**Abstract**

Although previous studies have shown that sulfate can either increase cation leaching or enhance cation adsorption in soil, little is known about the factors behind these phenomena. To learn more about them, calcium adsorption experiments were carried out with kaolinite and gibbsite at initial pH values 4 and 6 and in the presence of 1 or 20 mmolc L<sup>−</sup><sup>1</sup> of either nitrate or sulfate. The results indicated that limited sulfate-calcium coadsorption occurred on gibbsite when it was in contact with the dilute solution of CaSO4.2H2O at pH ~ 7. Regarding mineral and pH values, calcium adsorption from the concentrated solutions decreased with sulfate possibly because of the presence of ~31% of the CaSO4 0 ion pair in the concentrated CaSO4.2H2O solutions and the low free calcium activity therein. Calcium adsorption on kaolinite and gibbsite from all concentrated solutions was reduced when the initial pH changed from 4 to 6 suggesting a negative salt effect on that process. In addition to indicating negligible participation of gibbsite in calcium adsorption, our findings also suggest that higher amounts of gypsum applied to lime-amended oxisols reduce the effectiveness of the main oxisol clay-sized mineral capable of adsorbing cations, i.e., kaolinite, to impair calcium leaching. The uptake data were complemented with some zeta-potential measurements, which supported the lack of substantial uptake of calcium even in the presence of sulfate. Some modeling calculations using the only available model covering sulfate and calcium on gibbsite have been done to rationalize the experimental data, but the model is only able to involve pure electrostatic attraction of calcium, which is not sufficient to produce

substantial uptake. Finally, the aluminol basal plane that is present on both gibbsite and kaolinite has been additionally studied using second harmonic generation (SHG) down to 4°C, because the ion-pair formation decreases with decreasing temperature. The second harmonic results confirm the patterns observed in the electrokinetic measurements with kaolinite being quite comparable to the sapphire basal plane. Also and quite clearly, the presence of CaSO4 solutions caused temperature dependence different from pure CaCl2 and Na2SO4 solutions. The latter were essentially behaving like pure water. The difference between the calcium chloride and sulfate systems can be explained by sulfate interaction and might be linked to the temperature dependence of the formation of the CaSO4 ion pair. The temperature dependency study could be an important starting point for looking at ice nucleation in the presence of the three different solutions and more strongly link aqueous chemistry to ice nucleation processes.

**Keywords:** acidic soils, double-layer screening, gypsum, leaching, oxisols, ice nucleation

## **1. Introduction**

Gypsum (~95% m/m CaSO4.2H2O) can be used as soil amendment to lessen the phytotoxicity of available aluminum found in acidic subsoil depths that, although accessible to the plant roots, are not affected by surface liming [1]. Field and laboratory experiments have indicated that gypsum application on soil can increase the leaching of plant nutrients such as potassium, calcium, and magnesium [2–5]. Such an effect could be due to the formation of negative or neutral ion pairs (e.g., KSO4 −, CaSO4 0 , and MgSO4 0 ) in the soil solution as predicted by the hard and soft (Lewis) acid and base theory [6]. In contrast, experiments conducted with soil samples rich in gibbsite, hematite, goethite, or allophane have suggested that sulfate adsorption enhances calcium adsorption [7–13]. Fahrenhorst et al. [4] found that the gypsum potential of increasing the cation exchange capacity (CEC) of an oxisol from Amazon corresponded on a 1:1 basis with the sulfate adsorption by that soil. Okuma and Alves [14] observed, in the absence of phosphate, that the enhancing effect of sulfate on CEC increased with the soil contents of iron and aluminum oxides. Although Pearce and Sumner [11] have suggested that the equivalent retention of Ca2+ and SO4 <sup>2</sup><sup>−</sup> observed in kaolinitic subsoil could be due the combination of several mechanisms, including precipitation, specific adsorption, ionic strength-induced charging, and ion-pair adsorption, to date such phenomenon remains to be clarified.

Among the surface chemical equations used to model solution pH and ionic strength effects on ion complexation on a standard gibbsite surface (≡AlOH0 ), those referring to sulfate complexation are given by [15]

$$\equiv \text{AlOH}^{0} + \text{SO}\_{4}^{2-} + \text{H}^{\*} \leftrightarrow \equiv \text{AlSO}\_{4}{\cdot} + \text{H}\_{2}\text{O} \tag{1}$$

$$\equiv \text{AlOH}^{0} \star \text{SO}\_{4}^{2-} \leftrightarrow \equiv \text{AlOHSO}\_{4}^{2-} \tag{2}$$

**43**

suspensions (1 g L<sup>−</sup><sup>1</sup>

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration…*

lead, and cadmium with phosphate and ferrihydrite [17, 18]. Calcium complexation

≡ AlOH<sup>0</sup> + Ca2+ ↔ ≡ AlOCa<sup>+</sup> + H<sup>+</sup> (3)

Although the surface chemical Eqs. (1)–(3) refer to adsorption sites located on the gibbsite edge faces, the basal planes of plate-shaped gibbsite and kaolinite as well seem to be operational in calcium adsorption as demonstrated by surface force measurements [19–21]. Considering that hydrolyzed species are preferably adsorbed to oxide surfaces [22], cations that hydrolyze at high pH values, such as calcium, would be expected not to be adsorbed on gibbsite and kaolinite edge sites at low pH. However, the above results obtained with atomic force microscopes indicate that cations presenting high hydrolysis constants can be retained on basal

