Open AccessArticle

Contrasting Perfluorooctanoic Acid Removal by Calcite Before and After Heat Treatment

Zhaohui Li

^1,*

Yating Yang

²,

Yaqi Wen

Yuhan Li

²,

Jeremy Moczulewski

³,

Po-Hsiang Chang

^2,*

Stacie E. Albert

¹ and

Lori Allen

Department of Geosciences, University of Wisconsin–Parkside, 900 Wood Road, Kenosha, WI 53144, USA

College of Resources and Environment, Fujian Agriculture and Forestry University, Fuzhou 350002, China

Department of Chemistry, University of Wisconsin–Parkside, 900 Wood Road, Kenosha, WI 53144, USA

Authors to whom correspondence should be addressed.

Environments 2025, 12(1), 29; https://doi.org/10.3390/environments12010029

Submission received: 16 November 2024 / Revised: 2 January 2025 / Accepted: 16 January 2025 / Published: 17 January 2025

(This article belongs to the Special Issue The Occurrence, Behavior, Fate, Risk Assessment, and Treatment of per- and Polyfluoroalkyl Substances)

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Figure 1
Sorption of PFOA on Cal and HCal. The solid lines are the Langmuir model fitting of the observed data. The right y-axis with solid symbols is the percentage of PFOA sorbed. "> Figure 2
Sorption kinetics of PFOA on Cal and HCal without pH adjustment under an initial concentration of 100 mg L−1. The solid lines are pseudo-second-order model fitting of the observed data. "> Figure 3
PFOA removal by Cal as affected by equilibrium solution pH under an initial PFOA concentration of 100 mg L−1. "> Figure 4
PFOA removal by Cal and HCal as affected by equilibrium solution ionic strength under an initial PFOA concentration of 100 mg L−1. "> Figure 5
PFOA removal by Cal and HCal as affected by equilibrium solution temperature under an initial PFOA concentration of 100 mg L−1. "> Figure 6
FTIR spectra of Cal and HCal after in contact with different initial PFOA concentrations (H represents HCal), respectively. The numbers are the initial PFOA concentrations in mg L−1. "> Figure 7
XRD patterns of Cal (a) and HCal (b) after equilibrated with different initial concentrations of PFOA (numbers in mg L−1), and the standard samples of Ca(OH)2 and CaO. "> Figure 8
TG (a) and DTG (b) analyses of Cal and HCal. The number is the initial PFOA concentration in mg L−1. "> Figure 9
SEM images of Cal (a) and HCal (c) and their SEM images after their sorption of PFOA from an initial concentration of 200 mg L−1, respectively (b,d). The EDS spectra of face and point scans of samples after in contact with PFOA solution for 24 h (e). ">

Versions Notes

Abstract

Calcites before and after calcination at 1000 °C were evaluated for their potential removal of perfluorooctanoic acid (PFOA) from water. After heat treatment, the PFOA sorption capacity increased by 25%, from 3.2 to 3.9 mg g⁻¹, and the affinity increased by 2.7 times, from 0.03 to 0.08 L mg⁻¹. Kinetically, the initial rate, rate constant, and equilibrium sorption were 8.7 mg g⁻¹ h⁻¹, 2.6 g mg⁻¹ h⁻¹, and 1.8 mg g⁻¹ for heat treated calcite, in comparison to 6.4 mg g⁻¹ h⁻¹, 3.1 g mg⁻¹ h⁻¹, and 1.4 mg g⁻¹ for calcite without heat treatment. X-ray diffraction analyses showed phase changing from calcite to calcium oxide after calcination. However, after contact with PFOA solutions for 24 h, the major phase changed back to calcite with a minute amount of Ca(OH)₂. These results suggest that using raw cement materials derived from heat treatment of limestone may be a good option for the removal of PFOA from water. Thus, further studies are needed to confirm this claim.

Keywords:

perfluorooctanoic acid; heat treatment; calcite; calcium oxide; sorption

1. Introduction

Polyfluorinated substances, also known as per- and polyfluoroalkyl substances (PFAS), are a group of synthetic chemicals used extensively, beginning in the 1940s [1]. Since the 1950s, they were also widely used in consumer products [2]. Due to their potential toxicities to animals and human beings, studies for their removal using novel materials have attracted great attention in recent years [3]. Perfluorooctanoic acid (PFOA) is one of the PFAS studied considerably due to its extensive worldwide distribution, persistence in the environment, possible toxic effects, and potential bioaccumulation [4]. Removal of PFOA can be achieved via degradation or adsorption.

