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

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

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 constant database, the aqueous orthophosphate species (H2PO4 <sup>−</sup>) prevails from pH 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 values. This hypothesis will be discussed below.

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 mineral [20].

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

*Advanced Sorption Process Applications*

ence of 20 mmolc L<sup>−</sup><sup>1</sup>

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

The increase in the initial concentration of calcium sulfate (10 mM) reduced calcium adsorption on kaolinite by 50% relative to those measured in the pres-

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.

tion did not seem to promote an appreciable shift in the equilibrium of CaSO4

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.

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,

**Solution pH Ca2+**

*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* 

*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 gibbsite in the presence of nitrate and* 

0.5 mM Ca(NO3)2.4H2O 4.0 4.3 0.0a 0.0a 0.5 mM CaSO4.2H2O 4.0 4.5 0.0a 0.0a 0.5 mM Ca(NO3)2.4H2O 6.0 6.5 0.0*a* 0.0*a* 0.5 mM CaSO4.2H2O 6.0 6.9 0.2*b* 8.2*b* 10 mM Ca(NO3)2.4H2O 4.0 4.4 1.9A 4.3A 10 mM CaSO4.2H2O 4.0 4.6 0.9B 2.6B 10 mM Ca(NO3)2.4H2O 6.0 6.2 0.0*A* 0.0*A* 10 mM CaSO4.2H2O 6.0 6.8 0.0*B* 0.0*B*

of nitrate for the two initial pH conditions. Equilibrium

0

<sup>2</sup><sup>−</sup> activities in solution due to adsorp-

<sup>2</sup><sup>−</sup>, which could lead to comparable

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

comprises 31% of the

0

by equilibrium pH values (5.1 and 5.2) higher than the kaolinite IEP (4.9).

calculations (**Table 1**) indicate that the neutral pair CaSO4

Furthermore, the reduction of Ca2+ and SO4

formation toward free aqueous Ca2+ and SO4

*Calculated in relation to the initial calcium concentration.*

*concentrated solutions (10 mM).*

*solutions at initial pH 6.*

**3.3 Calcium adsorption on gibbsite**

**48**

**Table 3.**

*sulfate.*

*a*

the other hand, batch adsorption results cannot be associated with a given surface plane. The maximum diffuse layer potential from force measurements can likewise be explicated by constant or increasing calcium adsorption overcompensated by screening due to chloride adsorption in the Stern layer or simple double-layer screening.

In our experiments with the 10-mM solutions, the mean potential differences across the electrical double layer of gibbsite at equilibrium, as estimated by the Nernst Equation [48], were + 360 and + 242 mV for the more and less acidic solutions, respectively. These values, which are upper limit estimates, indicate that anion adsorption should be disfavored in the less acidic suspension. In the kaolinite scenarios with the 10-mM solutions, the mean Nernst surface potentials were also positive at equilibrium, i.e., +50 and + 5.6 mV for more and less acidic solutions, respectively. The lower values compared to those calculated for gibbsite, and a possible difference between the screening effects on silica and alumina basal planes, could cause a moderate reduction in calcium adsorption on kaolinite from the less acidic nitrate and sulfate solutions. For the experiments with the dilute solutions (0.5 mM), the low salt concentration was insufficient for promoting appreciable screening effects.

### **4.1 Electrokinetic measurements**

**Figure 4** displays the results of electrokinetic measurements carried out for kaolinite and gibbsite suspensions at pH 5. In general, the electrokinetic mobility (EM) values measured against increasing salt concentrations showed the same trends for both studied minerals. The increasing sulfate addition as Na2SO4 enhanced the negative surface charge of both kaolinite and gibbsite (**Figure 4(a)**). However, in the gibbsite scenario, sulfate caused charge reversal at the concentration of about 0.3 mM Na2SO4. Data obtained via the increasing addition of CaCl2 to kaolinite and gibbsite in the presence of 3 mM Na2SO4 demonstrated rather a shielding effect, mainly for the gibbsite scenario where no clear charge reversal was observed at 3 mM Ca2+ (**Figure 4(b)**). In a mixed CaSO4 system up to the solubility limit, we retrieved the sulfate effect of increasing the surface negative charge (**Figure 4(c)**). Electrophoretic mobility measurements are usually unstable when performed at pH value close to the mineral IEP; such an instability probably resulted in EM values somewhat different for kaolinite in the absence of Na2SO4 (**Figure 4(a)**) and CaSO4 (**Figure 4(c)**). Despite this, the obtained results corroborate the idea that co-adsorption of calcium due to sulfate adsorption is a minor feature.