Besides mineral surface properties, solution features must also be taken into account when evaluating ion adsorption. The formation of aqueous neutral ionic pairs can reduce the ion adsorption. The saturated solution of calcium sulfate (~15.8 mM), for instance, contains about 35% of total dissolved calcium and sulfate

formation, an increase in the ionic strength can, in some cases, decrease ion adsorp-

In this chapter, we studied the effects of the aqueous sulfate, salt concentration, and solution pH on the amounts of calcium adsorbed on kaolinite and gibbsite. These minerals were chosen because their joint contents in the oxisol clay fraction usually surpass 50% (m/m). Furthermore, while kaolinite is the most abundant cation-adsorbing mineral found in less weathered oxisols, gibbsite is the predominant clay-sized mineral in their deeply weathered counterparts [25]. The purpose of the present investigation was to evaluate the calcium adsorption behavior of kaolinite and gibbsite in oxisols amended with rich-sulfate sources such as gypsum.

A kaolinite sample from Minas Gerais, Brazil, and a laboratory-prepared sample

The mineralogical purity of the samples was checked through powder X-ray diffraction (XRD). The XRD patterns were measured using a Philips PW1830 diffractometer

) prepared in 0.01 M KNO3 solutions with different pH values,

of gibbsite were used in the experiments. The gibbsite sample was prepared via titration of 500 mL of 1 M Al3+ (added as AlCl3.6H2O) with 315 mL of 4 M NaOH followed by precipitate dialysis for 1 month against daily refreshed deionized water kept at 50°C [26]. Both samples were washed with deionized water, oven-dried (75°C/48 h), ground in an agate mortar, and passed through a 0.3-mm sieve.

(PANalytical, Almelo, the Netherlands) in continuous scan mode (0.02°2θ s<sup>−</sup><sup>1</sup>

to 90°2θ CuKα. The particle shapes were examined via scanning electron microscopy (SEM) using a LEO 435 VP microscope (LEO Electron Microscopy Inc., Thornwood, NY). The pH at the isoelectric point (IEP) was determined by graphical interpolation from zeta-potential (ζ) measurements. The ζ values were measured in triplicate with a NanoBrook Omni analyzer (Brookhaven Instruments, Holtsville, NY) in mineral

ion-pair. Thus, both ions can be less adsorbed if such retention does

ion-pair formation toward the formation of

<sup>2</sup><sup>−</sup>. Furthermore, even with limited aqueous ion-pair

) from 3

on gibbsite has in turn been modeled with the following equation [15]:

planes of gibbsite and kaolinite even under acidic condition.

0

*DOI: http://dx.doi.org/10.5772/intechopen.81273*

as the CaSO4

0

not shift the equilibrium of CaSO4

free aqueous ions Ca2+ and SO4

tion on the solid phase [23, 24].

**2. Material and methods**

**2.1 Minerals**

The above equations indicate that sulfate adsorption to gibbsite gives rise to negatively charged surface species, decreases the net positive surface charge, and increases the solution pH. Such features could favor sulfate-cation co-adsorption even at lower pH because sulfate adsorption on gibbsite increases as pH decreases [16]. Indeed, the acidic condition is required for the formation of ternary surface complexes of copper,

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration… DOI: http://dx.doi.org/10.5772/intechopen.81273*

lead, and cadmium with phosphate and ferrihydrite [17, 18]. Calcium complexation on gibbsite has in turn been modeled with the following equation [15]:

$$\equiv \text{AlOH}^{0} \star \text{Ca}^{2+} \leftrightarrow \equiv \text{AlOCa}^{\ast} \star \text{H}^{\ast} \tag{3}$$

Although the surface chemical Eqs. (1)–(3) refer to adsorption sites located on the gibbsite edge faces, the basal planes of plate-shaped gibbsite and kaolinite as well seem to be operational in calcium adsorption as demonstrated by surface force measurements [19–21]. Considering that hydrolyzed species are preferably adsorbed to oxide surfaces [22], cations that hydrolyze at high pH values, such as calcium, would be expected not to be adsorbed on gibbsite and kaolinite edge sites at low pH. However, the above results obtained with atomic force microscopes indicate that cations presenting high hydrolysis constants can be retained on basal planes of gibbsite and kaolinite even under acidic condition.

Besides mineral surface properties, solution features must also be taken into account when evaluating ion adsorption. The formation of aqueous neutral ionic pairs can reduce the ion adsorption. The saturated solution of calcium sulfate (~15.8 mM), for instance, contains about 35% of total dissolved calcium and sulfate as the CaSO4 0 ion-pair. Thus, both ions can be less adsorbed if such retention does not shift the equilibrium of CaSO4 0 ion-pair formation toward the formation of free aqueous ions Ca2+ and SO4 <sup>2</sup><sup>−</sup>. Furthermore, even with limited aqueous ion-pair formation, an increase in the ionic strength can, in some cases, decrease ion adsorption on the solid phase [23, 24].

In this chapter, we studied the effects of the aqueous sulfate, salt concentration, and solution pH on the amounts of calcium adsorbed on kaolinite and gibbsite. These minerals were chosen because their joint contents in the oxisol clay fraction usually surpass 50% (m/m). Furthermore, while kaolinite is the most abundant cation-adsorbing mineral found in less weathered oxisols, gibbsite is the predominant clay-sized mineral in their deeply weathered counterparts [25]. The purpose of the present investigation was to evaluate the calcium adsorption behavior of kaolinite and gibbsite in oxisols amended with rich-sulfate sources such as gypsum.