For degradative removal, UV irradiation was often applied. Using UV radiation at a wavelength of 254 nm and in the presence of 10 µM Fe³⁺, PFOA at a concentration of 48 µM could be photochemically decomposed by 47% within 4 h, resulting in a defluorination ratio of 15.4% [5]. In another study, photo-reductive defluorination of PFOA at an initial concentration of 25 µM and in the presence of 0.8 mM KI resulted in fluoride release up to 98% in 14 h [6]. In comparison, photochemical decomposition of 10 µM PFOA reached 87% in 2 h in the presence of 0.5 mM IO₄⁻ under the irradiation of 254 nm [7]. More recently, significant photodegradation of PFOA was observed in six different types of soils collected in China and their intrinsic properties, which included light-shielding effect, organic carbon content, and silica content, all affected PFOA photodegradation [4]. For selected inorganic catalysts, PFOA photodegradation was influenced in the following order: Ga₂O₃ > TiO₂ > CeO₂ > In₂O₃ > CdS due to a progressive decrease in band gap energy [8].

In addition to photochemical degradation, others like electrochemical and thermal technologies are also important methods for PFAS removal [9]. PFOA could be electrochemically decomposed using a boron-doped diamond electrode in the presence of sulfate electrolytes and the highest PFOA decomposition could be achieved using a high current density, a high speed of mixing, and an acidic condition [10]. Thermal decomposition under typical combustion conditions of time, temperature, and excess air level in a municipal incinerator showed no detectable amount of PFOA under typical municipal incineration conditions for 2 s [11]. Chemically, a new heterogeneous activation material produced by hydrogen peroxide and persulfate coupled with iron-modified diatomite showed great promise for the decomposition of PFOA [12]. In comparison to chemical processes, the degradation efficiency of biological processes is lower [13].

For adsorptive removal, selecting effective sorbents is the key for recent studies. The materials tested include natural and synthetic materials. In one study, it was shown that porous sorbents with amine groups may result in high PFOA removal due to formation of micelles/hemi-micelles [3]. Sorption of PFOA on alumina resulted in a capacity of 0.002 mg g⁻¹ at a pH value of 4.3 [14]. In another study, the distribution coefficient K_d values for PFOA sorption on montmorillonite and kaolinite coated with organic matter were only 0.2–2.0 L kg⁻¹ [15]. Sorption of PFOA on heat treated hydrotalcite also showed great promise [16].

Mechanisms of PFOA sorption could include electrostatic interaction, hydrophobic interaction, ligand exchange, and hydrogen bonding. Hydrophobic interactions played an important role where the K_d of PFOA sorption onto wetland soils increased with the fraction of organic carbon (f_oc) and ionic strengths of inorganic salts, but decreased as the concentration of humic acid increased [17]. In addition, electrostatic interaction was speculated based on the K_d values of PFOA sorption onto clays and the f_oc values of clays [15].

In addition, determination of perfluorinated alkylated substances in Lombardia region in Italy, including their profile levels and assessment via monitoring activities, was conducted for a whole year in 2018 [18]. Bioaccumulation, biodistribution, toxicology, and biomonitoring of organofluorine compounds in aquatic organisms were summarized recently [19]. Particularly, the presence and biodistribution of PFOA in Paracentrotus lividus highlighted its potential application for environmental biomonitoring [20]. And, the controlled uptake of PFOA in adult specimens of Paracentrotus lividus and gene expression in their gonads and embryos were also evaluated [21]. These results suggested it was critical for the removal of PFAS from the environment, including water.

Although extensive studies were conducted for PFOA removal from water, more researches in exploring Earth materials for this environmental application are still in great need. Using CaCO₃ (limestone) to make calcium oxides is one of the oldest known chemical processes in the world [22]. On a larger scale, calcite was studied as a substrate for permeable reactive barriers (PRBs) to remove fluoride from contaminated groundwater in one study [23]. In another study, granulated lime (Ca(OH)₂) and calcium carbonate (CaCO₃) used as coagulants were evaluated for the removal of heavy metals, with more than 98% of As and Ni removal by the former and 97% of Ni and Zn removal by the latter from artificially contaminated water [24].

