sunscreen [7], and polishing agents in manufacturing glass [8].

With the increasing use in products and applications, the presence of CeO₂ NPs in environmental components, such as soil, sediments, water, air, and biota, is inevitable. Although CeO₂ NPs in diesel fuel additives increase the efficiency of trapping particulate matter and decrease NO_x emissions during combustion in engines, CeO₂ NPs are present in the exhaust emissions that discharge into the air [7, 13, 14]. CeO₂ NPs released from fertilizers can also contaminate the soil. The estimated CeO₂ NPs concentrations in biosolids, which are calculated by using 2010’s market study production estimates in three metropolitan areas (New York, Shanghai, and London), range from 0.53 to 9.1 mg/kg [15, 16]. By using the same method of estimation, the range of the concentration of CeO₂ NPs released from wastewater treatment plants into waterbodies is 0.003–1.17 µg/L [15, 16]. Other possible sources of release of CeO₂ NPs into soil and water systems include discarded electronics or residue from coatings or sunscreen products [1, 17]. Once released into the environment, CeO₂ NPs may be contaminant carriers or serve as ecological and public health threats themselves.

The toxicity effects of CeO₂ NPs on plants, such as lettuce, rice, cucumber, tomato, and soybean, have been investigated [18–26]. For example, although the addition of CeO₂ NPs to fertilizer may enhance root growth in rice seedlings [27], it can alter the composition and nutritional values of crop plants, such as soybeans, cowpeas, corn, and mungbean [27–30]. Some studies have also involved determining the toxicity of CeO₂ NPs on microorganisms. Microbial communities were found to be inhibited by CeO₂ NPs during the wastewater treatment process [12]. The activities of soil enzymes in a soil–plant system were also affected by CeO₂ [31]. CeO₂ NP toxicity to invertebrates on land and in water systems has also been reported in studies involving earthworms, nematodes, algae, and phytoplankton [15, 32–34].

The transport of CeO₂ NPs in natural porous media (e.g., soil and aquifer systems) and in engineered porous media (e.g., sand filtration systems) was also the subject of study [35–40]. In terms of the physicochemical parameters, CeO₂ NP transport in saturated sand-packed columns was hindered in the presence of NaCl with high ionic strength (IS) values (larger than 10 mM) at pH 3 [37]. Organic matter enhanced the stability and mobility of CeO₂ NPs in the presence of 1 mM NaCl at pH 6.5 [41]. The increase of HA concentration from 0 to 10 mg/L or the decrease of IS from 100 to 1 mM markedly facilitated the transport of CeO₂ NPs more than the influence of the increase of pH from pH 7 to 10 [38]. The retention of CeO₂ NPs was observed in loamy sand as well as under high concentration levels of CaCl₂ and MgCl₂ at pH 8 [39].

Colloidal stability and aggregation state of engineered nanoparticles—such as CeO₂ NPs—are key criteria governing the environmental behavior of nanoparticles [42, 43]. The stability of nanoparticles controls their fate and transport in terrestrial and aquatic systems, bioavailability, and toxicity, as the size of the nanoparticles/nanoaggregates is essential in determining their reactivity [44, 45]. The stability of CeO₂ NPs is a function of their size [46–48], surface chemistry and capping agent [47, 49, 50], and surface charge [51, 52] which is influenced by the physicochemical conditions of the soil water environment the nanoparticles encounter, such as pH [45, 48, 53–56], IS [43, 53, 56–58], electrolyte types [48, 52, 56, 59, 60], and organic compounds [47, 52–54, 61, 62], including humic acid (HA) [43, 57, 63] and fulvic acid (FA) [45, 58, 63, 64].

In terms of various environmental parameters, one of the most importance of which was pH, the colloidal stability and aggregation of engineered nanoparticles were the subject of a number of studies [65–67]. The pH of suspensions affects the surface ionization of nanoparticles. The stability of NPs has been investigated in different electrostatic scenarios related to pH_PZC, specifically pH > pH_PZC, pH = pH_PZC, and pH < pH_PZC [45, 68, 69]. For example, the pH_PZC of the uncoated CeO₂ NPs, which had a primary particle diameter 28 ± 10 nm and a specific surface area 27.2 ± 0.9 m²/g, is equal to 6.8 ± 0.1 [45]. An increase of the pH from 3 to 10 decreased the zeta potential from positive to negative values. Further, the size of the CeO₂ NP nanoaggregates was lower than 200 nm at pH < pH_PZC and reached 800–1,500 nm at pH = pH_PZC.

IS plays an important role in controlling the behavior of NP stability [70–76]. The increase of IS, affecting NPs by electric double layer (EDL) compression, zeta potential reduction, and charge reversal by counterions, generally enhances NP aggregation [17, 77, 78]. Electrolyte valence is also an additional and important factor for determining the aggregation efficiency of electrolytes. For example, at pH 4.8 and 0.0128 M, CaCl₂ had a faster rate of aggregate formation of TiO₂ NPs than that of NaCl [79]. Other research also demonstrated that a lower critical coagulation concentration (CCC) was observed in CeO₂ NPs for Ca²⁺ (16 mM) than Na⁺ (80 mM) with a proportionality fraction of ^−2.3 where is the counterion valence [52].

Natural organic matter (NOM), which contains humic substances (HS), polysaccharides, algal, and bacterial residues, also affects the stability of NPs [45, 80]. As a major component of NOM, HS is present in soil, water, and geological organic deposits. HS in natural water systems, which compose 20%–50% of NOM in the aquatic environment, have three components: HAs, FAs, and humin [45, 81]. HS functional groups, such as carboxylate, phenolate, amino, and thiol, can attach themselves to minerals or metal oxide bulk particles and NPs [82, 83]. NPs stability is affected by the NOM coating formation types on NP surfaces [84]. In addition, the concentration, composition, molecular mass, hydrophobicity, and polarity of NOM and the couple effects of NOM and pH, or different types of cations, also determine NP stability [85–87]. For instance, the stability of CeO₂ NPs and complexes of FA and CeO₂ NPs were studied at three pH levels (3, 7, 10) that corresponded to the positive (pH < pH_PZC), neutral (pH = pH_PZC), and negative (pH > pH_PZC) surface charge of CeO₂ NPs, respectively [45]. The environmental amount of FA (2 mg/L) at pH 3 (pH < pH_PZC) was capable of stabilizing CeO₂ NPs. The CCC of CeO₂ NPs for KCl shifted to higher levels (>500 mM) as the SRHA concentration increased (from 0 to 10 mg/L) at pH 5.7 [47]. At pH 5.7 and at a low SRHA concentration (1 mg/L), SRHA stabilized the CeO₂ NPs at a high KCl concentration (500 mM) [47].

Despite significant advances in the understanding of the mechanisms and parameters that govern the fate and behavior of engineered nanomaterials (ENMs) in the environment, “many knowledge gaps remain” [42]. Studying ENMs in “more complex real-world systems” and “relevant environmental systems” [42], as in aquatic systems, is essential to further the understanding of the fate and transport of ENMs in the earth’s system of natural resources and to formulate regulations on their discharge in the aquatic, terrestrial, and atmospheric systems [42, 43].

Therefore, it is critical to further the understanding of the fate and behavior of prominent ENMs, like CeO₂ NPs, in complex environmental systems for the purpose of protecting the ecosystems and human health [42, 43]. Most of all, knowledge of the “aggregation state” and “characterization of homo- and heteroaggregates” of ENMs in aquatic and terrestrial systems needs to be improved, as outlined by Surette et al. [42] in their assessment of research needs in the field of ENMs and their interactions with the environment. Moreover, Surette et al. [42] highlighted the need to explore the effects of specific physicochemical parameters of aqueous systems on ENMs. These parameters include “IS/composition, pH, and organic matter” and relate to research needs focusing on “experimental condition” as proposed by Surette et al. [42]. It also responds to the necessity to conduct studies exploring the fate of ENMs in aquatic systems as mentioned by Bathi et al. [43].

Understanding of the synergistic effect of the physical and chemical properties of CeO₂ NPs and the chemistry of their environment—pH, ionic composition, and NOM (HA and FA)—is essential to further the understanding of the processes governing the colloidal stability, interfacial interactions, and aggregation of CeO₂ NPs in water and environmental systems. Interfacial interactions govern the behavior of CeO₂ NPs in the environment, and their threat to public health and the environment. Yet, the synergy between CeO₂ NPs as a function of pH, ionic composition, and organic matter that induces alterations in colloid systems and interfaces stability is complex and varied. Then, it is crucial to investigate the self-interactions between CeO₂ NPs, the interactions between CeO₂ NPs and monovalent/divalent cations, and the interactions between CeO₂ NPs and monovalent/divalent cations-humic substance systems under different pH conditions, as well as their ensuing assemblage of structures/aggregation in order to discover their affect on the colloidal phenomena and interfacial interactions of CeO₂ NPs in aqueous environments.

As discussed in our recent study titled “Colloidal phenomena and aggregation mechanisms of cerium oxide nanoparticles in aqueous systems: Effects of monovalent and divalent cations, and Suwanee River humic and fulvic acids,” colloid systems and interfaces stability of CeO₂ NPs in aqueous systems are impacted by electrolytes and organic matter [88]. Therefore, the objectives of this study are to (1) analyze the colloidal phenomena and interfacial interactions of CeO₂ NPs as a function of pH and determine the pH_PZC, (2) investigate the colloidal phenomena and interfacial interactions of CeO₂ NPs at three different pH regimes which relate to pH_PZC (below, at, and above pH_PZC) along with monovalent and metallic cation Na⁺ or divalent and metallic cation Ca²⁺, and (3) study the colloidal phenomena and interfacial interactions of CeO₂ NPs at three different pH regimes which relate to pH_PZC (below, at, and above pH_PZC) along with metallic cations (Na+ or Ca²⁺)—HS (HA or FA). A laboratory-scale batch experiment approach was conducted to assess and quantify the influence of pH, metal cations, IS, HA, and FA on the colloidal phenomena and interfacial interactions of CeO₂ NPs. Measurement techniques of dynamic light scattering (DLS) and electrophoretic light scattering (ELS) were utilized to determine the hydrodynamic diameter and zeta potential of CeO₂ NPs, respectively. Aggregation kinetics and attachment efficiency were calculated from the laboratory scale batch experimental data. The net energy versus interparticle distance for CeO₂ NPs in water and environmental systems were estimated using the Derjaguin, Landau, Verwey, Overbeek (DLVO) theory, i.e., the sum of van der Waals attraction and EDL repulsion between two approaching particles.