## **2. Material and methods**

#### **2.1 Minerals**

*Advanced Sorption Process Applications*

aqueous chemistry to ice nucleation processes.

ice nucleation

CaSO4 0

Ca2+ and SO4

**1. Introduction**

, and MgSO4

0

substantial uptake. Finally, the aluminol basal plane that is present on both gibbsite and kaolinite has been additionally studied using second harmonic generation (SHG) down to 4°C, because the ion-pair formation decreases with decreasing temperature. The second harmonic results confirm the patterns observed in the electrokinetic measurements with kaolinite being quite comparable to the sapphire basal plane. Also and quite clearly, the presence of CaSO4 solutions caused temperature dependence different from pure CaCl2 and Na2SO4 solutions. The latter were essentially behaving like pure water. The difference between the calcium chloride and sulfate systems can be explained by sulfate interaction and might be linked to the temperature dependence of the formation of the CaSO4 ion pair. The temperature dependency study could be an important starting point for looking at ice nucleation in the presence of the three different solutions and more strongly link

**Keywords:** acidic soils, double-layer screening, gypsum, leaching, oxisols,

Gypsum (~95% m/m CaSO4.2H2O) can be used as soil amendment to lessen the phytotoxicity of available aluminum found in acidic subsoil depths that, although accessible to the plant roots, are not affected by surface liming [1]. Field and laboratory experiments have indicated that gypsum application on soil can increase the leaching of plant nutrients such as potassium, calcium, and magnesium [2–5]. Such an effect could be due to the formation of negative or neutral ion pairs (e.g., KSO4

acid and base theory [6]. In contrast, experiments conducted with soil samples rich in gibbsite, hematite, goethite, or allophane have suggested that sulfate adsorption enhances calcium adsorption [7–13]. Fahrenhorst et al. [4] found that the gypsum potential of increasing the cation exchange capacity (CEC) of an oxisol from Amazon corresponded on a 1:1 basis with the sulfate adsorption by that soil. Okuma and Alves [14] observed, in the absence of phosphate, that the enhancing effect of sulfate on CEC increased with the soil contents of iron and aluminum oxides. Although Pearce and Sumner [11] have suggested that the equivalent retention of

mechanisms, including precipitation, specific adsorption, ionic strength-induced charging, and ion-pair adsorption, to date such phenomenon remains to be clarified. Among the surface chemical equations used to model solution pH and ionic strength effects on ion complexation on a standard gibbsite surface (≡AlOH0

those referring to sulfate complexation are given by [15]

≡ AlOH<sup>0</sup> + SO4

≡ AlOH<sup>0</sup> + SO4

) in the soil solution as predicted by the hard and soft (Lewis)

<sup>2</sup><sup>−</sup> observed in kaolinitic subsoil could be due the combination of several

2− + H<sup>+</sup> ↔ ≡ AlSO4

The above equations indicate that sulfate adsorption to gibbsite gives rise to negatively charged surface species, decreases the net positive surface charge, and increases the solution pH. Such features could favor sulfate-cation co-adsorption even at lower pH because sulfate adsorption on gibbsite increases as pH decreases [16]. Indeed, the acidic condition is required for the formation of ternary surface complexes of copper,

2− ↔ ≡ AlOHSO4

−,

),

<sup>−</sup> + H2O (1)

2− (2)

**42**

A kaolinite sample from Minas Gerais, Brazil, and a laboratory-prepared sample of gibbsite were used in the experiments. The gibbsite sample was prepared via titration of 500 mL of 1 M Al3+ (added as AlCl3.6H2O) with 315 mL of 4 M NaOH followed by precipitate dialysis for 1 month against daily refreshed deionized water kept at 50°C [26]. Both samples were washed with deionized water, oven-dried (75°C/48 h), ground in an agate mortar, and passed through a 0.3-mm sieve.

The mineralogical purity of the samples was checked through powder X-ray diffraction (XRD). The XRD patterns were measured using a Philips PW1830 diffractometer (PANalytical, Almelo, the Netherlands) in continuous scan mode (0.02°2θ s<sup>−</sup><sup>1</sup> ) from 3 to 90°2θ CuKα. The particle shapes were examined via scanning electron microscopy (SEM) using a LEO 435 VP microscope (LEO Electron Microscopy Inc., Thornwood, NY). The pH at the isoelectric point (IEP) was determined by graphical interpolation from zeta-potential (ζ) measurements. The ζ values were measured in triplicate with a NanoBrook Omni analyzer (Brookhaven Instruments, Holtsville, NY) in mineral suspensions (1 g L<sup>−</sup><sup>1</sup> ) prepared in 0.01 M KNO3 solutions with different pH values,

which were stabilized by additions of 0.01 M HCl or 0.01 M NaOH. The five-point N2-BET-specific surface areas were measured in the samples with a Micromeritics ASAP 2010 surface area analyzer (Micromeritics, Norcross, GA).