Water treatment residuals (Al-WTR) generated during drinking water treatment using alum salts were evaluated for the removal of perfluorooctanesulfonic acid (PFOS) and PFOA, with a PFOA sorption capacity of <0.3 mg g⁻¹ [25]. As for the removal of PFOS using calcite, rarely any reports could be found. In one study it was noticed that PFOS can be efficiently removed by a membrane in the presence of calcium cations, due to the bridging between Ca²⁺ and the sulfonate functional group of PFOS, which resulted in a large molecular size [26]. During a thermal treatment, PFOS gasified when the temperature was increased to around 425 °C, and the temperature decreased to 350 °C in the presence of Ca(OH)₂ [27]. In another study, it was found that different calcium compounds showed different mineralization effects of PFASs during a thermal process, owing to the different reaction mechanisms and calcium hydroxide was the most effective Ca reagent for PFAS defluorination due to the possible conversion of carbon–fluorine bonds in PFAS to carbon–hydrogen bonds via the hydrodefluorination reaction [28].

In comparison to calcite, the critical but overlooked influence of dolomite in perfluoroalkyl acids (PFAA) adsorption may affect the PFAA transport within a groundwater system [29]. Aquifer materials of higher dolomite content showed significantly stronger affinity for PFAS than the aquifer materials of low dolomite content [30]. Still more research is needed to study the removal of contaminants by calcite and dolomite after heat treatment. Initial studies showed great removal of cationic and anionic color dyes using heat treated calcite and dolomite [31,32]. In this study the potential use of calcite before and after heat treatment was evaluated. Although the results are preliminary, they could provide further insights into the use of lime materials produced from limestone for the removal of PFAS from water. This will extend the application of either lime materials or cement for contaminant immobilization in addition to their major function as construction materials.

2. Materials and Methods

2.1. Materials

The PFOA used was pentadecafluorooctanoic acid in ammonium salt and was purchased from Sigma-Aldrich. It has a CAS # of 335-67-1, a formula of C₈HF₁₅O₂, a molecular mass of 414.07 g/mol, a solubility of 9500 mg L⁻¹ [33]. The calcite used was obtained from Guzhang Shan Lin Shi Yu Mineral Co., Ltd., Changzhou, Hunan province, China. It has a CAS # of 13397-26-7 and purity > 99.5%. The sample was pure after X-ray diffraction (XRD) analyses (See Section 3.4.2). For comparison, the raw calcite (Cal) was used as a control. It was ground to <200 meshes. For heat treatment, the ground Cal was heat to 1000 °C and maintained for 3 h. Then, the samples were cooled down in the oven over night. The treated samples are named HCal. The HCal was X-rayed to investigate its transformation and it was found that the Cal was converted to CaO (See Section 3.4.2) after heat treatment and before in contact with water.

2.2. Experimental Methods

For the isotherm study, the initial PFOA concentrations varied from 0 to 200 mg L⁻¹. For all other studies, the initial PFOA concentration was fixed at 100 mg L⁻¹. The solution volume used was 20 mL and the solid mass used was 1.00 g per sample. The mixing time was 24 h for the isotherm study and was 0.25, 0.5, 1.0, 2.0, 4.0, 8.0, 16.0, and 24.0 h for the kinetic study. For the pH-effect study, the equilibrium solution pH was adjusted to values between 5 and 11 by adding 2 M NaOH or 2 M HCl drop-wise. However, due to the strong basic nature of HCal, the pH study was difficult to conduct. The solution pH was not adjusted in the isotherm and kinetic studies. Final pH values were between 6–7 for Cal and about 12 for HCal in the isotherm study. The mixtures were centrifuged at 5000 rpm for 5 min, and the supernatants were passed through 0.22 μm filters before analyses. All experiments were run in duplicate.

2.3. Instrumental Analyses

The equilibrium PFOA concentrations were analyzed by Shimadzu liquid chromatograph LC-30A and liquid chromatograph mass spectrometer LCMS-8050 system. The other components include valve unit FCV-20AH₂, SIL-30AC autosampler, prominence column oven CTO-20AC, degassing unit DGU-20A_5R, and back pressure regulator SFC-30A. The column is Shim-pack GIST-HP C18-AQ (2.1 × 100 mm, 1.9 μm). The mobile phase consisted of 100% water with 2 mM ammonium acetate (mobile phase A) and 100% acetonitrile (mobile phase B). The gradient of analysis was started from 28% B, increased to 90% B for 3 min, and then maintained for 1 min. It was then increased again to 100% B for 0.1 min and maintained for 1 min. The gradient elution was then decreased back to 28% B for 0.1 min and maintained for 2.8 min. The flow rate was 0.3 mL min⁻¹, the column temperature was 40 °C, and the retention time was 4.4 min for PFOA. The standards were up to 100 µg L⁻¹ with an r² greater than 0.99. The detection limit of PFOA was 0.5 µg L⁻¹ in this study. Data were obtained using a mass spectrometer with the following parameters: DL temperature at 250 °C; heat block temperature at 400 °C; nebulizing gas flow at 3 L min⁻¹; drying gas flow at 10 L min⁻¹; heating gas flow at 10 L min⁻¹; interface voltage at 3.0 kV; and focus voltage at 2.12 kV. A suitable mass spectrometer transition and collision energy was applied for each compound. The calibration solution was prepared from the standard mixture solution, with both PFOA and internal standard-¹³CM₄PFOA concentrations ranging from 1 to 200 μg L⁻¹, and each internal standard at a concentration of 5 µg L⁻¹. The ratios of the peak area of the target compound over the peak area of its internal standard were plotted against the ratios of the concentrations for the target and its internal standard in the solution. By optimizing the parameters for precursor transmission and fragmentation, transmissions of 412.8 and 459 were selected for quantitation of PFOA, respectively. The whole samples including the standards were diluted by methanol (Optima@ LC/MS, LOT 225738) to be further analyzed.