The novelty of this research consists of the analysis of the mechanisms governing the interfacial interactions between the surface of CeO₂ NPs as a function of pH, ionic composition, and Suwanee river humic and fulvic acids, the description and measurement of the dynamic growth of these nanoaggregate–CeO₂ complexes, and the evaluation of their colloidal stability in aqueous environment.

2. Materials and Methods

2.1. Materials

2.1.1. CeO₂ NP Stock Suspensions

CeO₂ NPs were obtained in the form of cerium (IV) oxide nanoparticles from Sigma–Aldrich (Saint Louis, MO). According to the manufacturer, the average particle size and density of CeO₂ NPs are below 25 nm and 7.13 g/cm³ at 25°C, respectively. The particle size of CeO₂ NPs was obtained by application of the Brunauer, Emmett, and Teller (BET) theory, and the structural conformity of the CeO₂ NPs was measured by X-ray diffraction, as stated in the product information from Sigma–Aldrich. Every CeO₂ NP stock suspension (250 mg/L) was prepared according to the following dispersion protocol. First, a UMT2 ultramicrobalance (Mettler Toledo, Columbus, OH) in a Nano Enclosure Xpert®, 38,872 series (Labconco, Kansas City, MO) was used to weigh 250 mg of cerium (IV) oxide nanoparticles. Next, the weighted cerium oxide (IV) nanoparticles were dispersed in 80 mL ultrapure deionized (DI) water (18.2 MΩ cm) (Millipore Corporation, Billerica, MA) by ultrasonication, and diluted to a 1 L suspension. An ultrasonic liquid processor (S-4000, output power 600 W Max and output frequency 20 KHz, Misonix, Newton, CT) was used for the ultrasonication of the CeO₂ NP stock suspensions with the following settings: an amplitude of 45, a power of 45 W, and a sonication of 4 min.

2.1.2. HA and FA Stock Suspensions

Suwannee River humic acid (SRHA) (standard II, 2S101H) and Suwannee River fulvic acid (SRFA) (standard II, 2S101F) were obtained from the International Humic Substances Society, St. Paul, MN. A 50 mg/L SRHA stock suspension was first made by dissolving 50 mg of SRHA in 1 L of ultrapure DI water. A 1.41 mg/L of SRFA stock suspension was next prepared by dissolving 1.41 mg of SRFA powder in 1 L of ultrapure DI water. Both SRHA and SRFA suspensions were then filtered through Whatman™ hardened ashless filter papers, i.e., cotton filters made of high-quality cotton linters (Whatman quantitative filter paper, hardened ashless, Grade 542, GE Healthcare Life Sciences, Little Chalfont, UK) with nominal particle retention rating of 2.7 μm. The 50 mg/L SRHA and 1.41 mg/L SRFA suspensions were stored in the dark at 4°C in a refrigerator.

2.2. Preparation of CeO₂ NP Suspensions

CeO₂ NP stock suspension was sonicated in an ultrasonic bath (CPX2800H, Branson Ultrasonics Corporation, Danbury, CT) with settings of 110 W for power (maximum power) and 40 KHz for frequency for 30 min prior to use. A diluted CeO₂ NP suspension (25 mg/L) was obtained by mixing 10 mL of CeO₂ NP stock suspension with 90 mL of solution with specific IS in the absence and presence of SRHA or SRFA. The pH values, which ranged from 1 to 14 for the diluted CeO₂ NP suspensions (25 mg/L), were adjusted with HCl or NaOH (BDH®, VWR International, LLC, Radnor, PA) at different concentrations (0.01, 0.1, and 1 M) and measured with a FiveEasy™ Plus pH meter (Mettler-Toledo, Colombus, OH). Once the pH_PZC was determined from the analysis of the zeta potential measurements, two other pHs, which were above and below the pH_PZC, were selected for the later three pH region experiments.

Four concentrations (0.05, 0.1, 0.5, and 1 mM) and one control (0 mM) of NaCl (Fisher Scientific, Fair Lawn, NJ) and CaCl₂ (Amresco, Solon, OH) were prepared to obtain specific IS for each CeO₂ NP suspension. In addition, HA (5 mg/L) and FA (0.14 mg/L) were prepared using the specific IS concentration solutions (0, 0.1, and 1 mM) to obtain specific HS and IS conditions for each CeO₂ NP suspension.

Three pHs, greater than, equal to, and less than the pH_PZC, were then adjusted to obtain different pH–IS and pH–HS system suspensions. All the suspensions in the different pH–IS or pH–HS systems were prepared by diluting 250 mg/L CeO₂ NP stock suspension with specific IS or specific HS suspension and IS solution in 100 mL beakers employing a magnetic stirring system. The HCl and NaOH were used to adjust the pH of each suspension. The suspension was stirred until reaching the specific pH. All the suspensions were produced in four replicates (sets) for subsequent DLS and ELS analyses. A summary of the solution chemistry of the CeO₂ NP suspensions analyzed in this study is given in Table 1.

Table 1 

Solutions created for experimentation on the stability and aggregation of cerium oxide NPs in aqueous systems as a function of pH, electrolyte type and concentration, and the absence and presence of humic substances.

2.3. Aggregation Experiments

These diluted 25 mg/L CeO₂ NP suspensions were sonicated with the Misonix S-4000 ultrasonic liquid processor (50 amplitude and 45 W power) for 4 min before testing. A NanoBrook 90Plus Zeta Particle Size Analyzer (Brookhaven Instruments Corporation, Holtsville, NY) was used to characterize the aggregation behavior of the CeO₂ NPs. The maximum value of automatic measurements for each run in this analyzer’s setting was 10 with manual operation required between two runs to obtain continuous data in approximately 1 hr. DLS was used to analyze the hydrodynamic diameters of CeO₂ NPs in one set of pH–IS or pH–HS system suspensions over approximately 1 hr. ELS also over an hour was then used to measure the zeta potentials of CeO₂ NPs in another set of both pH–IS and pH–HS system suspensions.

2.3.1. Particle Size Measurements

A NanoBrook 90Plus Zeta Particle Size Analyzer from Brookhaven Instruments Corporation (Holtsville, NY) was employed to analyze the particle size and to determine the hydrodynamic diameter of the CeO₂ nanoaggregates as a function of pH, ionic composition, SRHA, and SRFA. DLS measurements were used to establish the CeO₂ nanoaggregate sizes with the DLS measured by introducing the suspensions into BI-SCP plastic cells (10 W × 10 L × 52 H mm, 3.5 mL). Performed over a period of approximately 1 hr for each sample, the DLS was measured in an array of 31 runs, with each composed of 10 measurements with a 10 s time interval between them. The time separating the DLS runs was accounted for when reporting the DLS measurements. All DLS measurements were taken at 25°C.

2.3.2. Zeta Potential Measurements

A NanoBrook 90Plus Zeta Particle Size Analyzer from Brookhaven Instruments Corporation (Holtsville, NY) was employed to measure the zeta potential of the CeO₂ nanoaggregates with respect to pH, electrolytes, IS, SRHA, and SRFA. The zeta potential of the CeO₂ nanoaggregates was analyzed by placing 1.5 mL of the nanoparticles suspension in BI-SCP plastic cells followed by ELS measurements. The ELS was measured at 25°C for each sample within 1 hr to establish the zeta potential of these CeO₂ nanoaggregates. ELS measurements were taken continuously with the data collected and categorized every 10 min. The number of zeta potential measurements within 1 hr varied because the setting for data reading was observations of less than 0.05 relative residual. The time elapsing between the ELS runs was accounted for when reporting the ELS measurements. All ELS measurements were taken at 25°C.

2.4. Aggregation Kinetics Analysis

Nanoparticle aggregation kinetics were determined by monitoring the hydrodynamic radius as a function of time (t). The initial rate of change of was proportional to the initial aggregation rate constant , as well as the initial particle concentration , and expressed as follows:

The was determined by using a linear correlation function to the experimental data during the early stage of aggregation (i.e., the first 600 s) [89–91]. The aggregation kinetics of each suspension were calculated using the data of the hydrodynamic radius of CeO₂ NPs obtained by DLS.

3. Results

3.1. Effect of pH on Stability of CeO₂ NPs

To establish the effect of pH on the stability of CeO₂ NPs, ELS, and DLS methods were used to measure the variation of surface charge and average hydrodynamic diameter. Figure 1 shows the average zeta potential and hydrodynamic diameter of CeO₂ NPs of approximately 1 hr measurements as a function of pH values ranging from extremely acidic (pH 1) to extremely basic (pH 14). The zeta potential of CeO₂ NPs was 10.28 ± 24.48 mV, and the average hydrodynamic diameter of CeO₂ NPs was 1,056.85 ± 345.63 nm at pH 1. When the pH of aqueous suspension changed from acidic to basic conditions (pH 2–12), the average zeta potential of CeO₂ NPs varied from positive values (e.g., 52.53 ± 10.26 mV at pH 4) to negative values (e.g., −52.76 ± 12.46 mV at pH 12), and the size of CeO₂ nanoaggregate was stable around 100 nm except at pH 10. The pH_PZC of CeO₂ NPs was 10.2, where the surface charge of CeO₂ NPs was neutralized, i.e., zeta potential was close to 0 mV. The CeO₂ NP nanoaggregate size increased significantly (1,202.14 ± 442.86 nm) at pH 10, which was close to pH_PZC (Figure 1). Additionally, the hydrodynamic diameter reached about 1,500 nm, and the absolute average zeta potential was less than 10 mV at pH 13 and 14. Figure 2 presents results from the aggregation profiles, i.e., a series of aggregation experiments measuring the hydrodynamic diameters of CeO₂ NPs as a function of time by DLS, over pH values ranging from 1 to 14. After approximately 1 hr of incubation (i.e., ∼1 hr), the CeO₂ NPs increased to the micron level at pH 1, 10, 13, and 14 (Figure 2). However, the hydrodynamic diameter of CeO₂ NPs was constant at other pH values (2–9 and 11–12) and close to 100 nm (Figure 2). The average hydrodynamic diameter of CeO₂ NPs during the first 100 s and at ∼1 hr at pH ranging from 1 to 14 is displayed in Table S1. The size distribution represents the hydrodynamic diameter of NPs in different size ranges in order to analyze the size of NPs distributed at a time interval. The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH ranging from 1 to 14 is presented in Figure S1. The initial constant aggregation rate represents the aggregation kinetics, which ranged between 0.80 and 1.50 nm/s at pH 1, 10, 13, and 14 (Figure 3). However, was less than 0.10 nm/s at other pH values (2–9 and 11–12).