#### **2.2 Calcium adsorption experiments**

Before the adsorption experiments, aqueous mineral suspensions (20 g L<sup>−</sup><sup>1</sup> ) were titrated with 0.1 M HCl or 0.1 M NaOH until their pH values stabilized at 4 and 6 to mimic an unsuitable (pH 4) and a proper (pH 6) acidity condition for the most of crops. Then, the suspensions were centrifuged, and the minerals were oven-dried at 75°C for 48 h, ground in an agate mortar, and passed through a 0.3-mm sieve. Calcium solutions (0.5 and 10 mM; pH 4 and 6) were prepared from reagent-grade calcium sulfate and calcium nitrate. The above concentrations were chosen to allow for respective low and high ion pairing in the sulfate solutions. Aqueous calcium speciation (**Table 1**) was calculated using the formation constants available in the built-in database (NIST 46.7) of Visual MINTEQ [27].

Batch calcium adsorption experiments were performed without CO2 exclusion as follows: 0.2 g of mineral and 20 mL of solution were placed into 50-mL polypropylene centrifuge tubes. The suspensions were shaken end over end at 30 rpm for 24 h at 20 ± 2°C. The pH was measured in the suspensions with a Thermo Orion 4-star pH meter (Thermo Orion Inc., Beverly, MA) after a two-point calibration with standard buffer solutions (pH 4.0 and 7.0). The tubes were centrifuged for 30 min at 10,000 rpm, and the supernatants were filtered through 0.22-μm cellulose membranes prior to calcium analysis in an atomic absorption spectrophotometer (Varian AA 240 FS, Agilent Technologies, Mulgrave, Australia). A cross-check of the initial calcium concentrations in the applied solutions was concurrent with the measurements performed in the remaining solutions. Both absolute and relative amounts of adsorbed calcium were calculated from the difference between the calcium initial and equilibrium aqueous concentrations.


*a I = 1/2 ×* ∑*Mizi 2 , where I is the solution ionic strength calculated via Visual MINTEQ, mM; Mi and zi are the respective concentration, mM, and charge of the dissolved specie i.*

*b Proportions of the main dissolved calcium species calculated via Visual MINTEQ in relation to the total calcium concentration.*

*c Free calcium activity in solution calculated via Visual MINTEQ from activity coefficients assessed with the Davies equation.*

**45**

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration…*

Electrokinetic experiments were carried out for kaolinite and gibbsite suspen-

sodium sulfate. Additional measurements were carried out with calcium sulfate solutions including a saturated one (~15.8 mM) to maximize the formation of the

were freshly prepared and at least 10 runs per sample were performed. These series clearly indicated the absence of time dependence, i.e., equilibration was typically on

Second harmonic generation (SHG) experiments were carried out with an alumina prism exposing the c-cut to the aqueous solution. This crystal plane is structurally equivalent to the gibbsite basal plane on both gibbsite and kaolinite. The setup has been described in detail elsewhere [28, 29]. The sapphire prism was obtained from Kyburz (Safnern, Switzerland) and cleaned based on procedures established earlier. The PTFE cell was initially cleaned as well with acetone and ethanol and copiously washed with MilliQ water. After each experimental series, the solution in the cell was replaced several times with MilliQ water until the signal

Uptake experiments were carried out in triplicate. The paired *t*-test was applied to evaluate if the mean difference between the amounts of calcium adsorbed to each mineral in the presence of nitrate and sulfate, respectively, differed from zero at *P* = 0.05. Calculations were performed using SAS software, version 9.3 [30].

The electrokinetic experiments involved at least 10 runs per sample. From this the averaged electrophoretic mobility was used, and the standard error was noted as

The second harmonic generation experiments allow clear observation of equilib-

rium states as well. As for the electrokinetic experiments, transient effects were absent when the reported data points were collected. The data were scaled to the signal in pure water, which was measured before and after a concentration series. The temperature dependence was studied with a setup that was previously used in

The XRD pattern*s* (**Figure 1**) show some accessory quartz with kaolinite and no other phase with gibbsite [31]. The plate-shaped morphologies of the mineral particles conform to the pseudohexagonal plates of kaolinite [32] and the plate-like

the density and particle dimensions of a gibbsite sample prepared using the same

g<sup>−</sup><sup>1</sup>

reported for this mineral [34]. The gibbsite surface

) is close to that calculated by Rosenqvist and Casey [35] from

) is within the

lozenges or plate-like gibbsite hexagons [33] (**Figure 2**).

g<sup>−</sup><sup>1</sup>

The specific surface area of the kaolinite sample (30.9 m<sup>2</sup>

, pH 5) presenting increasing amounts of either calcium chloride or

ion pair. The Brookhaven PALS apparatus was used. Suspensions

*DOI: http://dx.doi.org/10.5772/intechopen.81273*

**2.4 Second harmonic generation measurements**

obtained remained stable at the initial value of MilliQ water.

**2.3 Electrokinetic measurements**

0

sions (1 g L<sup>−</sup><sup>1</sup>

aqueous CaSO4

the order of 5 min.

**2.5 Data analysis**

specified by the software.

ice-nucleation studies.

**3. Results and discussion**

**3.1 Mineral characterization**

interval from 5 to 39 m<sup>2</sup>

g<sup>−</sup><sup>1</sup>

area (26.6 m<sup>2</sup>

#### **Table 1.**

*Calcium solutions used in the sorption experiments.*

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration… DOI: http://dx.doi.org/10.5772/intechopen.81273*

#### **2.3 Electrokinetic measurements**

*Advanced Sorption Process Applications*

**2.2 Calcium adsorption experiments**

and equilibrium aqueous concentrations.