Powder XRD patterns of Cal and HCal after PFOA sorption were recorded in the 2θ range of 15–65° using a Rigaku Ultima IV diffractometer (CuKα radiation) equipped with a D/teX Ultra detector, using a step size of 0.01° and a scan speed of 2° min⁻¹. The thermogravimetric (TG) analyses were performed on a Netzsch STA 449 F3 Jupiter (Selb, Germany) with a heating rate of 10 °C min⁻¹ under nitrogen conditions. Fourier transform infrared (FTIR) spectra were acquired from 400 to 4000 cm⁻¹ by accumulating 256 scans at a resolution of 4 cm⁻¹ on a Thermo Nicolet iS50 spectrometer using the KBr pressing method. The scanning electron microscopy (SEM) is Hitachi SU8010 and energy dispersive X-ray spectroscopy (EDS) analysis was conducted using X-Max^N Silicon Drift detector.

3. Results

3.1. Isotherm Study

With the initial PFOA concentrations used, PFOA removal by Cal varied from 54% to 80%. For HCal, PFOA removal varied from 76% to 94% (Figure 1). The data were fitted to several isotherm models, and the Langmuir model fitted the best with the r² greater than 0.99 (Figure 1). The Langmuir model has the form:

C_{S} = \frac{K_{L} {S_{m} C}_{L}}{1 + K_{L} C_{L}}

(1)

where C_L is the equilibrium PFOA concentration in solution (mg L⁻¹) and C_S and S_m are the amount of PFOA sorbed on solid surfaces at equilibrium and the PFOA sorption capacity (mg g⁻¹), and K_L is the Langmuir coefficient (L mg⁻¹), reflecting the affinity of PFOA on solid surfaces. Equation (1) can be rearranged to a linear form:

\frac{C_{L}}{C_{S}} = \frac{1}{K_{L} S_{m}} + \frac{C_{L}}{S_{m}}

(2)

so that the K_L and S_m values can be determined by a linear regression. The fitted Langmuir parameters were S_m = 3.9 mg g⁻¹ and K_L = 0.08 L mg⁻¹ for PFOA sorption on HCal, in comparison to S_m = 3.2 mg g⁻¹ and K_L = 0.03 L mg⁻¹ for PFOA sorption on Cal. Similarly, Freundlich and Langmuir models were used to fit the PFOA sorption on zeolite and on granular activated carbon (GAC) [34]. It can also be found in Figure 1 is that at low equilibrium concentrations (<20 mg L⁻¹) the sorption was more effective when HCal was used.

3.2. Kinetic Study

Relatively speaking, the sorption of PFOA was fast, with equilibrium reached in 5 h (Figure 2). The data fitted to the pseudo-second-order kinetic model best with r² > 0.9998. The pseudo-second-order kinetic model is expressed as:

q_{t} = \frac{k q_{e}^{2} t}{1 + {k q}_{e} t}

(3)

where k is the sorption rate constant (g mg⁻¹ h⁻¹),

k q_{e}^{2}

is the initial rate (mg g⁻¹ h⁻¹), and q_t and q_e (mg g⁻¹) are the amount of PFOA sorbed at time t and at equilibrium. Equation (3) can be re-arranged into a linear form:

\frac{t}{q_{t}} = \frac{1}{k q_{e}^{2}} + \frac{1}{q_{e}} t

(4)