Figure 1 

Zeta potential and hydrodynamic diameter of CeO₂ NPs as a function of pH (1–14). The results are presented as an average value of around 1 hr measurements with the corresponding standard deviation.

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Figure 2 

Hydrodynamic diameter of CeO₂ NPs as function of pH (1–14) over about 1 hr: (a) pH 1–3, (b) pH 4–6, (c) pH 7–9, (d) pH 10−12, and (e) pH 13 and 14.

Figure 3 

Initial aggregation rate constant k₁₁ of CeO₂ NPs as a function of pH (1–14) during the first 600 s.

3.2. Impact of pH and NaCl on Stability of CeO₂ NPs

The effects of NaCl concentrations on CeO₂ NPs stability in three pH domains representing the three electrostatic states—pH = 8.2 < pH_PZC, pH = 10.2 = pH_PZC, and pH = 12.2 > pH_PZC—were examined by measuring the surface charge and size of nanoparticles using ELS and DLS methods.

3.2.1. Positively Charged CeO₂ NPs in the Presence of NaCl (pH = 8.2 < pH_PZC)

The aggregation experiments were conducted at pH 8.2 (pH < pH_PZC), where CeO₂ NPs were positively charged in the absence of NaCl (Figure 1). The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 8.2 and in the presence of monovalent NaCl salt at various concentrations are presented in Figures 4(a) and 4(b) and Figures 5(a) and 5(b), respectively.

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Figure 4 

Hydrodynamic diameter of CeO₂ NPs as function of time in varying concentrations of NaCl ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) 0, 0.05, and 0.1 mM NaCl and pH 8.2; (b) 0.5 and 1 mM NaCl and pH 8.2; (c) 0, 0.05, and 0.1 mM NaCl and pH 10.2; (d) 0.5 and 1 mM NaCl and pH 10.2; (e) 0, 0.05, and 0.1 mM NaCl and pH 12.2; (f) 0.5 and 1 mM NaCl and pH 12.2.

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Figure 5 

Zeta potential of CeO₂ NPs as function of time in varying concentrations of NaCl ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) 0, 0.05 and 0.1 mM NaCl, and pH 8.2; (b) 0.5 and 1 mM NaCl, and pH 8.2; (c) 0, 0.05 and 0.1 mM NaCl, and pH 10.2; (d) 0.5 and 1 mM NaCl, and pH 10.2; (e) 0, 0.05 and 0.1 mM NaCl, and pH 12.2; (f) 0.5 and 1 mM NaCl, and pH 12.2.

The surface charge reflects the inclination of aggregation or stabilization of NPs. In terms of time, the surface charge of CeO₂ NPs in the presence of NaCl at each NaCl concentration at each 10 min interval in 1 hr was relatively stable. It is independent of time at pH < pH_PZC (Figures 5(a) and 5(b) and Table S2). However, when considering the concentration of NaCl, the zeta potential decreased from around 21 mV to around 10 mV when the concentration of NaCl augmented from 0 to 1 mM. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 8.2 in concentration of NaCl 0, 0.05, 0.1, 0.5, and 1 mM are 0.10, 0.84, −0.28, 0.56 and −0.76, respectively.

The results of hydrodynamic diameter measurements were the same as predicted by the results of surface charge. CeO₂ NPs were stable at pH < pH_PZC when the concentration of NaCl was less than 1 mM (Figures 4(a) and 4(b)). When concentration of NaCl reached 1 mM, the hydrodynamic diameter of CeO₂ NPs in about 1 hr increased rapidly by 122.28% during the first 100 s (188.74 ± 29.89 nm) and 1,098.98% (1,346.21 ± 221.29 nm) at ∼1 hr, when compared to the controls during the first 100 s (84.91 ± 3.61 mm) and at ∼1 hr (112.28 ± 2.11 mm), respectively (Table S3). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 8.2 and in the presence of monovalent NaCl salt at various concentrations is presented in Figures S2(a) and S2(b).

The aggregation rate constant an indicator of the stability of NPs under different conditions is obtained by using linear regression of hydrodynamic radius over 10 min in this experiment. The results of also show the same prediction as surface charge measurements. The was 0.35 nm/s at 1 mM NaCl, which was larger than values (<0.02 nm/s) at other NaCl concentrations (Table 2).

Table 2 

Initial aggregation rate constant k₁₁ obtained from the linear regression of the experimental data during the first 600 s, under varying concentrations of NaCl or CaCl₂ ranging from 0 to 1 mM in the absence and presence of HA (5 mg/L) or FA (0.14 mg/L) at three different pH values (8.2, 10.2, and 12.2).

The net energy barrier is also an indicator of the stability of NPs. The larger net energy barrier means more stable NPs, while no barrier indicates the inclination of the aggregation of NPs. The results show that the net energy barrier disappeared at 1 mM NaCl (Figure 6(a)), which is the same prediction as surface charge and hydrodynamic diameter measurements of CeO₂ NPs.

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Figure 6 

Net energy versus interparticle distance according to DLVO theory, showing the influence of representative monovalent electrolyte NaCl or divalent electrolyte CaCl₂ concentrations varying from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2, for CeO₂ NPs suspended in aqueous systems: (a) NaCl and pH 8.2; (b) NaCl and pH 10.2; (c) NaCl and pH 12.2; (d) CaCl₂ and pH 8.2; (e) CaCl₂ and pH 10.2; and (f) CaCl₂ and 12.2. Calculation of DLVO theory includes the contributions of the van der Waals and electric double layer (EDL) interactions. The net energy (Net) shown is the sum of EDL repulsion and van der Waals attraction between two approaching particles.

3.2.2. Uncharged CeO₂ NPs in the Presence of NaCl (pH = 10.2 = pH_PZC)

At pH = pH_PZC, the surface of CeO₂ NPs is neutralized in the absence of NaCl (Figure 1). The aggregation profiles and average zeta potentials of CeO₂ NPs at pH 10.2 and in the presence of monovalent NaCl salt at various concentrations are presented in Figures 4(c) and 4(d) and Figures 5(c) and 5(d), respectively.

In the presence of NaCl, the absolute average zeta potential of CeO₂ NPs varied around 0 mV and was less than 20 mV (Figures 5(c) and 5(d) and Table S2). However, the surface charge of CeO₂ NPs was negative at 0, 0.05, and 1 mM of NaCl, but positive at 0.1 and 0.5 mM of NaCl. The surface charge is independent of time at pH = pH_PZC. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 10.2 in concentration of NaCl 0, 0.05, 0.1, 0.5, and 1 mM are −0.76, −0.48, −0.79, −0.67 and 0.61, respectively. The results of hydrodynamic diameter, , size distribution, and net energy barrier were the same as the prediction from the surface charge results, which suggests that CeO₂ NPs were unstable in the presence of NaCl at pH = pH_PZC. The hydrodynamic diameter of CeO₂ NPs under all the NaCl concentrations grew at high rates and reached several microns after approximately 1 hr of incubation, at ∼1 hr (Figures 4(c) and 4(d)). For example, the hydrodynamic diameter of CeO₂ NPs at 0.1 mM NaCl increased by 618.28% at ∼1 hr (1,430.52 ± 159.61 nm) when compared to the hydrodynamic diameter during the first 100 s (199.16 ± 23.37 nm) (Table S3). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 10.2 and in the presence of monovalent NaCl salt at various concentrations is presented in Figures S2(c) and S2(d). The under all the NaCl concentrations (0.33–0.45 nm/s) was greater than that at pH 8.2 and pH 12.2, and no concentration-dependent increase was observed (Table 2). No energy barriers were present when the Na⁺ concentration was larger than 0.1 mM (Figure 6(b)).

3.2.3. Negatively Charged CeO₂ NPs in the Presence of NaCl (pH = 12.2 > pH_PZC)

At pH = 12.2 (pH > pH_PZC), the CeO₂ NPs exhibited negative zeta potential values in the absence of NaCl (Figure 1). The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 8.2 and in the presence of monovalent NaCl salt at various concentrations are presented in Figures 4(e) and 4(f) and Figures 5(e) and 5(f), respectively.

In the presence of NaCl, the average zeta potential of CeO₂ NPs in 1 hr (0–60 min) was less than −30 mV (Figures 5(e) and 5(f) and Table S2). This finding indicates that CeO₂ NPs are stable at pH > pH_PZC, which is also demonstrated by the results of hydrodynamic diameter, size distribution, and net energy barrier. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 12.2 in concentration of NaCl 0, 0.05, 0.1, 0.5, and 1 mM are −0.64, −0.88, 0.97, −0.47, and −0.74, respectively. The hydrodynamic diameter of CeO₂ NPs at this pH grew at low rates for all NaCl concentrations (Figures 4(e) and 4(f)). For all NaCl concentrations, the average hydrodynamic diameter of CeO₂ NPs was less than 110 nm after approximately 1 hr of incubation, at ∼1 hr. For example, the hydrodynamic diameter of CeO₂ NPs at 0.5 mM NaCl increased by only 19.87% at ∼1 hr (102.39 ± 4.81 nm) when compared to the hydrodynamic diameter during the first 100 s (85.42 ± 3.99 nm) (Table S4). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 12.2 and in the presence of monovalent NaCl salt at various concentrations is presented in Figures S2(e) and S2(f). The was very small (less than 0.02 nm/s) for concentrations of NaCl ranging from 0 to 1 mM (Table 2). The energy barriers were 3.10, 6.00, 4.73, and 2.30 kJ, for 0.05, 0.1, 0.5 and 1 mM of Na⁺, respectively (Figure 6(c)).