**Solution pH I**

*respective concentration, mM, and charge of the dissolved specie i.*

*Calcium solutions used in the sorption experiments.*

MINTEQ [27].

which were stabilized by additions of 0.01 M HCl or 0.01 M NaOH. The five-point N2-BET-specific surface areas were measured in the samples with a Micromeritics ASAP

Before the adsorption experiments, aqueous mineral suspensions (20 g L<sup>−</sup><sup>1</sup>

Batch calcium adsorption experiments were performed without CO2 exclusion as follows: 0.2 g of mineral and 20 mL of solution were placed into 50-mL polypropylene centrifuge tubes. The suspensions were shaken end over end at 30 rpm for 24 h at 20 ± 2°C. The pH was measured in the suspensions with a Thermo Orion 4-star pH meter (Thermo Orion Inc., Beverly, MA) after a two-point calibration with standard buffer solutions (pH 4.0 and 7.0). The tubes were centrifuged for 30 min at 10,000 rpm, and the supernatants were filtered through 0.22-μm cellulose membranes prior to calcium analysis in an atomic absorption spectrophotometer (Varian AA 240 FS, Agilent Technologies, Mulgrave, Australia). A cross-check of the initial calcium concentrations in the applied solutions was concurrent with the measurements performed in the remaining solutions. Both absolute and relative amounts of adsorbed calcium were calculated from the difference between the calcium initial

0.5 mM CaSO4.2H2O 4.0 02 93.3 06.7 - 0.38 0.5 mM CaSO4.2H2O 6.0 02 93.2 06.2 - 0.38 0.5 mM Ca(NO3)2.4H2O 4.0 02 99.7 - 0.3 0.42 0.5 mM Ca(NO3)2.4H2O 6.0 02 99.7 - 0.3 0.42 10 mM CaSO4.2H2O 4.0 30 69.3 30.7 - 3.68 10 mM CaSO4.2H2O 6.0 30 69.2 30.8 - 3.68 10 mM Ca(NO3)2.4H2O 4.0 30 96.8 - 3.2 5.07 10 mM Ca(NO3)2.4H2O 6.0 30 96.8 - 3.2 5.07

*, where I is the solution ionic strength calculated via Visual MINTEQ, mM; Mi and zi are the* 

*Proportions of the main dissolved calcium species calculated via Visual MINTEQ in relation to the total calcium* 

*Free calcium activity in solution calculated via Visual MINTEQ from activity coefficients assessed with the Davies* 

**<sup>a</sup> Ca2+ CaSO4**

**<sup>0</sup> CaNO3**

**mM %b mM**

**<sup>+</sup> Ca2+ <sup>c</sup>**

were titrated with 0.1 M HCl or 0.1 M NaOH until their pH values stabilized at 4 and 6 to mimic an unsuitable (pH 4) and a proper (pH 6) acidity condition for the most of crops. Then, the suspensions were centrifuged, and the minerals were oven-dried at 75°C for 48 h, ground in an agate mortar, and passed through a 0.3-mm sieve. Calcium solutions (0.5 and 10 mM; pH 4 and 6) were prepared from reagent-grade calcium sulfate and calcium nitrate. The above concentrations were chosen to allow for respective low and high ion pairing in the sulfate solutions. Aqueous calcium speciation (**Table 1**) was calculated using the formation constants available in the built-in database (NIST 46.7) of Visual

)

2010 surface area analyzer (Micromeritics, Norcross, GA).

**44**

*a*

*b*

*c*

*I = 1/2 ×* ∑*Mizi*

*concentration.*

*equation.*

**Table 1.**

*2*

Electrokinetic experiments were carried out for kaolinite and gibbsite suspensions (1 g L<sup>−</sup><sup>1</sup> , pH 5) presenting increasing amounts of either calcium chloride or sodium sulfate. Additional measurements were carried out with calcium sulfate solutions including a saturated one (~15.8 mM) to maximize the formation of the aqueous CaSO4 0 ion pair. The Brookhaven PALS apparatus was used. Suspensions were freshly prepared and at least 10 runs per sample were performed. These series clearly indicated the absence of time dependence, i.e., equilibration was typically on the order of 5 min.

#### **2.4 Second harmonic generation measurements**

Second harmonic generation (SHG) experiments were carried out with an alumina prism exposing the c-cut to the aqueous solution. This crystal plane is structurally equivalent to the gibbsite basal plane on both gibbsite and kaolinite. The setup has been described in detail elsewhere [28, 29]. The sapphire prism was obtained from Kyburz (Safnern, Switzerland) and cleaned based on procedures established earlier. The PTFE cell was initially cleaned as well with acetone and ethanol and copiously washed with MilliQ water. After each experimental series, the solution in the cell was replaced several times with MilliQ water until the signal obtained remained stable at the initial value of MilliQ water.

#### **2.5 Data analysis**

Uptake experiments were carried out in triplicate. The paired *t*-test was applied to evaluate if the mean difference between the amounts of calcium adsorbed to each mineral in the presence of nitrate and sulfate, respectively, differed from zero at *P* = 0.05. Calculations were performed using SAS software, version 9.3 [30].

The electrokinetic experiments involved at least 10 runs per sample. From this the averaged electrophoretic mobility was used, and the standard error was noted as specified by the software.