#### **4.2 Surface charge modeling applied to gibbsite**

Some of our assumptions also agree with gibbsite surface charge trends simulated with the diffuse layer model (DLM) [49] using the calcium/gibbsite and sulfate/gibbsite surface complexation constants available in [15] as well as Eqs. (1)–(3). The simulations were performed for standard gibbsite suspensions (10 g L<sup>−</sup><sup>1</sup> ) in solutions containing dual cation-anion combinations with sodium, calcium, nitrate, and sulfate, where nitrate and sodium ions do not interact with the surfaces via chemical reactions. Binary systems were chosen as references of both inert ions (sodium and nitrate) and their mixtures with sulfate and calcium. Although the standard gibbsite as treated by Karamalidis and Dzombak [15] has an IEP of 9 (i.e., lower than in the present case), the calculation results illuminate calcium adsorption in the mixed system. Unfortunately, self-consistent data for a more complex surface complexation model are not available. The authors are well

**51**

**Figure 4.**

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

aware about the drawback of using a simple DLM. With the available data, it was not possible to design a separate model. The present modeling, for example, disregards the available information about the nature of the surface complexes (inner sphere vs. outer sphere) or the calcium adsorption on the basal planes, because the

*of (a) Na2SO4, (b) CaCl2 + 3 mM Na2SO4, and (c) CaSO4.*

*Electrophoretic mobility values of kaolinite and gibbsite suspensions measured under increasing concentrations* 

*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 4.**

*Advanced Sorption Process Applications*

screening.

screening effects.

**4.1 Electrokinetic measurements**

**4.2 Surface charge modeling applied to gibbsite**

the other hand, batch adsorption results cannot be associated with a given surface plane. The maximum diffuse layer potential from force measurements can likewise be explicated by constant or increasing calcium adsorption overcompensated by screening due to chloride adsorption in the Stern layer or simple double-layer

In our experiments with the 10-mM solutions, the mean potential differences across the electrical double layer of gibbsite at equilibrium, as estimated by the Nernst Equation [48], were + 360 and + 242 mV for the more and less acidic solutions, respectively. These values, which are upper limit estimates, indicate that anion adsorption should be disfavored in the less acidic suspension. In the kaolinite scenarios with the 10-mM solutions, the mean Nernst surface potentials were also positive at equilibrium, i.e., +50 and + 5.6 mV for more and less acidic solutions, respectively. The lower values compared to those calculated for gibbsite, and a possible difference between the screening effects on silica and alumina basal planes, could cause a moderate reduction in calcium adsorption on kaolinite from the less acidic nitrate and sulfate solutions. For the experiments with the dilute solutions (0.5 mM), the low salt concentration was insufficient for promoting appreciable

**Figure 4** displays the results of electrokinetic measurements carried out for kaolinite and gibbsite suspensions at pH 5. In general, the electrokinetic mobility (EM) values measured against increasing salt concentrations showed the same trends for both studied minerals. The increasing sulfate addition as Na2SO4 enhanced the negative surface charge of both kaolinite and gibbsite (**Figure 4(a)**). However, in the gibbsite scenario, sulfate caused charge reversal at the concentration of about 0.3 mM Na2SO4. Data obtained via the increasing addition of CaCl2 to kaolinite and gibbsite in the presence of 3 mM Na2SO4 demonstrated rather a shielding effect, mainly for the gibbsite scenario where no clear charge reversal was observed at 3 mM Ca2+ (**Figure 4(b)**). In a mixed CaSO4 system up to the solubility limit, we retrieved the sulfate effect of increasing the surface negative charge (**Figure 4(c)**). Electrophoretic mobility measurements are usually unstable when performed at pH value close to the mineral IEP; such an instability probably resulted in EM values somewhat different for kaolinite in the absence of Na2SO4 (**Figure 4(a)**) and CaSO4 (**Figure 4(c)**). Despite this, the obtained results corroborate the idea that co-adsorption of calcium due to sulfate adsorption is a minor

Some of our assumptions also agree with gibbsite surface charge trends simulated with the diffuse layer model (DLM) [49] using the calcium/gibbsite and sulfate/gibbsite surface complexation constants available in [15] as well as Eqs. (1)–(3). The simulations were performed for standard gibbsite suspensions

calcium, nitrate, and sulfate, where nitrate and sodium ions do not interact with the surfaces via chemical reactions. Binary systems were chosen as references of both inert ions (sodium and nitrate) and their mixtures with sulfate and calcium. Although the standard gibbsite as treated by Karamalidis and Dzombak [15] has an IEP of 9 (i.e., lower than in the present case), the calculation results illuminate calcium adsorption in the mixed system. Unfortunately, self-consistent data for a more complex surface complexation model are not available. The authors are well

) in solutions containing dual cation-anion combinations with sodium,

**50**

feature.