3.3. Impact of pH and CaCl₂ on Stability of CeO₂ NPs

The effects of CaCl₂ concentrations on CeO₂ NPs stability in three pH domains representing the three electrostatic states—pH = 8.2 < pH_PZC, pH = 10.2 = pH_PZC, and pH = 12.2 > pH_PZC—were examined by measuring the surface charge and size of nanoparticles using ELS and DLS methods.

3.3.1. Positively Charged CeO₂ NPs in the Presence of CaCl₂ (pH = 8.2 < pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 8.2 and in the presence of divalent CaCl₂ salt at various concentrations are given in Figures 7(a) and 7(b) and Figures 8(a) and 8(b), respectively. At pH 8.2 (pH < pH_PZC), at 0–10 min, the average zeta potential of CeO₂ NPs diminished from 20.46 ± 11.12 mV to 16.21 ± 11.37 mV when the concentration increased from 0 to 0.1 mM of CaCl₂, but it then increased to 21.41 ± 13.45 mV when the CaCl₂ concentration raised to 1 mM (Figures 8(a) and 8(b) and Table S2). However, the manifestation of different trends occurred at 50–60 min. At this point, the average zeta potential increased from 19.88 ± 12.47 mV to 27.65 ± 7.76 mV with the increase of CaCl₂ concentration (increased from 0 to 0.1 mM). It then decreased to 12.71 ± 13.08 mV when the concentration of CaCl₂ was 0.5 mM, and increased to 22.40 ± 7.50 mV when the CaCl₂ concentration was 1 mM. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 8.2 in concentration of CaCl₂ 0, 0.05, 0.1, 0.5, and 1 mM are 0.10, 0.92, 0.82, −0.85, and 0.60, respectively. No obvious prediction was possible from the results of zeta potential when CeO₂ NPs were in the presence of CaCl₂ at pH 8.2.

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Figure 7 

Hydrodynamic diameter of CeO₂ NPs as function of time in varying concentrations of CaCl₂ ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) 0, 0.05, and 0.1 mM CaCl₂ and pH 8.2; (b) 0.5 and 1 mM CaCl₂ and pH 8.2; (c) 0, 0.05, and 0.1 mM CaCl₂ and pH 10.2; (d) 0.5 and 1 mM CaCl₂ and pH 10.2; (e) 0, 0.05, and 0.1 mM CaCl₂ and pH 12.2; and (f) 0.5 and 1 mM CaCl₂, and pH 12.2.

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Figure 8 

Zeta potential of CeO₂ NPs as function of time in varying concentrations of CaCl₂ ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) 0, 0.05, and 0.1 mM CaCl₂ and pH 8.2; (b) 0.5 and 1 mM CaCl₂ and pH 8.2; (c) 0, 0.05, and 0.1 mM CaCl₂ and pH 10.2; (d) 0.5 and 1 mM CaCl₂ and pH 10.2; (e) 0, 0.05, and 0.1 mM CaCl₂ and pH 12.2; and (f) 0.5 and 1 mM CaCl₂ and pH 12.2.

The hydrodynamic diameter of CeO₂ NPs was greater at 1 mM CaCl₂ than at other CaCl₂ concentrations (Figures 7(a) and 7(b)). At ∼1 hr, the average hydrodynamic diameter at 1 mM CaCl₂ (922.91 ± 89.41 nm) was 721.97% (112.28 ± 2.11 nm), 763.99% (106.82 ± 2.29 nm), 615.05% (129.07 ± 5.38 nm), and 426.62% (175.25 ± 11.86 nm) greater than at 0, 0.05, 0.1, and 0.5 mM CaCl₂, respectively (Table S3). This finding indicates that CeO₂ NPs are unstable at 1 mM CaCl₂. The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 8.2 and in the presence of monovalent CaCl₂ salt at various concentrations is presented in Figures S3(a) and S3(b). The same prediction can be demonstrated by and size distribution. When the concentration of CaCl₂ was less than 1 mM, was very small (0.01 nm/s) (Table 2). At 1 mM CaCl₂, increased to 0.08 nm/s.

However, the prediction from the result of hydrodynamic diameter differed little from the prediction from the net energy analysis, possibly affected by the values of both the zeta potential and IS. The net energy barrier disappeared at 0.5 and 1 mM of CaCl₂ (Figure 6(d)). This absence indicates that the CeO₂ NPs are stable when the concentration of CaCl₂ is less than 0.5 mM, but unstable when the concentration of CaCl₂ reaches 0.5 and 1 mM.

3.3.2. Uncharged CeO₂ NPs in the Presence of CaCl₂ (pH = 10.2 = pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 10.2 and in the presence of divalent CaCl₂ salt at various concentrations are shown in Figures 7(c) and 7(d) and Figures 8(c) and 8(d). At pH 10.2 (pH = pH_PZC), the average zeta potential of CeO₂ NPs in the presence of CaCl₂ varied about 0 mV (Figures 8(c) and 8(d) and Table S2), which indicates that CeO₂ NPs are unstable and the surface charge of CeO₂ NPs is independent of time at pH = pH_PZC. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 10.2 in concentration of CaCl₂ 0, 0.05, 0.1, 0.5, and 1 mM are −0.76, −0.62, −0.83, −0.73, and −0.40, respectively. The unstabilty prediction of CeO₂ NPs at pH = pH_PZC is also demonstrated by the results of hydrodynamic diameter, , size distribution, and net energy barrier. At pH 10.2, hydrodynamic diameters under all the CaCl₂ concentrations grew at very high rates (Figures 7(c) and 7(d)). At ∼1 hr, the average hydrodynamic diameters were 1,670.11 ± 229.21 nm, 1,318.10 ± 125.18 nm, 1,690.85 ± 210.69 nm, 1,845.22 ± 175.04 nm, and 3,568.23 ± 552.86 nm, at 0, 0.05, 0.1, 0.5, and 1 mM of CaCl₂, respectively (Table S3). For example, the hydrodynamic diameter of CeO₂ NPs at 0.1 mM CaCl₂ increased by 167.77% at ∼1 hr (1,690.85 ± 210.69 nm) when compared to the hydrodynamic diameter during the first 100 s (631.45 ± 49.47 nm). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 10.2 and in the presence of monovalent CaCl₂ salt at various concentrations is presented in Figures S3(c) and S3(d). in the presence of CaCl₂ was 0.25–0.50 nm/s at pH 10.2 (Table 2), which was greater than at pH 8.2 and 12.2. No energy barrier was present at pH 10.2 (Figure 6(e)).

3.3.3. Negatively Charged CeO₂ NPs in the Presence of CaCl₂ (pH = 12.2 > pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 12.2 and in the presence of divalent CaCl₂ salt at various concentrations are shown in Figures 7(e) and 7(f) and Figures 8(e) and 8(f), respectively. At pH 12.2 (pH > pH_PZC), the average zeta potential in 1 hr was approximately −30 mV at 0, 0.05, 0.1, and 1 mM CaCl₂ (Figures 8(e) and 8(f) and Table S2). However, the zeta potential in 1 hr was −5.65 ± 20.67 mV at 0.5 mM CaCl₂. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 12.2 in concentration of CaCl₂ 0, 0.05, 0.1, 0.5, and 1 mM are −0.64, −0.86, −0.25, 0.13, and −0.91, respectively. This finding indicates that although CeO₂ NPs were unstable at 0.5 mM CaCl₂, they remained stable at other CaCl₂ concentrations. The same prediction was obtained from the results of the other three parameters, which are the hydrodynamic diameter, , and the size distribution.

At pH 12.2 (pH > pH_PZC), the growth rates were low for all CaCl₂ concentrations, except for 0.5 mM of CaCl₂ (Figures 7(e) and 7(f)). At ∼1 hr, the average hydrodynamic diameters were 89.09 ± 1.56 nm, 99.74 ± 2.37 nm, 182.14 ± 5.52 nm, 1,700.32 ± 287.78 nm, and 101.85 ± 2.72 nm under 0, 0.05, 0.1, 0.5, and 1 mM of CaCl₂, respectively (Table S3). This finding demonstrates that CeO₂ NPs in the presence of CaCl₂ are stable at pH > pH_PZC, except for 0.5 mM of CaCl₂. For example, the hydrodynamic diameter of CeO₂ NPs at 1 mM CaCl₂ increased only by 2.64% at ∼1 hr (101.85 ± 2.72 nm) when compared to the hydrodynamic diameter during the first 100 s (99.23 ± 5.87 nm). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 12.2 and in the presence of monovalent CaCl₂ salt at various concentrations is presented in Figures S3(e) and S3(f). The was very low (<0.02 nm/s) at 0, 0.05, 0.1, and 1 mM CaCl₂ (Table 2). However, was 0.47 nm/s at 0.5 mM CaCl₂. The values of energy barriers were 6.60 and 0.51 kJ, for 0 and 0.05 mM CaCl₂, respectively (Figure 6(f)). However, the prediction from all the parameters above differed from the prediction from the result of the net energy barrier at 0.1 and 1 mM of CaCl₂. No energy barrier was found for 0.1, 0.5, and 1 mM of CaCl₂.

3.4. Impact of pH and HA on Stability of CeO₂ NPs

The effects of HA and cation (Na⁺ and Ca²⁺) concentrations on CeO₂ NPs stability in three pH domains representing the three electrostatic states (i.e., pH = 8.2 < pH_PZC, pH = 10.2 = pH_PZC, and pH = 12.2 > pH_PZC) were examined via ELS and DLS measurements of the nanoparticle surface charges and sizes.