The second harmonic generation experiments allow clear observation of equilibrium states as well. As for the electrokinetic experiments, transient effects were absent when the reported data points were collected. The data were scaled to the signal in pure water, which was measured before and after a concentration series. The temperature dependence was studied with a setup that was previously used in ice-nucleation studies.

## **3. Results and discussion**

#### **3.1 Mineral characterization**

The XRD pattern*s* (**Figure 1**) show some accessory quartz with kaolinite and no other phase with gibbsite [31]. The plate-shaped morphologies of the mineral particles conform to the pseudohexagonal plates of kaolinite [32] and the plate-like lozenges or plate-like gibbsite hexagons [33] (**Figure 2**).

The specific surface area of the kaolinite sample (30.9 m<sup>2</sup> g<sup>−</sup><sup>1</sup> ) is within the interval from 5 to 39 m<sup>2</sup> g<sup>−</sup><sup>1</sup> reported for this mineral [34]. The gibbsite surface area (26.6 m<sup>2</sup> g<sup>−</sup><sup>1</sup> ) is close to that calculated by Rosenqvist and Casey [35] from the density and particle dimensions of a gibbsite sample prepared using the same

**Figure 1.** *X-ray diffraction patterns of (a) kaolinite and (b) gibbsite samples. Qz = quartz.*

**Figure 2.** *Scanning electron micrographs of (a) kaolinite and (b) gibbsite samples.*

procedure followed here (25 m<sup>2</sup> g<sup>−</sup><sup>1</sup> ). However, their reported N2-BET surface area is somewhat lower (19.6 m<sup>2</sup> g<sup>−</sup><sup>1</sup> ).

The respective IEP values found for kaolinite and gibbsite are 4.9 and 10.6 (**Figure 3**). The kaolinite IEP is close to the value compiled by Stumm and Morgan [36] (4.6). The gibbsite one is in between the pristine point of zero charge at pH 10.0 found by Hiemstra et al. [26] and the IEP at pH 11.3 reported by Adekola et al. [37] for gibbsites with surface areas above 25 m2 g<sup>−</sup><sup>1</sup> .

#### **3.2 Calcium adsorption on kaolinite**

The batch experiments were carried out without attempting to exclude CO2. In principle, calcite bulk precipitation could also decrease the equilibrium aqueous calcium concentrations. However, according to equilibrium calculations, the pH threshold value needed for calcite formation in the 10-mM calcium solutions in the absence of calcium adsorption is about 7.7. Considering that this pH value increases as the initial calcium concentration in solution decreases, calcite formation in our mineral suspensions can be safely excluded because the highest equilibrium pH was 6.9. Furthermore, we assumed that the aqueous bicarbonate (HCO3 <sup>−</sup>) (~2.4 μM as calculated with Visual MINTEQ ) did not affect the sulfate adsorption. Helyar et al. [38] have found that 0.58 mM HCO3 <sup>−</sup> did not influence phosphate adsorption on a gibbsite sample.

The amounts of calcium adsorbed to kaolinite from the dilute and concentrated solutions are presented in **Table 2**. The results for the dilute solutions (0.5 mM) with initial pH 4 showed that higher sulfate adsorption on kaolinite at low pH [16] did not enhance calcium adsorption via potential co-adsorption. This finding

**47**

**Figure 3.**

*a*

**Table 2.**

*sulfate.*

*Calculated in relation to the initial calcium concentration.*

*concentrated solutions (10 mM).*

*solutions at initial pH 6.*

*IEP = isoelectric point.*

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration…*

*Zeta potential values measured in kaolinite and gibbsite 0.01-M KNO3 suspensions at different pH conditions.* 

0.5 mM Ca(NO3)2.4H2O 4.0 4.2 0.0a 00.0a 0.5 mM CaSO4.2H2O 4.0 4.4 0.0a 00.0a 0.5 mM Ca(NO3)2.4H2O 6.0 5.1 0.4*b* 21.7*b* 0.5 mM CaSO4.2H2O 6.0 5.2 0.4*b* 22.7*b* 10 mM Ca(NO3)2.4H2O 4.0 4.0 3.2A 09.4A 10 mM CaSO4.2H2O 4.0 4.1 1.6B 05.3B 10 mM Ca(NO3)2.4H2O 6.0 4.7 1.6*A* 05.1*A* 10 mM CaSO4.2H2O 6.0 4.9 0.8*B* 02.6*B*

**Initial Equilibrium μmol m<sup>−</sup><sup>2</sup> (%)a**

**Solution pH Ca2+**

suggests that a possible increase in negative surface charges following from sulfate adsorption was insufficient to overcome the electrostatic repulsion to calcium ions exerted by positive surface charges. In this case, positive charges superseded negative ones because the equilibrium pH values remained below the kaolinite IEP. These results also indicate that even starting to deprotonate at pH 3 [39], the silanol surface groups of kaolinite did not adsorb calcium. The local electrostatic repulsion from the positively charged aluminol surface groups of kaolinite, which deprotonate above pH 8.7 [39], may have outweighed the attraction exerted by the negatively charged silanol groups on calcium at low pH and low ionic strength. Calcium adsorption to kaolinite from the 0.5-mM solutions with initial pH 6 was

*Normal letters refer to comparisons between solutions at initial pH 4; italicized letters refer to comparisons between* 

*Initial and equilibrium pH values and amounts of calcium adsorbed on kaolinite in the presence of nitrate and* 