When Equation (4) was used to fit the kinetic data, the initial rate, rate constant, and equilibrium sorption were 6.4 mg g⁻¹ h⁻¹, 3.1 g mg⁻¹ h⁻¹, and 1.4 mg g⁻¹ for PFOA sorption on Cal, respectively. In comparison, sorption of PFOA on HCal resulted in an initial rate of 8.7 mg g⁻¹ h⁻¹, a rate constant of 2.6 g mg⁻¹ h⁻¹ with an equilibrium sorption of 1.8 mg g⁻¹, respectively. Separately, uptake of PFOA on GAC and powdered activated carbon (PAC) was also fitted with the pseudo-second-order kinetic model, having equilibrium sorptions of 0.30 and 0.42 mmol g⁻¹, initial rates of 0.01 and 2.45 mmol g⁻¹ h⁻¹, and rate constants of 0.07 and 13.9 g mmol⁻¹ h⁻¹ at a pH value of 7, respectively [35]. PFOA sorption on Al-WTR also followed the pseudo-second-order kinetic model with equilibrium sorption at <0.3 mg g⁻¹ [25].

3.3. Influence of Equilibrium Solution pH, Ionic Strength, and Temperature

As the equilibrium, solution pH was very high (>12) when HCal was used; in this study, the influence of solution pH on PFOA sorption by HCal was not assessed. In contrast, for Cal, as the equilibrium solution pH increased from 5 to 11, the PFOA sorption decreased by 50%, from about 1.5 to 0.7 mg g⁻¹ (Figure 3). As the PFOA is in anionic form, higher solution pH will result in more negative charges on Cal surfaces, increasing repulsion between Cal surfaces and the monoanion form of PFOA⁻, reducing PFOA sorption. Similarly, sorption of PFOA on multi-walled carbon nanotubes decreased as solution pH increased [36]. Moreover, sorption of PFOA on alumina decreased by as much as 90% as pH increased from 4 to 7, as variable surface charge was associated with alumina [14]. Also, on the mineral goethite, PFOS sorption decreased as the solution pH increased [37]. In this study, the equilibrium solution pH value was about 12 for HCal. Thus, under high pH conditions, HCal showed much higher PFOA removal in comparison to Cal.

Influence of solution ionic strength on PFOA sorption was minimal for both Cal and HCal (Figure 4). Still, for HCal, PFOA sorption was about 1.8 mg g⁻¹ in comparison to 1.4 mg g⁻¹ for Cal. This indicated that the neutral ionic salt has minimal influence on the sorption of monoanion forms of PFOA⁻ on solid surfaces. On the contrary, PFOS sorption on goethite increased as the solution ionic strength increased from 0.001 to 0.1 M as NaCl [37].

For Cal, as the temperature increased from 30 to 60 °C, the PFOA sorption increased from 1.4 to 1.9 mg g⁻¹. In contrast, the PFOA sorption increased only slightly from 1.8 to 1.9 mg g⁻¹ for HCal. Again, it showed that HCal is slightly better than Cal for PFOA removal under elevated temperature.

The solute distribution coefficient K_d between the solid and the solution is related to the thermodynamic parameters by:

L n K_{d} = - \frac{Δ H}{R T} + \frac{Δ S}{R}

(5)

where the changes in enthalpy and entropy after PFOA sorption are represented by ΔH and ΔS, respectively; and the gas constant and the equilibrium temperature in K are represented by R and T. Their relation is demonstrated in Figure 5. The free energy of sorption ΔG is realted to ΔH and ΔS by:

Δ G = Δ H - T Δ S

(6)

The calculated thermodynamic values are listed in Table 1. Overall, the ΔG values were between –25 and –37 kJ mol⁻¹ for Cal and –32 and –36 kJ mol⁻¹ for HCal, suggesting that PFOA sorption on both minerals was endothermic. The ΔS° values were 0.40 and 0.15 kJ mol⁻¹ K⁻¹, for PFOA sorption by Cal and HCal, respectively, indicating more random orientation of sorbed PFOA molecules on the mineral surfaces. The ΔH° values were 96 and 13 kJ mol⁻¹ for PFOA sorption on Cal and HCal, indicating endothermic PFOA sorption on Cal and HCal (Table 1).

3.4. Instrumental Characterization of Minerals After PFOA Sorption

3.4.1. FTIR Analyses

The three peaks at 1149, 1201, and 1244 cm⁻¹ were attributed to vibrations of the CF₂ and CF₃ groups for raw PFOA [38]. In this study, they were located at 1150, 1210, and 1244 cm⁻¹ (Figure 6). However, these peaks did not show up on the Cal and HCal after PFOA sorption. This is because the amount of PFOA sorbed was less than 1% of the total mass of the Cal and HCal as determined in the isotherm study.