(10 g L<sup>−</sup><sup>1</sup>

*Electrophoretic mobility values of kaolinite and gibbsite suspensions measured under increasing concentrations of (a) Na2SO4, (b) CaCl2 + 3 mM Na2SO4, and (c) CaSO4.*

aware about the drawback of using a simple DLM. With the available data, it was not possible to design a separate model. The present modeling, for example, disregards the available information about the nature of the surface complexes (inner sphere vs. outer sphere) or the calcium adsorption on the basal planes, because the DLM only includes inner-sphere surface complexes and the model does not distinguish crystal planes. Also when the data base was generated, the information about the Ca adsorption at low pH had not been available. Furthermore, in the following the comparisons between solutions containing 1:1, 2:1 and 2:2 electrolyte solutions are difficult to directly compare, since at the same time the ionic strength and the cation concentrations might be changing for a given concentration of the electrolyte (this is also true for the subsequent sections).

The enhancement of an inert electrolyte concentration increases the gibbsite charge at a fixed pH with no shift in the IEP (**Figure 5(a)**). The diffuse layer model in the database used for the calculations causes a steep charge increase due to the absence of a Stern layer. This behavior is hidden when plotting titration raw data (i.e., total acid and base added versus pH). As depicted in **Figure 5(b)**, calcium does not significantly affect the curves relative to **Figure 5(a)**. The charge for a given cation concentration and pH increases because the negative counterion concentration is higher in the calcium presence, allowing for better shielding and increased charge. The enhanced charge is not due to calcium adsorption, which is insignificant.

**Figure 5(c)** shows two effects relative to the gibbsite surface charge as a function of sodium sulfate concentration. One is sulfate uptake, which greatly decreases surface charge density. For the two lowest sulfate concentrations, all sulfate is adsorbed with nearly no negatively charged counterion to shield positive charges, thereby limiting the number of positively charged surface groups. The situation changes for the high-sulfate concentration: first, the IEP is decreased due to adsorbed sulfate. Less than 15% sulfate is adsorbed at the lowest pH, leaving a substantial negative charge to allow the increase in protonated surface groups as pH decreases. Calcium sulfate (**Figure 5(d)**) behaves similarly to sodium sulfate; according to the model, calcium adsorption does not occur at low pH. The formation of the CaSO4 0 ion pair in solution affects ionic strength and free sulfate activity; hence the charge is higher in the sodium sulfate system. The model does not predict enhanced calcium uptake and indicates that no system charge reversal occurred at the pH values investigated experimentally. Even so, a possible relative increase in neutral aluminol groups (≡AlOH0 ) at pH ~ 7 occurred in our experiments to favor the sulfate-calcium coadsorption on gibbsite from the 0.5 mM CaSO4 solution.

#### **4.3 Considerations from kaolinite titration data**

Riese [50] titrated kaolinite in CaCl2, NaNO3, and Na2SO4 solutions, and we show the resulting curves in **Figure 6**. The results for NaNO3 (**Figure 6(a)**) showed a typical trend with increasing salt concentration, i.e., a higher titrable charge in a higher electrolyte solution. The absence of spread as a function of salt level observed at low Na2SO4 concentrations (**Figure 6(b)**) could indicate sulfate adsorption when related to the calculations for gibbsite (**Figure 5(c)**). However, for the previous gibbsite charge simulations, the total charge was assessed, whereas for kaolinite the titrable one was determined. The titration data in CaCl2 showed a very small spread and surprisingly low values for the titrable charge as compared to NaNO3 (**Figure 6(c)**) perhaps because the interaction of calcium with the kaolinite surface hindered proton release. Such an interaction would not apply to the kaolinite edge; also, oxidic surface functional groups on edge faces are not expected to bind to calcium at low pH values. Thus, the interaction is more likely with basal planes. Kumar et al. [21] found similar features on the kaolinite gibbsite face and reported visible co-adsorption of chloride at 100 mM CaCl2 concentration. The influence of calcium adsorption on basal surfaces on proton release must be studied to understand the systems. Clearly, the available data suggest that both sulfate and

**53**

**Figure 5.**

*and (d) CaSO4.*

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

calcium each interact with kaolinite at pH 4 and 6, rendering cooperative interactions possible; this trend supports experimental data from our CaSO4 systems.