*Means followed by the same letter in the same column do not differ at P = 0.05 according to the paired t-test. Lowercase letters refer to comparisons of dilute solutions (0.5 mM); uppercase letters refer to comparisons of* 

*DOI: http://dx.doi.org/10.5772/intechopen.81273*

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration… DOI: http://dx.doi.org/10.5772/intechopen.81273*

#### **Figure 3.**

*Advanced Sorption Process Applications*

**Figure 1.**

**Figure 2.**

procedure followed here (25 m<sup>2</sup>

area is somewhat lower (19.6 m<sup>2</sup>

**3.2 Calcium adsorption on kaolinite**

[38] have found that 0.58 mM HCO3

g<sup>−</sup><sup>1</sup>

*X-ray diffraction patterns of (a) kaolinite and (b) gibbsite samples. Qz = quartz.*

 g<sup>−</sup><sup>1</sup> ). The respective IEP values found for kaolinite and gibbsite are 4.9 and 10.6 (**Figure 3**). The kaolinite IEP is close to the value compiled by Stumm and Morgan [36] (4.6). The gibbsite one is in between the pristine point of zero charge at pH 10.0 found by Hiemstra et al. [26] and the IEP at pH 11.3 reported by Adekola

6.9. Furthermore, we assumed that the aqueous bicarbonate (HCO3

The batch experiments were carried out without attempting to exclude CO2. In principle, calcite bulk precipitation could also decrease the equilibrium aqueous calcium concentrations. However, according to equilibrium calculations, the pH threshold value needed for calcite formation in the 10-mM calcium solutions in the absence of calcium adsorption is about 7.7. Considering that this pH value increases as the initial calcium concentration in solution decreases, calcite formation in our mineral suspensions can be safely excluded because the highest equilibrium pH was

calculated with Visual MINTEQ ) did not affect the sulfate adsorption. Helyar et al.

The amounts of calcium adsorbed to kaolinite from the dilute and concentrated solutions are presented in **Table 2**. The results for the dilute solutions (0.5 mM) with initial pH 4 showed that higher sulfate adsorption on kaolinite at low pH [16] did not enhance calcium adsorption via potential co-adsorption. This finding

et al. [37] for gibbsites with surface areas above 25 m2

*Scanning electron micrographs of (a) kaolinite and (b) gibbsite samples.*

). However, their reported N2-BET surface

 g<sup>−</sup><sup>1</sup> .

<sup>−</sup> did not influence phosphate adsorption on a

<sup>−</sup>) (~2.4 μM as

**46**

gibbsite sample.

*Zeta potential values measured in kaolinite and gibbsite 0.01-M KNO3 suspensions at different pH conditions. IEP = isoelectric point.*


*a Calculated in relation to the initial calcium concentration.*

*Means followed by the same letter in the same column do not differ at P = 0.05 according to the paired t-test. Lowercase letters refer to comparisons of dilute solutions (0.5 mM); uppercase letters refer to comparisons of concentrated solutions (10 mM).*

*Normal letters refer to comparisons between solutions at initial pH 4; italicized letters refer to comparisons between solutions at initial pH 6.*

#### **Table 2.**

*Initial and equilibrium pH values and amounts of calcium adsorbed on kaolinite in the presence of nitrate and sulfate.*

suggests that a possible increase in negative surface charges following from sulfate adsorption was insufficient to overcome the electrostatic repulsion to calcium ions exerted by positive surface charges. In this case, positive charges superseded negative ones because the equilibrium pH values remained below the kaolinite IEP. These results also indicate that even starting to deprotonate at pH 3 [39], the silanol surface groups of kaolinite did not adsorb calcium. The local electrostatic repulsion from the positively charged aluminol surface groups of kaolinite, which deprotonate above pH 8.7 [39], may have outweighed the attraction exerted by the negatively charged silanol groups on calcium at low pH and low ionic strength. Calcium adsorption to kaolinite from the 0.5-mM solutions with initial pH 6 was

the same in the presence of either nitrate or sulfate. Therefore, no sulfate-calcium co-adsorption occurred under the net negative surface charge condition provided by equilibrium pH values (5.1 and 5.2) higher than the kaolinite IEP (4.9).

The increase in the initial concentration of calcium sulfate (10 mM) reduced calcium adsorption on kaolinite by 50% relative to those measured in the presence of 20 mmolc L<sup>−</sup><sup>1</sup> of nitrate for the two initial pH conditions. Equilibrium calculations (**Table 1**) indicate that the neutral pair CaSO4 0 comprises 31% of the total calcium dissolved in the 10 mM calcium sulfate solution for both pH values, which decreases free calcium ion activity in the sulfate solutions to about 63% of those calculated for the nitrate solutions. The enhanced formation of neutral ion pairs containing an adsorptive and/or lowering the activity of its ionic free form in solution may decrease its adsorption [23, 40] and outweigh potential co-adsorption. Furthermore, the reduction of Ca2+ and SO4 <sup>2</sup><sup>−</sup> activities in solution due to adsorption did not seem to promote an appreciable shift in the equilibrium of CaSO4 0 formation toward free aqueous Ca2+ and SO4 <sup>2</sup><sup>−</sup>, which could lead to comparable conditions of free ions in the nitrate solutions. Finally, unlike the observed for the dilute solutions presenting initial pH 4, the enhancement of calcium loading to 10 mM resulted in calcium adsorption on kaolinite overcoming the electrostatic repulsion from the net positive surface charge of that mineral at equilibrium.