For Cal, the sharp bands at 714 and 877 cm⁻¹ and a broad band at 1440 cm⁻¹ are characteristics of calcite. The 877 cm⁻¹ was attributed to the CO₃ bending vibration band while the 1440 cm⁻¹ was attributed to the CO₃ stretching vibration band and they are typical fundamental bands for calcite [39]. For HCal, there is one more band at 3652 cm⁻¹. It was assigned to the band of stretching vibration ν₁ of Ca(OH)₂ [40]. The presence of Ca(OH)₂ was also confirmed in the XRD study (See Section 3.4.2). The Ca(OH)₂ should be produced from CaO after reaction with water and then consumed to the conversion to calcite.

3.4.2. XRD Analyses

For Cal, the XRD patterns matched well with the JCPDS card No. 47-1743 [40]. The original sample is very pure without peaks of other minerals. After PFOA sorption, no changes in mineral phase were observed (Figure 7a). For HCal without in contact with water, the XRD pattern matched well with that of CaO, suggesting the decomposition of calcite after heating and formation of calcium oxides (Figure 7b). Also shown in Figure 7b is the XRD pattern of Ca(OH)₂. For the HCal after contact with water or PFOA solution, the major phase was calcite with minute amounts of Ca(OH)₂. The XRD results suggested that conversion of CaO to mainly calcite via Ca(OH)₂ as an intermediate product after 24 h of equilibration with water may contribute to the elevated PFOA removal by HCal. In a previous study, based on TG and XRD analyses, it was found that CaO was completely converted to Ca(OH)₂ and then to CaCO₃ in ambient conditions within several days [41].

3.4.3. Thermogravimetric Analyses

The TG and DTG curves of PFOA (Figure 8) matched well with published data [42]. The peak decomposition of PFOA occurred at 175 °C (Figure 8b), at which 95% of the mass was lost. For raw Cal and Cal after sorbing PFOA from 200 mg L⁻¹ solution, the decomposition temperature was at 783 and 772 °C, agreeing well with the calcite that displayed maximum weight loss at 750 to 880 °C, and was attributed to the CO₂ removal and formation of CaO [43]. For raw HCal, decomposition temperature was at 398 °C. This could be due to the decomposition of Ca(OH)₂ between 350 and 400 °C [41]. For the HCal sample after PFOA sorption from 200 mg L⁻¹ solution, the DTG curves showed two peaks, one at 418 °C and the other at 736 °C with the former due to decomposition of Ca(OH)₂, and the latter due to decomposition of CaCO₃ (Figure 8). Calcium oxide is a reactive material in ambient conditions, and it will carbonate and hydrate in a timeframe of hours [41]. A TG analysis of CaO stored in ambient conditions for several days showed 39% mass loss in two steps with the first due to decomposition of Ca(OH)₂ between 350 and 400 °C (20% of the total weight loss) and the second, due to decomposition of CaCO₃ above 550 °C (80% of the total weight loss) [41], implying almost complete conversion from CaO to Ca(OH)₂ and then to CaCO₃ in ambient conditions within several days. Thus, the TG and DTG curves together with the XRD results confirm that CaO and/or Ca(OH)₂ played a critical role for PFOA sorption on HCal.

3.4.4. SEM Observation and EDS Elemental Analyses

The SEM images of Cal before and after contact with the 200 mg L⁻¹ PFOA solution showed similar morphology, with particle size ranging from 0.5 to 5 μm and well-developed cleavages along {10–14} directions (Figure 9a,b). In contrast, the particle size was more uniform, at about 0.5 μm for HCal before and after contact with the PFOA solution. Meanwhile, the {10–14} cleavage is not visible in the HCal samples (Figure 9c,d), indicating a change from Cal to CaO after heating. Even after equilibration with PFOA solution, the {10–14} cleavage is rarely seen while the XRD pattern showed calcite phase (Figure 7b). This may indicate the particles under the SEM were newly formed calcite without well-developed crystal morphology or cleavage. The EDS spectra of face and point scans of the samples showed the presence of C in the samples, again agreeing well with the XRD analyses, indicating phase changes from calcium oxide into calcite after contact with the PFOA solution. A similar SEM morphological observation was found for CaO after being exposed to ambient conditions for several days [41].