3.4.1. Positively Charged CeO₂ NPs in the Presence of HA (pH = 8.2 < pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 8.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations, are given in Figures 9(a) and 10(a) and Figures 11(a) and 11(b), respectively. At pH 8.2 (pH < pH_PZC), at 0–60 min, in the presence of HA, when the concentration of NaCl increased from 0 to 1 mM, the average zeta potential of CeO₂ NPs increased from 7.98 ± 20.71 mV to 13.03 ± 9.68 mV (Figures 11(a) and 11(b) and Table S2). However, it increased initially when the CaCl₂ concentration increased from 0 to 0.01 mM, then decreased when concentration of CaCl₂ increased to 1 mM. The average zeta potential of CeO₂ NPs was less than 20 mV and close to 0 mV, which indicates the aggregation of CeO₂ NPs in the presence of HA and CaCl₂ at pH 8.2. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 8.2 in the presence of HA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are 0.01, −0.66, −0.35, −0.17, and −0.88, respectively. However, the other parameters did not show the same prediction as zeta potential measurements.

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Figure 9 

Hydrodynamic diameter measurements of CeO₂ NPs in the presence of HA (5 mg/L) or FA (0.14 mg/L) as a function of time, and NaCl concentration ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) HA, NaCl, and pH 8.2; (b) HA, NaCl, and pH 10.2; (c) HA, NaCl, and pH 12.2; (d) FA, NaCl, and pH 8.2; (e) FA, NaCl, and pH 10.2; and (f) FA, NaCl, and pH 12.2.

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Figure 10 

Hydrodynamic diameter measurements of CeO₂ NPs in the presence of HA (5 mg/L) or FA (0.14 mg/L) as a function of time, and CaCl₂ concentration ranging from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2: (a) HA, CaCl₂, and pH 8.2; (b) HA, CaCl₂, and pH 10.2; (c) HA, CaCl₂, and pH 12.2; (d) FA, CaCl₂, and pH 8.2; (e) FA, CaCl₂, and pH 10.2; and (f) FA, CaCl₂, and pH 12.2.

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Figure 11 

Zeta potential of CeO₂ NPs as a function of time in varying concentrations of NaCl or CaCl₂ in the presence of HA at pH 8.2, pH 10.2, and pH 12.2: (a) HA, NaCl, and pH 8.2; (b) HA, CaCl₂, and pH 8.2; (c) HA, NaCl, and pH 10.2; (d) HA, CaCl₂, and pH 10.2; (e) HA, NaCl, and pH 12.2; and (f) HA, CaCl₂, and pH 12.2.

At pH < pH_PZC, in the presence of HA, the average hydrodynamic diameter of CeO₂ NPs was smaller at 0.1 mM of NaCl/CaCl₂ than at both 0 and 1 mM of NaCl/CaCl₂ (Figures 9(a) and 10(a)). For instance, at ∼1 hr, the hydrodynamic diameter of CeO₂ NPs at 1 mM CaCl₂ (1,397.18 ± 209.50 nm) was 649.24% greater than at 0.1 mM (186.48 ± 8.47 nm) (Table S4). At ∼1 hr, the average hydrodynamic diameter in the presence of HA was larger than or close to it in the absence of HA. For instance, at ∼1 hr, the hydrodynamic diameter of CeO₂ NPs at 0.1 mM NaCl in the presence of HA (183.78 ± 7.76 nm) was 65.05% greater than in the absence of HA (111.35 ± 2.83 nm). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 8.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations, is presented in Figures S4(a), S4(b), S4(g), and S4(h).

First decreased then increased when the electrolyte concentration increased from 0 to 1 mM. It was smaller at 0.1 and 1 mM NaCl/CaCl₂ than at 0 mM of NaCl/CaCl₂ (Table 2). In the presence of HA, at 0.1 mM of salt, in the presence of NaCl was smaller than in the presence of CaCl₂. In the presence of HA, at 1 mM of salt, in the presence of NaCl was greater than in the presence of CaCl₂.

No energy barrier was found at 1 mM of NaCl/CaCl₂ in the presence of HA (Figures 12(a) and 12(d)). The values of energy barriers were 4.23, 2.83, and 2.96 kJ at 0 mM of NaCl, 0.1 mM of NaCl, and 0.1 mM of CaCl₂, respectively (Figures 12(a) and 12(d)). This finding indicates that at pH < pH_PZC, CeO₂ NPs are unstable at 1 mM of NaCl/CaCl₂.

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Figure 12 

Net energy versus interparticle distance according to DLVO theory, showing the influence of HA (5 mg/L) and representative monovalent electrolyte NaCl or divalent electrolyte CaCl₂ concentrations varying from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2, for CeO₂ NPs suspended in aqueous systems: (a) HA, NaCl, and pH 8.2; (b) HA, NaCl, and pH 10.2; (c) HA, NaCl, and pH 12.2; (d) HA, CaCl₂, and pH 8.2; (e) HA, CaCl₂, and pH 10.2; and (f) HA, CaCl₂, and 12.2. Calculation of DLVO theory includes the contributions of the van der Waals and electric double layer (EDL) interactions. The net energy (Net) shown is the sum of EDL repulsion and van der Waals attraction between two approaching particles.

3.4.2. Uncharged CeO₂ NPs in the Presence of HA (pH = 10.2 = pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations, are presented in Figures 9(b) and 10(b) and Figures 11(c) and 11(d), respectively. At pH 10.2 (pH = pH_PZC), the absolute average zeta potential of CeO₂ NPs varied about 0 mV and was less than 22 mV (Figures 11(c) and 11(d) and Table S2). The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 10.2 in the presence of HA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are 0.89, −0.97, 0.72, −0.87, and −0.75, respectively. CeO₂ NPs were unstable in the presence of HA at pH = pH_PZC. This conclusion can also be predicted by the results of hydrodynamic diameter, , and size distribution. The average hydrodynamic diameter of CeO₂ NPs grew at high rates and reached the micron level in the presence of HA after approximately 1 hr of incubation (Figures 9(b) and 10(b)). For example, in the presence of HA, the hydrodynamic diameter of CeO₂ NPs at 0.1 mM CaCl₂ increased by 236.61% at ∼1 hr (1,584.86 ± 223.48 nm) when compared to the hydrodynamic diameter during the first 100 s (470.83 ± 39.34 nm) (Table S4). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations, is presented in Figures S4(c), S4(d), S4(i), and S4(j).

The was in the range of 0.37–0.54 nm/s under all concentrations of NaCl/CaCl₂ in the presence of HA (Table 2).

No energy barrier was found at 0.1 mM of NaCl and at 0.1 and 1 mM of CaCl₂ (Figures 12(b) and 12(e)). The values of energy barriers were 6.87 and 4.70 kJ at 0 and 1 mM of NaCl (Figure 12(b)). This result indicates that at pH = pH_PZC, CeO₂ NPs are unstable in the presence of HA at 0.1 mM of NaCl, and at 0.1 and 1 mM of CaCl₂. This prediction is different from the prediction from the results of the other parameters above.

3.4.3. Negatively Charged CeO₂ NPs in the Presence of HA (pH = 12.2 > pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 12.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations, are presented in Figures 9(c) and 10(c) and Figures 11(e) and 11(f), respectively. At pH 12.2 (pH > pH_PZC), the average zeta potential of CeO₂ NPs was less than −30 mV during the 1 hr incubation in the presence of HA and NaCl/CaCl₂ (Figures 11(e) and 11(f) and Table S2), which indicates that CeO₂ NPs are stable in the presence of HA at pH > pH_PZC. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 12.2 in the presence of HA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are 0.31, 0.11, −0.24, −0.43, and 0.35, respectively.

The growth of hydrodynamic diameter of CeO₂ NPs was slow for all suspensions in the presence of HA over 1 hr (Figures 9(c) and 10(c)). At ∼1 hr, the average hydrodynamic diameter of CeO₂ NPs was less than 120 nm. For example, in the presence of HA, the hydrodynamic diameter of CeO₂ NPs at 1 mM NaCl increased only by 20.74% at ∼1 hr (93.86 ± 1.91 nm) when compared to the hydrodynamic diameter during the first 100 s (77.74 ± 2.31 nm) (Table S4). This finding also indicates the stability of CeO₂ NPs in the presence of HA at pH > pH_PZC with the same prediction obtained from the other parameters. The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 12.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRHA, all at various concentrations is presented in Figures S4(e), S4(f), S4(k), and S4(l). The was less than 0.02 nm/s at pH 12.2 for all suspensions in the presence of HA (Table 2). The net energy barriers were 10.4, 27.3, and 18.4 kJ, at 0, 0.1, and 1 mM of NaCl, respectively (Figure 12(c)). The net energy barriers were 20.5 and 4.16 at 0.1 and 1 mM of CaCl₂ (Figure 12(f)).

3.5. Impact of pH and FA on Stability of CeO₂ NPs

The effects of FA and cation (Na⁺ and Ca²⁺) concentrations on CeO₂ NPs stability in three pH domains representing the three electrostatic states—pH = 8.2 < pH_PZC, pH = 10.2 = pH_PZC, and pH = 12.2 > pH_PZC—were examined by measuring the surface charge and size of nanoparticles using ELS and DLS methods.

3.5.1. Positively Charged CeO₂ NPs in the Presence of FA (pH = 8.2 < pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 8.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, are presented in Figures 9(d) and 10(d) and Figures 13(a) and 13(b), respectively. At pH < pH_PZC, in the presence of FA, the average zeta potential of CeO₂ NPs was less than −35 mV at 0, 0.1, and 1 mM of NaCl and at 0.1 mM of CaCl₂ (Figures 13(a) and 13(b) and Table S2). However, the average zeta potential of CeO₂ NPs ranged from −23 to −26 mV at 1 mM of CaCl₂. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 8.2 in the presence of FA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are −0.58, −0.21, −0.11, −0.67, and 0.56, respectively.

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Figure 13 

Zeta potential of CeO₂ NPs as a function of time in varying concentrations of NaCl or CaCl₂ in the presence of FA at pH 8.2, pH 10.2, and pH 12.2: (a) FA, NaCl, and pH 8.2; (b) FA, CaCl₂, and pH 8.2; (c) FA, NaCl, and pH 10.2; (d) FA, CaCl₂, and pH 10.2; (e) FA, NaCl, and pH 12.2; and (f) FA, CaCl₂, and pH 12.2.