*suspensions as a function of pH, salt concentration, and composition: (a) NaNO3, (b) Ca(NO3)2, (c) Na2SO4,* 

 *aqueous* 

*Net surface charge (σo) values calculated using the double layer model for standard gibbsite in 10-g L<sup>−</sup><sup>1</sup>*

**Figure 7** shows results of second harmonic generation (SHG) experiments at pH 6 with the c-cut of sapphire. The setup has been described in much detail before [28, 29]. As pointed out before, this crystal plane is more or less structurally equivalent to the gibbsite basal plane, which is present on both gibbsite and kaolinite. Numerous investigations have shown that this plane has a low isoelectric point.

**4.4 Second harmonic generation experiments at room temperature**

*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 5.**

*Advanced Sorption Process Applications*

insignificant.

(≡AlOH0

(this is also true for the subsequent sections).

DLM only includes inner-sphere surface complexes and the model does not distinguish crystal planes. Also when the data base was generated, the information about the Ca adsorption at low pH had not been available. Furthermore, in the following the comparisons between solutions containing 1:1, 2:1 and 2:2 electrolyte solutions are difficult to directly compare, since at the same time the ionic strength and the cation concentrations might be changing for a given concentration of the electrolyte

The enhancement of an inert electrolyte concentration increases the gibbsite charge at a fixed pH with no shift in the IEP (**Figure 5(a)**). The diffuse layer model in the database used for the calculations causes a steep charge increase due to the absence of a Stern layer. This behavior is hidden when plotting titration raw data (i.e., total acid and base added versus pH). As depicted in **Figure 5(b)**, calcium does not significantly affect the curves relative to **Figure 5(a)**. The charge for a given cation concentration and pH increases because the negative counterion concentration is higher in the calcium presence, allowing for better shielding and increased charge. The enhanced charge is not due to calcium adsorption, which is

**Figure 5(c)** shows two effects relative to the gibbsite surface charge as a function of sodium sulfate concentration. One is sulfate uptake, which greatly decreases surface charge density. For the two lowest sulfate concentrations, all sulfate is adsorbed with nearly no negatively charged counterion to shield positive charges, thereby limiting the number of positively charged surface groups. The situation changes for the high-sulfate concentration: first, the IEP is decreased due to adsorbed sulfate. Less than 15% sulfate is adsorbed at the lowest pH, leaving a substantial negative charge to allow the increase in protonated surface groups as pH decreases. Calcium sulfate (**Figure 5(d)**) behaves similarly to sodium sulfate; according to the model,

0

ion pair

calcium adsorption does not occur at low pH. The formation of the CaSO4

adsorption on gibbsite from the 0.5 mM CaSO4 solution.

**4.3 Considerations from kaolinite titration data**

in solution affects ionic strength and free sulfate activity; hence the charge is higher in the sodium sulfate system. The model does not predict enhanced calcium uptake and indicates that no system charge reversal occurred at the pH values investigated experimentally. Even so, a possible relative increase in neutral aluminol groups

Riese [50] titrated kaolinite in CaCl2, NaNO3, and Na2SO4 solutions, and we show the resulting curves in **Figure 6**. The results for NaNO3 (**Figure 6(a)**) showed a typical trend with increasing salt concentration, i.e., a higher titrable charge in a higher electrolyte solution. The absence of spread as a function of salt level observed at low Na2SO4 concentrations (**Figure 6(b)**) could indicate sulfate adsorption when related to the calculations for gibbsite (**Figure 5(c)**). However, for the previous gibbsite charge simulations, the total charge was assessed, whereas for kaolinite the titrable one was determined. The titration data in CaCl2 showed a very small spread and surprisingly low values for the titrable charge as compared to NaNO3 (**Figure 6(c)**) perhaps because the interaction of calcium with the kaolinite surface hindered proton release. Such an interaction would not apply to the kaolinite edge; also, oxidic surface functional groups on edge faces are not expected to bind to calcium at low pH values. Thus, the interaction is more likely with basal planes. Kumar et al. [21] found similar features on the kaolinite gibbsite face and reported visible co-adsorption of chloride at 100 mM CaCl2 concentration. The influence of calcium adsorption on basal surfaces on proton release must be studied to understand the systems. Clearly, the available data suggest that both sulfate and

) at pH ~ 7 occurred in our experiments to favor the sulfate-calcium co-

**52**

*Net surface charge (σo) values calculated using the double layer model for standard gibbsite in 10-g L<sup>−</sup><sup>1</sup> aqueous suspensions as a function of pH, salt concentration, and composition: (a) NaNO3, (b) Ca(NO3)2, (c) Na2SO4, and (d) CaSO4.*

calcium each interact with kaolinite at pH 4 and 6, rendering cooperative interactions possible; this trend supports experimental data from our CaSO4 systems.