3.5. Possible PFOA Removal Mechanism by Cal and HCal

The critical micelle concentration (CMC) of PFOA was 25 mM [44]. It was also reported as 3460 mg L⁻¹ [45]. Thus, the PFOA molecules were in monomer form under the studied concentrations and the removal involved in interactions between PFOA monomers and the Cal or HCal surfaces. Calcium hydroxide was considered to be the most effective Ca reagent for PFAS defluorination as the carbon−fluorine bonds in PFASs could be converted to carbon−hydrogen bonds via the hydrodefluorination reaction [29]. As summarized by Cagnetta et al. [46], calcium oxide was effective for the degradation of PFOS and PFOA in one study, but poor with low fluoride recovery in another study [47] and this difference was attributed to the milling conditions adopted for the treatment [46]. In this study, it was found that HCal performed better than Cal, indicating that CaO and/or Ca(OH)₂ played an important role in PFAS removal. It is not so clear whether it is due to sorptive or degrative removal. Thus, further work is needed to determine whether the sorbed PFOA could be desorbed or not under different conditions. In addition, an HPLC-MS analysis should be deployed to determine whether any degradation products could be detected or not. If so, the speciation and quantity should be determined.

3.6. Implication of HCal for PFOA Removal

After calcination, the PFOA sorption capacity increased from 3.2 to 3.9 mg g⁻¹, an increase of only about 25%. In comparison, the affinity of PFOA for the sorbent increased from 0.03 to 0.08 L mg⁻¹ after calcination. These values are significantly smaller than the PFOA sorption on calcined hydrotalcite (CHT), which had a value of 1587 mg g⁻¹ [16]. However, in China, the median and maximum PFOS and PFOA concentrations in environmental water were 0.4 and 2.4 and 0.1 and 1.3 ng L⁻¹ for remote areas, but increased to 4.0 and 14.1 and 3.9 and 30.8 ng L⁻¹ for urban areas, respectively [48], as over the past 30 years, China had rapid increase in PFAS production and usage [49]. In this study the PFOA removal was more than 90% with initial concenttrations up to 100 mg L⁻¹. Thus, at such low concentrations, the HCal could definitely achieve a better job for the removal of PFOA from natural water.

Kinetic-wise, the initial rate, rate constant, and equilibrium sorption were 6.4 mg g⁻¹ h⁻¹, 3.1 g mg⁻¹ h⁻¹, and 1.4 mg g⁻¹ for PFOA sorption on Cal, respectively. And these values changed to 8.7 mg g⁻¹ h⁻¹, 2.6 g mg⁻¹ h⁻¹, and 1.8 mg g⁻¹, still, only about a 30% increase. Thus, one question is whether it is worth calcining calcite up to 1000 °C for the removal of PFOA. The answer is absolutely, yes. For cement making, raw materials made of mainly limestone, clay, sand, and iron are mixed together and ground to form a homogeneous blend which is then processed in a rotary kiln by calcining the raw materials to decompose CaCO₃ (at about 900 °C), through which CO₂ is released and CaO is produced [50]. CaO is one of the most promising sorbents for CO₂ capture [51]. Addition of CaO before smouldering in treatment of PFAS in sewage sludge could reduce PFAS content in emissions by 97–99% [52]. As such, the results from this study may indicate that raw cement materials can be used to remove PFOA from water, a new discovery for the use of cement materials, and a further extension of the use of CaO. For this reason, raw materials for cement should be tested in the future for comparison.

4. Conclusions

The study showed that calcite was able to remove PFOA from water at the capacity of 3.2 mg g⁻¹. After the calcite was heated to 1000 °C for 3 h, at which calcite was converted to calcium oxide, the PFOA removal increase by 30% to 3.9 mg g⁻¹. However, the PFOA affinity for the sorbent increased from 0.03 to 0.08 L mg⁻¹, by 2.7 times. The initial rate for PFOA removal also increased 30% from 6.4 mg g⁻¹ h⁻¹ to 8.7 mg g⁻¹ h⁻¹, after heat treatment. The finding showed the following promises: (1) calcite could form a natural attenuation filtration media to reduce the PFOA concentration, if the groundwater contains elevated PFOA level; (2) as a component for concrete production, the raw materials produced from heat treatment of limestone, it also showed improved PFOA removal. Regardless the mechanism of PFOA removal by Cal or HCal, this study showed initial optimistic results for PFOA removal by Earth materials and heat-treated Earth materials. Further detailed studies on transport of PFOA through the limestone or dolostone aquifer materials would warrant a best assessment of PFAS in the environment. Moreover, transport of PFAS thought HCal could be further explored to evaluate whether it could serve as packing media for water infiltration in household to larger scales.