At ∼1 hr, the average hydrodynamic diameter of CeO₂ NPs in the presence of FA was less than 110 nm (Figures 9(d) and 10(d)). For example, in the presence of FA, the hydrodynamic diameter of CeO₂ NPs at 0.1 mM NaCl only increased by 9.12% at ∼1 hr (101.87 ± 3.36 nm) when compared to the hydrodynamic diameter during the first 100 s (93.36 ± 1.95 nm) (Table S5). This result indicates that FA stabilizes CeO₂ NPs at pH < pH_PZC. The same prediction is obtainable from the results of size distribution (Figure S5). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 8.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, is presented in Figures S5(a), S5(b), S5(g), and S5(h). Thus was less than 0.02 nm/s at pH 8.2 for all suspensions in the presence of FA (Table 2).

However, at 1 mM of CaCl₂, the result of the net energy barrier was different from the other parameters in the presence of FA (Figures 14(a) and 14(d)). Here, energy barriers existed at all salt concentrations, except at that level (Figures 14(a) and 14(d)).

(a)

(b)

(c)

(d)

(e)

(f)

Figure 14 

Net energy versus interparticle distance according to DLVO theory, showing the influence of FA (0.14 mg/L) and representative monovalent electrolyte NaCl or divalent electrolyte CaCl₂ concentrations varying from 0 to 1 mM at pH 8.2, pH 10.2, and pH 12.2, for CeO₂ NPs suspended in aqueous systems: (a) FA, NaCl, and pH 8.2; (b) FA, NaCl, and pH 10.2; (c) FA, NaCl, and pH 12.2; (d) FA, CaCl₂, and pH 8.2; (e) FA, CaCl₂, and pH 10.2; and (f) FA, CaCl₂, and 12.2. Calculation of DLVO theory includes the contributions of the van der Waals and electric double layer (EDL) interactions. The net energy (Net) shown is the sum of EDL repulsion and van der Waals attraction between two approaching particles.

3.5.2. Uncharged CeO₂ NPs in the Presence of FA (pH = 10.2 = pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, are presented in Figures 9(e) and 10(e) and Figures 13(c) and 13(d), respectively. At pH = pH_PZC, the average zeta potential of CeO₂ NPs was less than −30 mV in the presence of FA and NaCl (Figures 13(c) and 13(d) and Table S2). The average zeta potential of CeO₂ NPs was −15 to −32 mV in the presence of FA and CaCl₂. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 10.2 in the presence of FA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are 0.07, −0.20, 0.77, 0.23, and 0.17, respectively.

The results of hydrodynamic diameter, , and size distribution indicated the stability of CeO₂ NPs in the presence of FA at pH 10.2. At ∼1 hr, the average hydrodynamic diameter of CeO₂ NPs in the presence of FA was less than 105 nm (Figures 9(e) and 10(e)). For example, in the presence of FA, the hydrodynamic diameter of CeO₂ NPs at 1 mM CaCl₂ increased only by 25.28% at ∼1 hr (97.02 ± 2.96 nm) when compared to the hydrodynamic diameter during the first 100 s (77.44 ± 4.25 nm) (Table S5). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, is presented in Figures S5(c), S5(d), S5(i), and S5(j). The was less than 0.02 nm/s at pH 10.2 for all suspensions in the presence of FA (Table 2).

However, at pH 10.2, the net energy barrier existed in the presence of FA and NaCl/CaCl₂, except that the net energy barrier was 0 kT at 1 mM CaCl₂ (Figures 14(b) and 14(e)). This finding indicates that CeO₂ NPs are not stable at 1 mM of CaCl₂.

3.5.3. Negatively Charged CeO₂ NPs in the Presence of FA (pH = 12.2 > pH_PZC)

The aggregation profiles and average zeta potentials of CeO₂ NPs at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, are presented in Figures 9(f) and 10(f) and Figures 13(e) and 13(f), respectively. At pH 12.2 (pH > pH_PZC), at 0–1 hr, the average zeta potential varied from −24 to −38 mV for all suspensions in the presence of FA (Figures 13(e) and 13(f) and Table S2), a finding that shows the stability of CeO₂ NPs at pH 12.2. The Pearson’s correlation coefficients for the ξ potential values of CeO₂ NPs and time at pH 12.2 in the presence of FA in concentration of NaCl 0, 0.1, and 1 mM, CaCl₂ 0.1 and 1 mM are 0.78, −0.96, 0.46, −0.33, and −0.72, respectively. The same prediction is obtainable from results of all the parameters below as the result of zeta potential measurements.

The average hydrodynamic diameter of the CeO₂ NPs grew at low rates in the presence of FA, and the average hydrodynamic diameter was less than 120 nm (Figures 9(f) and 10(f)). For example, in the presence of FA, the hydrodynamic diameter of the CeO₂ NPs at 1 mM NaCl only increased by 14.36% at ∼1 hr (94.35 ± 3.49 nm) when compared to the hydrodynamic diameter during the first 100 s (82.50 ± 1.77 nm) (Table S5). The particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr at pH value of 10.2 and in the presence of monovalent NaCl or divalent CaCl₂ salts as well as SRFA, all at various concentrations, is presented in Figures S5(e), S5(f), S5(k), and S5(l). The was less than 0.02 nm/s at pH 12.2 for all suspensions in the presence of FA (Table 2). In the presence of FA, net energy barriers existed under all electrolyte concentrations at pH 12.2 (Figures 14(c) and 14(f)).

4. Discussion

The stability of NPs derives from their ability to remain unchanged over time under certain conditions. The importance of studying the stability of the CeO₂ NPs is that their stability affects their transport, retention, bioavailability, and toxicity while posing potential risks to the environment and human health.

4.1. Effect of pH on the Stability of CeO₂ NPs (Homoaggregation)

The pH is one of the water chemistry parameters that affect NPs aggregation by altering the charge on the surface of NPs [69]. At both low and high pH levels, the aggregation was extremely constrained but was enhanced as the pH approached pH_PZC [92].

Homoaggregation, a term which refers to the aggregation of two similar CeO₂ NPs [93], was affected by the changing pH conditions presented in this study. The stability of CeO₂ NPs in aqueous systems differed at various pH domains. Although the CeO₂ NPs were stable at pH < pH_PZC (except for 1 mM of NaCl/CaCl₂) and at pH > pH_PZC (except for 0.5 mM CaCl₂), the aggregation was enhanced at pH = pH_PZC.

The phenomena and explanations of the aggregation and stability of CeO₂ NPs in three pH domains (pH < pH_PZC, pH = pH_PZC, and pH > pH_PZC) in our research are summarized below.

At pH < pH_PZC, the average hydrodynamic diameter of CeO₂ NPs at 1 mM of salt during the first 100 s (approximately 165 nm) was much lower than that at ∼1 hr (above 1,000 nm) (Figures 4(b) and 7(b)). The CeO₂ NPs were stable at either no or low concentration of electrolytes (<1 mM) and were unstable at higher concentration (1 mM) of electrolytes. This conclusion was demonstrated by the results of DLVO theory analysis on the net energy interaction among CeO₂ NPs: the repulsive energy barriers were found at 0–0.5 mM NaCl/CaCl₂ but not shown at 1 mM NaCl/CaCl₂ (Figures 6(a) and 6(d)).

One possible explanation is that at pH < pH_PZC, surface protonation occurs when the protons adsorb to the hydration layer capped on CeO₂ NPs, thus preventing the further aggregation of CeO₂ NPs. The occurrence of suspensions in either the absence of electrolyte or in the presence of low electrolyte results in a dominance of the repulsive interaction from the large distance separations between CeO₂ NPs. However, the increase in electrolyte concentration led to the compression of the double layer of CeO₂ NPs and reduced the stability of CeO₂ NPs. The strongly positively charged surface of CeO₂ NPs at pH < pH_PZC resulted in the predominated role of electrostatic repulsion over attraction; thus the stability of CeO₂ NPs was enhanced. Similarly, due to the same mechanism above, the low pH enhanced the CeO₂ NP mobility, a result that is attributable to the increased stability of the CeO₂ NPs [40]. The similar influence of electrolyte on the stability of CeO₂ NPs was also noted in that the CeO₂ NPs were stable at low NaCl concentrations (<10 mM) and unstable at higher NaCl concentrations (>10 mM). Further, the positive zeta potentials determined for three different synthesized CeO₂ NPs had isoelectric points at pH 5.8, 7.6, and 6.5 [52]. The uncoated CeO₂ NPs exhibited positive zeta potentials at pH less than its pH_PZC (pH 6.8 ± 0.1) [44]. In addition, at pH 4.5 (i.e., pH < pH_PZC), TiO₂ NPs nanoaggregate size was substantially greater, and the size distribution became broader at a higher IS (0.165 M) in the presence of NaCl/CaCl₂ [79]. However, no obvious aggregation was evident at a lower IS (0.0045 M).

At a pH = pH_PZC, the largest aggregation sizes were obtained. This finding was in correlation to the close-to-zero average zeta potentials (Table S2), the large values of the average hydrodynamic diameter of CeO₂ NPs (1,300–3,600 nm) at ∼1 hr (Figures 4(c) and 4(d), Figures 7(c) and 7(d), and Table S3), and the lack of a repulsive energy barrier (0 kT) (Figures 6(b) and 6(e)). This outcome suggests that at pH = pH_PZC, the electrostatic attraction is predominant over the repulsion, and destabilization behavior exists in the suspensions. Similarly, a pH level approximate to the pH_PZC (pH 6.5) increased the protons that then reacted with the functional groups (e.g., carboxyl groups) with a negative surface charge, thus reducing the CeO₂ NP surface charges to the benefit of the aggregation [40]. However, in their study, Oriekhova and Stoll [44] reported that the average hydrodynamic diameter of CeO₂ NPs was near 2,000 nm at pH = pH_PZC. Quik et al. [63] also noted a discrepancy in pH_PZC that was equal to pH 8 in 10 mg/L CeO₂ NP samples. The differences in these studies of the pH_PZC of CeO₂ NP are perhaps due to the purchasing of nanoparticles created with various synthesizing techniques, and a difference in the concentrations, solution/suspension conditions, and mediums used in the experiments. Similarly, the pH_PZC values for the TiO₂ NPs also varied due to the technique used and constituents introduced during the synthesis [94]. In addition, the enhanced aggregation behavior of other types of NPs also occurred when the surface charge of the NPs was near the pH_PZC. When the pH reached the range of 5–8, large TiO₂ nanoaggregates formed, suggesting neutralization as the primary catalyst in nanoaggregate formation [95].