### **3.3 Calcium adsorption on gibbsite**

The amounts of calcium adsorbed on gibbsite from all solutions are presented in **Table 3**. Although according to Eqs. (1) and (2) sulfate adsorption to gibbsite could favor sulfate-calcium coadsorption, such a trend was not observed for the dilute and acidic conditions (0.5 mM; pH 4). At equilibrium pH ~ 7, calcium adsorption on gibbsite from the dilute solutions only occurred in the presence of sulfate. Because of the high IEP of the studied gibbsite, all adsorption experiments corresponded to a net positive surface charge. Sulfate adsorption and pH enhancement reduce the positive charges and concomitant electrostatic repulsion of calcium. Therefore,


*a Calculated in relation to the initial calcium concentration.*

*Means followed by the same letter in the same column do not differ at P = 0.05 according to the paired t-test. Lowercase letters refer to comparisons of dilute solutions (0.5 mM); uppercase letters refer to comparisons of concentrated solutions (10 mM).*

*Normal letters refer to comparisons between solutions at initial pH 4; italicized letters refer to comparisons between solutions at initial pH 6.*

#### **Table 3.**

*Initial and equilibrium pH values and amounts of calcium adsorbed on gibbsite in the presence of nitrate and sulfate.*

**49**

mineral [20].

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration…*

our results suggest that besides sulfate adsorption, the net positive surface charge, which also depends on solution pH and ionic strength, affected calcium sulfate co-adsorption on gibbsite. Sulfate also decreased calcium adsorption on gibbsite from the more concentrated acidic solutions (10 mM; pH = 4), presumably due to

**4. pH and salt concentration effects on calcium adsorption on kaolinite** 

Unlike the enhancing pH effect on calcium adsorption on kaolinite observed for the 0.5-mM solutions containing either nitrate or sulfate, an opposite effect was found for the two more concentrated solutions (10 mM) containing those anions. In these cases, an initial pH increase from 4 to 6 decreased absolute calcium adsorp-

This negative pH effect on calcium adsorption from the 10-mM solutions can be considered analogous to the positive pH effect on orthophosphate adsorption on kaolinite observed in previous papers for acidic condition (pH < 7) [41–43]. Such an effect concurs with negative charge enhancement on the kaolinite surface, which disfavors orthophosphate adsorption. On the other hand, the orthophosphate adsorption to kaolinite from dilute solutions decreases as solution pH increases, which resembles the P adsorption behavior of iron oxides [44]. He et al. [45] suggested that the aqueous P speciation may be related to the positive pH effect on orthophosphate adsorption to kaolinite. However, such an association does not seem plausible: according to equilibrium calculations based on NIST 46.7 stability

ranging from 3 to 7. Likewise, for pH ranging from 4 to 7, calcium speciation in the studied 10-mM solutions does not differ from results in **Table 1**. Therefore, higher salt concentration should be associated with low calcium adsorption at higher pH

While the basal planes of kaolinite have been considered to hold permanent negative charges [46], recent surface force measurements carried out with atomic force microscopes have shown that the surface charges of the silica basal plane of kaolinite and the alumina basal plane of kaolinite and gibbsite react to changes in solution pH and salt concentration as the variable charges found on edge faces of those minerals [19–21]. Furthermore, the solution effects on basal silica surface charges differ from those observed for alumina faces [19, 21], and even for a given basal plane, the magnitudes of the surface charges can change when the dissolved cation constitutes the sole difference among the solutions in contact with the

Such complex, anisotropic charge behavior may be relevant to our findings given the apparent unexpected negative pH effect on calcium adsorption. Weak calcium hydrolysis suggests that uptake on oxide-like surface groups, such as those on kaolinite and gibbsite edges, only occurs at relatively high pH [15, 47]. Therefore, the basal planes of kaolinite and gibbsite may be the main adsorption sites of that cation at pH < 7 [19, 21]. Although Siretanu et al. [20] observed calcium adsorption from CaCl2 solutions at pH 6 for the alumina face of nanosized gibbsite, their surface force measurements indicated that calcium adsorption from solutions with increasing concentrations of that cation initially increased followed by a decrease in the extracted charge densities with a maximum between 5 and 10 mM CaCl2. The authors argued that the concurrent increasing co-adsorption of chloride ions could explain the decrease in surface force. Unfortunately, the amounts of adsorbed calcium cannot be assessed in such experiments as from batch adsorption studies. On

<sup>−</sup>) prevails from pH

*DOI: http://dx.doi.org/10.5772/intechopen.81273*

**and gibbsite**

the same mechanisms as those proposed for kaolinite.

tion by 50% on kaolinite and by 100% on gibbsite.

values. This hypothesis will be discussed below.

constant database, the aqueous orthophosphate species (H2PO4

*Calcium Uptake on Kaolinite and Gibbsite: Effects of Sulfate, pH, and Salt Concentration… DOI: http://dx.doi.org/10.5772/intechopen.81273*

our results suggest that besides sulfate adsorption, the net positive surface charge, which also depends on solution pH and ionic strength, affected calcium sulfate co-adsorption on gibbsite. Sulfate also decreased calcium adsorption on gibbsite from the more concentrated acidic solutions (10 mM; pH = 4), presumably due to the same mechanisms as those proposed for kaolinite.