Author Contributions

Conceptualization, Z.L. and P.-H.C.; methodology, Z.L. and P.-H.C.; validation, Z.L., P.-H.C. and L.A.; formal analysis, Y.Y., Y.W., Y.L. and J.M.; investigation, Y.Y., Y.W., Y.L. and J.M.; resources, Z.L. and P.-H.C.; data curation, Y.Y., Y.W. and Y.L.; writing—original draft preparation, Z.L.; writing—review and editing, S.E.A. and Z.L.; supervision, Z.L., P.-H.C. and L.A.; project administration, Z.L.; funding acquisition, Z.L. All authors have read and agreed to the published version of the manuscript.

Funding

This work was supported by the following grants: (1) Scientist in Residency grant from WiSys; and (2) a grant from Fresh Water Collaborative of Wisconsin to Engaging Undergraduate Students in Research on using Earth Materials for contaminant removal including dyes and Per- and Polyfluorinated substances.

Data Availability Statement

Data are available upon request.

Conflicts of Interest

The authors declare no conflicts of interest.

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Figure 1. Sorption of PFOA on Cal and HCal. The solid lines are the Langmuir model fitting of the observed data. The right y-axis with solid symbols is the percentage of PFOA sorbed.

Figure 2. Sorption kinetics of PFOA on Cal and HCal without pH adjustment under an initial concentration of 100 mg L⁻¹. The solid lines are pseudo-second-order model fitting of the observed data.

Figure 3. PFOA removal by Cal as affected by equilibrium solution pH under an initial PFOA concentration of 100 mg L⁻¹.

Figure 4. PFOA removal by Cal and HCal as affected by equilibrium solution ionic strength under an initial PFOA concentration of 100 mg L⁻¹.

Figure 5. PFOA removal by Cal and HCal as affected by equilibrium solution temperature under an initial PFOA concentration of 100 mg L⁻¹.

Figure 6. FTIR spectra of Cal and HCal after in contact with different initial PFOA concentrations (H represents HCal), respectively. The numbers are the initial PFOA concentrations in mg L⁻¹.

Figure 7. XRD patterns of Cal (a) and HCal (b) after equilibrated with different initial concentrations of PFOA (numbers in mg L⁻¹), and the standard samples of Ca(OH)₂ and CaO.

Figure 8. TG (a) and DTG (b) analyses of Cal and HCal. The number is the initial PFOA concentration in mg L⁻¹.

Figure 9. SEM images of Cal (a) and HCal (c) and their SEM images after their sorption of PFOA from an initial concentration of 200 mg L⁻¹, respectively (b,d). The EDS spectra of face and point scans of samples after in contact with PFOA solution for 24 h (e).

Table 1. Thermodynamics of PFOA sorption on Cal and HCal.

Minerals	ΔG° (kJ mol⁻¹)			ΔH° (kJ mol⁻¹)	ΔS° (kJ mol⁻¹ K⁻¹)
	303 K	318 K	333 K
Cal	−24.8	−30.8	−36.7	96.2	0.40
HCal	−32.0	−34.2	−36.5	13.4	0.15

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Li, Z.; Yang, Y.; Wen, Y.; Li, Y.; Moczulewski, J.; Chang, P.-H.; Albert, S.E.; Allen, L. Contrasting Perfluorooctanoic Acid Removal by Calcite Before and After Heat Treatment. Environments 2025, 12, 29. https://doi.org/10.3390/environments12010029

AMA Style

Li Z, Yang Y, Wen Y, Li Y, Moczulewski J, Chang P-H, Albert SE, Allen L. Contrasting Perfluorooctanoic Acid Removal by Calcite Before and After Heat Treatment. Environments. 2025; 12(1):29. https://doi.org/10.3390/environments12010029

Chicago/Turabian Style

Li, Zhaohui, Yating Yang, Yaqi Wen, Yuhan Li, Jeremy Moczulewski, Po-Hsiang Chang, Stacie E. Albert, and Lori Allen. 2025. "Contrasting Perfluorooctanoic Acid Removal by Calcite Before and After Heat Treatment" Environments 12, no. 1: 29. https://doi.org/10.3390/environments12010029

APA Style

Li, Z., Yang, Y., Wen, Y., Li, Y., Moczulewski, J., Chang, P.-H., Albert, S. E., & Allen, L. (2025). Contrasting Perfluorooctanoic Acid Removal by Calcite Before and After Heat Treatment. Environments, 12(1), 29. https://doi.org/10.3390/environments12010029

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