The enhanced aggregation behavior at pH = pH_PZC in our research is explained by the occurrence of the surface charge neutralization of CeO₂ NPs from the interaction of hydroxyl ions with the CeO₂ NP surfaces. The condensation reaction, which is the inverse of hydrolysis, is expressed below [96]:with M-OH₂⁺ indicating the products after the interactions among CeO₂ NPs—water interface and H⁺.

The hydrolysis reaction mechanism was used to demonstrate the impact of pH on aggregation of TiO₂ NPs by adding FeCl₃ [97] and on aggregation of iron oxide nanoparticles [98]. The hydrolysis complex structure of C₆₀ NPs was also considered as the reason for the initial decrease in pH caused by the decrease of the absolute zeta potential [99]. In addition, at pH = pH_PZC, the maximum aggregation sizes were observed for TiO₂, titanate nanotubes, and titanate nanotubes-TiO₂ (800–1,300 nm) [100] and for ZnO NPs (1,802 nm) [101]. The destabilization was also present for ZnO NPs at pH 9, which was close to pH_PZC (9.2) [102].

At pH > pH_PZC, the surface deprotonation continues processing the large amount of OH⁻ interacting with surface of CeO₂ NPs. The large amount of OH⁻ adsorbed on the CeO₂ NP surfaces in turn yields strongly negatively charged and stable CeO₂ NPs (close to or less than −30 mM) at pH 12.2 (Table S2). Small CeO₂ nanoaggregates were present after an hour incubation, which was less than 190 nm, as shown in Figures 4(e) and 4(f) and Figures 7(e) and 7(f). Similarly, the homoaggregation of CeO₂ NPs at pH 11, which was larger than its pH_PZC, was enhanced with the increase of NaCl and CaCl₂ until they reached their CCC values [52]. In addition, the negatively charged TiO₂ NPs (with an approximate zeta potential of −55 mV) were stable (with an average hydrodynamic diameter near 120 nm) at a high pH [95]. In correspondence to the small CeO₂ nanoaggregates, we obtained small aggregate rates (<0.02 nm/s) in our study. A low aggregation rate (i.e., close to 0 nm/min) was also obtained for negatively charged TiO₂ NPs (with a zeta potential of −38.7 mV) at pH > pH_PZC [103]. At pH 11–12 (pH > pH_PZC = 7.8), small nanoaggregates (<50 nm) formed at a low concentrations of iron oxide NPs (<50 mg/L) [104]. Moreover, at pH 9 (pH > pH_PZC), the Fe₃O₄ NPs were more unstable in the presence of divalent cations (e.g., Ca²⁺) than in the presence of monovalent cations (Na⁺) [105]. The size of Ag NPs was below 100 nm—at pH 8 because of the negative zeta potential from the negatively charged OH⁻ [106].

4.2. Effect of HA and FA on the Stability of CeO₂ NPs (Heteroaggregation)

Although the HA stabilized CeO₂ NPs under pH > pH_PZC, the aggregation was enhanced at pH = pH_PZC and in the presence of 0 and 1 mM of NaCl/CaCl₂ at pH < pH_PZC. Further, the FA stabilized CeO₂ NPs at all three pH levels (pH>, = , and <pH_PZC) and under all electrolyte concentrations (0–1 mM of NaCl or CaCl₂).

Heteroaggregation is a phenomenon that is characterized by dissimilar particle aggregation [93], the mechanisms of which between SRHA or SRFA and CeO₂ NPs differ from the mechanisms of homoaggregation between CeO₂ NPs. The adsorbed NOM on the NP surface was not only found to increase electrostatic repulsive forces between NPs but also cause steric forces, thus reducing NP aggregation [107–109]. In addition, the NOM sorption on the particles can facilitate the cation bridging and enhance both NOM and particle heteroaggregation (e.g., graphene oxide) [110]. Moreover, NOM can also neutralize the positive surface charge of NPs and enhance the aggregation [94]. Therefore, the complex nature of the heteroaggregation state and stability of CeO₂ NPs in the presence of HA or FA was also affected by both the pH and ionic composition and concentration.

4.2.1. Stable CeO₂ NPs with HA

The CeO₂ NPs under study here exhibited stability after a 1 hr incubation period in terms of the zeta potential (with the absolute value either near, or larger than, 30 mV). Further, the average hydrodynamic diameter was less than 120 nm for CeO₂ NPs in the presence of HA at all NaCl/CaCl₂ concentrations (0–1 mM) at pH 12.2 (Figures 9(c) and 10(c) and Table S2). The dispersion of CeO₂ NPs by HA at pH > pH_PZC may account for the electrosteric stabilization, which is caused by the steric repulsion between the complexes of CeO₂ NPs and negatively charged HA. This finding agrees with the studies involving the use of NOM and particularly HA in CeO₂ NP stabilization [47, 49] and other types of NPs [68, 69, 97, 111, 112]. Both the electrostatic and steric stabilization mechanisms were also determined as catalysts for enhancing the NP stabilization [101, 113–116]. Further, the low average hydrodynamic diameter (<1,000 nm) of CeO₂ NPs in the presence of NOM was present at a high pH and a low IS [49]. HA was an important catalyst in stabilizing the TiO₂ NPs and forming small-size nanoaggregates (of approximately 250 nm) in natural waters due to the effects of both the electrostatic and steric repulsion [69]. The TiO₂ NPs surface was covered by the HA molecules, and the aggregation was inhibited at a high pH [97]. We also found that the stability of CeO₂ NPs was higher in the presence of HA at pH > pH_PZC than at other pHs in our study. At a pH > pH_PZC, the increase of HA (from 0 to 1 mg/L) decreased the zeta potentials from <−20 to −25.8 mV and −30 to −35.65 mV for both the anatase and rutile TiO₂ NPs, with a 120 nm size at 1 mg/L HA, indicating that HA stabilized these TiO₂ NPs [94]. However, at pH < pH_PZC, the increase of HA (0–1 mg/L) neutralized the positively charged TiO₂ NPs, resulting in close to 0 mV zeta potentials (−4 mV for the anatase TiO₂ NPs and −9 mV for the rutile TiO₂ NPs) and unstable TiO₂ NPs. In another example, at basic pH (pH 8 and 10), Ag NPs were stable with a small hydrodynamic diameter (around 120 nm) in the presence of SRHA [117].

Inhibited aggregation was also in the presence of HA at 0.1 mM of salt (NaCl or CaCl₂) at pH 8.2 (Figures 9(a) and 10(a) and Table S2) due to the low values of CeO₂ NPs (less than 0.05 nm/s). The stability of CeO₂ NPs is higher than predicted by the values of average zeta potential at pH < pH_PZC, a finding that indicates that steric stabilization inhibits the growth of CeO₂ nanoaggregates [17].

4.2.2. Unstable CeO₂ NPs with HA

Four concepts indicated the instability of the CeO₂ NPs in the presence of HA under 0 and 1 mM of salt (NaCl or CaCl₂) at pH 8.2 and in the presence of HA at all NaCl/CaCl₂ concentrations (0–1 mM) at pH 10.2. These are (i) an average hydrodynamic diameter larger than 1,500 nm (Figures 9(a) and 9(b) and Figures 10(a) and 10(b)); (ii) an absolute average zeta potential value less than 20 mV (Table S2); (iii) an aggregation rate larger than 0.35 nm/s (Table 2); and (iv) the values of repulsive energy barriers smaller than 7 kT (Figures 12(a), 12(b), 12(d), and 12(e)). A summary of the possible mechanisms and explanations for the unstable CeO₂ NPs at different pH and salt concentrations is below.

The predominance of protons that neutralize the CeO₂ NP surface charges may cause the CeO₂ NP instability in the presence of HA at all Na⁺ or Ca²⁺ concentrations (0–1 mM) at pH = pH_PZC, though negatively charged HA exists in the systems. This instability also indicates the insufficiency of 5 mg/L HA in stabilizing the CeO₂ NPs (25 mg/L) at pH 10.2.

HA adsorption on the partial surface of CeO₂ NPs may be the cause of CeO₂ NP destabilization at pH < pH_PZC in the presence of HA and the absence of salt, thus increasing the CeO₂ NP hydrophobicity while promoting aggregation, also stated by Ghosh et al. [68] and Chen et al. [94]. HA neutralized the positive charges on the CeO₂ NP surfaces at pH < pH_PZC, which decreased the repulsive forces between the NPs and increased aggregation. HA can also lead to charge neutralization and NP destabilization [118, 119]. Similar phenomena and mechanisms were also identified for magnetite NPs [120]. In their study, Hu et al. [120] determined that HA increased the magnetite NP aggregation under changing pH conditions. In the presence of low HA concentrations (2 and 3 mg/L), the aggregation was enhanced at low pHs. The enhanced aggregation of CeO₂ NPs at 1 mM NaCl, as compared to the 0.1 mM NaCl, is from the compression of EDLs caused by an increase in IS. Similarly, Quik et al. [63] found that the steric repulsion of NOM had little effect upon the stability of the CeO₂ NPs. The aggregation of CeO₂ NPs was enhanced (i.e., the average hydrodynamic diameter increased from 173 to 253 nm) with the increase of electrolytes content in NOMs [63]. Also, a study of transport and deposition of SRHA-formed Ag NPs noted a marked instability of Ag NPs at a low acidic pH and a high IS [117]. In their study of CeO₂ NP destabilization, Li and Chen [47] noted an enhanced CeO₂ NP aggregation in the presence of HA, at a high CaCl₂ concentration of (80 mM) and a 5.7 pH (pH < pH_PZC). This type of CeO₂ NP aggregation is due to the bridging effect. In the bridge effect, interactions between HA and Ca²⁺ result in the formation of large HA aggregates via Ca²⁺ complexation which can bridge NPs and induce their aggregation [121]. In our study, we noted the enhanced aggregation in the presence of HA at 1 mM of Ca²⁺ at pH 8.2. Here the average zeta potential was approximately 10 mV (Table S2), the average hydrodynamic diameter was larger than 1,300 nm after a 1 hr incubation period (Figure 10(a)), and the slope of growth of CeO₂ NPs in the hydrodynamic diameter was larger than 0.10 nm/s (Table 2). Similarly, the bridging attraction between CeO₂ NPs induced by the aggregation of HA via Ca²⁺ was also found when the CeO₂ NPs were in the presence of HA and under high concentrations of Ca²⁺ [47]. The CeO₂ NPs were more unstable at a higher IS in the presence of NOM, and the CeO₂ nanoaggregates were larger than 1,000 nm [49]. Similar aggregation behaviors due to the inter-molecule bridging effect has been found among other types of NPs, such as C₆₀ NPs [121] and Ag NPs [122]. In addition, NPs (e.g., Ag NPs or C₆₀) in the systems of SRHA or SRNOM were more unstable with high divalent cations than with monovalent cations [123, 124]. The bridging effect of NOM and Ca²⁺ increased the aggregation of other types of NPs, such as citrate-coated gold nanoparticles [125], iron oxide NPs [126], and other types of NPs [127]. In another example, the aggregation of silicon NPs was enhanced in the presence of SRHA and Ca²⁺ due to the bridging effect of SRHA with Ca²⁺ overweighing the compression of the EDL between NPs by Ca²⁺ [128].

4.2.3. Stable CeO₂ NPs with FA

In this research, the CeO₂ NPs were very stable in the presence of FA (0.14 mg/L) at all three pHs (8.2, 10.2, and 12.2) and under all electrolyte concentrations (0–1 mM of NaCl or CaCl₂) (Figures 9(d)–9(f) and Figures 10(d)–10(f)). This finding indicates that lower concentration of FA in our study can completely coat the surface of CeO₂ NPs and strongly stabilize CeO₂ NPs due to steric hindrance and increased electrostatic repulsion as indicated by the high absolute zeta potential values. Recent studies on the influence of FA upon the stability of CeO₂ NPs also noted the similar prevention of CeO₂ NP aggregation. Specifically, Oriekhova and Stoll [44] found that CeO₂ NPs were stable at pH 8 (the average hydrodynamic diameter < 210 nm) in the presence of 2 mg/L of FA. They also reported that the negative zeta potential and the small average hydrodynamic diameter (no larger than 220 nm) were due to the formation of CeO₂ NP–FA complexes, as indicated on SEM imagery [45]. Similar stable nanoaggregates were also present among other types of NPs. The increase of surface charge and formation of surface coating caused by FA was also evidenced in the study of cit-AgNPs, where adsorption of FA on cit-AgNPs led to electrosteric stabilization [89]. In the presence of FA and a high IS, the TiO₂ NPs were stabilized [129]. Similarly, the FA adsorption to other types of NPs (e.g., TiO₂ NPs, C₆₀ NPs, Ag NPs, iron oxide NPs) and stabilizing effects of FA on NPs were attributed to steric effects and compression of the diffusive layer of both NPs and FAs [89, 129, 130].

It is known that FA has a lower molecular weight (500–2,000 g/mol) than HA (1,000–10,000 g/mol) with HA in many cases deemed more effective than FA in stabilizing NPs [87, 115, 130–133]. However, in our research, CeO₂ NPs in the presence of FA (0.14 mg/L) exhibited a greater degree of stability than CeO₂ NPs in the presence of HA (5 mg/L) at all three alkaline pH levels. This phenomenon suggests that FA has stronger electrostatic interactions than HA at those alkaline pH levels. The reason may be that FA is more aromatic and thus has increased the phenolic groups ionization and low hydrogen interaction-forming open structures than HA in alkaline pH [134, 135].

5. Conclusion

Our research provides a framework for understanding colloid systems and interfaces stability of engineered nanoparticles, CeO₂ NPs, in aqueous environments. The authors conducted a series of experiments to elucidate the physical and chemical interaction mechanisms that govern the colloidal stability and aggregation kinetics of CeO₂ NPs under the influence of soil and water abiotic factors—pH, ionic composition (monovalent NaCl and divalent CaCl₂ salts), and SRHA and SRFA. Experimental results showed that the solution chemistry affected the colloidal stability and aggregation kinetics of CeO₂ NPs. As a first key finding, this research demonstrated the effects of pH and salt on colloid systems and interfaces stability of CeO₂ NPs. The zeta potential of CeO₂ NPs, with pH_PZC of 10.2, decreased (from positive to negative) with increasing solution pH. The diameter of CeO₂ NP aggregates was ∼1,700 nm in the region of pH_PZC, and decreased with pH at pH < pH_PZC or pH > pH_PZC to ∼100 nm, except at pHs 1, 13, and 14, where it reached ∼1,500–2,250 nm. The impacts of Na⁺ and Ca²⁺ cations and HA and FA on the levels and rates of aggregation were pH-dependent.

Additionally, in the presence of salts, CeO₂ NPs were stable at pH < pH_PZC (except for 1 mM of NaCl/CaCl₂) and pH > pH_PZC (except for 0.5 mM CaCl₂); however, the aggregation was enhanced at pH = pH_PZC, with a diameter of CeO₂ NPs in the ∼1,300–3,600 nm range. Next, HA stabilized CeO₂ NPs under pH > pH_PZC; however, the aggregation was again enhanced at pH = pH_PZC with a diameter of CeO₂ NPs in the ∼1,500–1,900 nm range, and in the presence of 0 and 1 mM of NaCl/CaCl₂ at pH < pH_PZC. FA (0.14 mg/L) showed more efficiency in stabilizing the CeO₂ NPs than HA (5 mg/L) at three pH levels (8.2, 10.2, and 12.2) and under all different electrolyte concentrations (0–1 mM of NaCl or CaCl₂). The diameter of CeO₂ NPs in the presence of FA grew at low rates and was ∼95–115 nm at all three pHs and under all different electrolyte concentrations.

The significant impact of the solution chemistry (i.e., pH, salts, HA, and FA) in which CeO₂ NPs were suspended is linked to the development of interfacial complexation in aqueous environments, this finding is the second key result of this research. Particularly, besides the EDL compression effect by Ca²⁺, between CeO₂ NPs, the ion bridging effect between CeO₂ NPs in Ca²⁺-HA systems was a key controlling mechanism of the stability of CeO₂ NPs. Furthermore, FA inhibited the aggregation of CeO₂ NPs by enhancing the energy barrier, therefore allowing CeO₂ NPs to remain stable at pH values of 8.2, 10.2, and 12.2, and with NaCl and CaCl₂ concentrations ranging from 0 to 1 mM.

The risk assessment for CeO₂ NPs contamination, according to their aggregation and stability state in surface water and in the subsurface soils and aquifers, is of great importance for future regulation and evaluation of CeO₂ NPs waste disposal and applications in products. In rivers and oceans that are abundant with HA or FA, CeO₂ NPs are likely to be suspended and freely transported from one aquatic region to another, posing risks for aquatic animals and human swimming in a CeO₂ NPs-contaminated water. However, under some specific conditions with different solution chemistries, CeO₂ NPs form aggregates, settle at the bottom of the water compartments, become sediment contaminants uptake by sediment-dwelling animals, and enter the food chains. Therefore, additional testing of the CeO₂ NPs aggregation behavior and stability and other possible interactions under different solution chemistry is required to develop a data collection bank for CeO₂ NPs for the accurate and precise analysis, interpretation, and prediction for environmental risk assessment and protection.

Data Availability

The datasets generated during and/or analyzed during the current study are available from the corresponding author upon reasonable request.

Disclosure

This article is based on a chapter from the Ph.D. dissertation titled “Stability and Toxicity of Cerium Oxide Nanoparticles in Water-Soil-Plant Systems” by Linlin Mu, Clemson University, Clemson, SC, United States (2019) [136].

Conflicts of Interest

The authors declare that they have no conflicts of interest.

Authors’ Contributions

L.M. and C.J.G.D. contributed to conceptualization, data curation, and methodology. L.M., M.G., and C.J.G.D. contributed to formal analysis. L.M., P.C.B., and C.J.G.D. contributed to investigation. L.M. contributed to writing—original draft preparation. L.M., M.G., P.C.B., and C.J.G.D. contributed to writing—review and editing. C.J.G.D. contributed to funding acquisition.

Acknowledgments

We wish to convey our appreciation to Clemson University for supporting this work. Open access funding enabled and organized by Carolinas 2023.

Supplementary Materials

The supplementary materials contain additional information on the hydrodynamic diameter of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements for each pH (1–14); the particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements for each pH (1–14); the zeta potential of CeO₂ NPs under varying concentrations of NaCl/CaCl₂ in the absence/presence of HA/FA at three different pHs (8.2, 10.2, and 12.2) and different time intervals of around 1 hr measurements; the hydrodynamic diameter of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements at three different pHs (8.2, 10.2, and 12.2) and each NaCl/CaCl₂ concentration; the particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements in the presence of NaCl at three different pHs (8.2, 10.2, and 12.2); the particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements in the presence of CaCl₂ at three different pHs (8.2, 10.2, and 12.2); the hydrodynamic diameter of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements at three different pHs (8.2, 10.2, and 12.2) and each NaCl/CaCl₂ concentration in the presence of HA; the particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements in the presence of HA, NaCl, and CaCl₂ at three different pHs (8.2, 10.2, and 12.2); the hydrodynamic diameter of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements at three different pHs (8.2, 10.2, and 12.2) and each NaCl/CaCl₂ concentration in the presence of FA; and the particle size distribution of CeO₂ NPs during the first 100 s and at ∼1 hr of around 1 hr measurements in the presence of FA, NaCl, and CaCl₂ at three different pHs (8.2, 10.2, and 12.2). (Supplementary Materials)