#### **4.4 Second harmonic generation experiments at room temperature**

**Figure 7** shows results of second harmonic generation (SHG) experiments at pH 6 with the c-cut of sapphire. The setup has been described in much detail before [28, 29]. As pointed out before, this crystal plane is more or less structurally equivalent to the gibbsite basal plane, which is present on both gibbsite and kaolinite. Numerous investigations have shown that this plane has a low isoelectric point.

#### **Figure 6.**

*Titrable surface charge (Δσ) values obtained by Riese [50] for kaolinite KGa-1 suspended in solutions of (a) NaNO3, (b) Na2SO4, and (c) CaCl2.*

So it is probable that at pH 6, the surface is negatively charged. **Figure 7(a)** shows the results for Na2SO4 addition that correspond to those shown in **Figure 4(a)**. The experiment is started from a water solution to which Na2SO4 is added, and the signal is referenced to the signal in water (this is also the case for all the other data in **Figure 7**). The increase in the SHG signal indicates less positive or more negative charge. Assuming that the sapphire surface is negative at this pH value, it would mean more negative charge and correspond to data on kaolinite rather than those on gibbsite. Also the change in signal is quite small, which also is the case for kaolinite in **Figure 4(a)**. The unexpected adsorption of anions on the negatively charged basal plane of sapphire has also been reported by others. Starting from a solution containing about 1.5 mM sulfate, the addition of CaCl2 results in a much

**55**

**Figure 7.**

*(a) Na2SO4, (b) CaCl2 + 1.5 mM Na2SO4, and (c) CaSO4.*

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

stronger change in the signal (**Figure 7(b)**). This is the same as what was observed in **Figure 4(b)** for both gibbsite and kaolinite. The SHG results on the basal plane clearly mimic the electrokinetic behavior. Furthermore, addition of CaCl2 to the bare surface shows the same behavior irrespective of the absence and presence of sulfate as can be seen in **Figure 7(b)**. This would suggest that the divalent cation is strongly shielding the negative charge, whereas it is not clear whether charge reversal occurs. The SHG data cannot indicate charge reversal. The electrokinetic data

*Second harmonic generation electric field of sapphire basal planes measured under increasing concentrations 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 7.**

*Advanced Sorption Process Applications*

**54**

**Figure 6.**

*(a) NaNO3, (b) Na2SO4, and (c) CaCl2.*

So it is probable that at pH 6, the surface is negatively charged. **Figure 7(a)** shows the results for Na2SO4 addition that correspond to those shown in **Figure 4(a)**. The experiment is started from a water solution to which Na2SO4 is added, and the signal is referenced to the signal in water (this is also the case for all the other data in **Figure 7**). The increase in the SHG signal indicates less positive or more negative charge. Assuming that the sapphire surface is negative at this pH value, it would mean more negative charge and correspond to data on kaolinite rather than those on gibbsite. Also the change in signal is quite small, which also is the case for kaolinite in **Figure 4(a)**. The unexpected adsorption of anions on the negatively charged basal plane of sapphire has also been reported by others. Starting from a solution containing about 1.5 mM sulfate, the addition of CaCl2 results in a much

*Titrable surface charge (Δσ) values obtained by Riese [50] for kaolinite KGa-1 suspended in solutions of* 

*Second harmonic generation electric field of sapphire basal planes measured under increasing concentrations of (a) Na2SO4, (b) CaCl2 + 1.5 mM Na2SO4, and (c) CaSO4.*

stronger change in the signal (**Figure 7(b)**). This is the same as what was observed in **Figure 4(b)** for both gibbsite and kaolinite. The SHG results on the basal plane clearly mimic the electrokinetic behavior. Furthermore, addition of CaCl2 to the bare surface shows the same behavior irrespective of the absence and presence of sulfate as can be seen in **Figure 7(b)**. This would suggest that the divalent cation is strongly shielding the negative charge, whereas it is not clear whether charge reversal occurs. The SHG data cannot indicate charge reversal. The electrokinetic data

for gibbsite might suggest charge reversal at higher Ca concentrations, which in turn would agree with the observation of Siretanu et al. [20] who directly observed the specific interaction of Ca with gibbsite basal planes. As for **Figure 7(a)**, the results for the basal plane from **Figure 7(b)** concur with the data for the particles from **Figure 4(b)**, which could suggest the relevance of the basal planes in the reactions of the particles. Finally in **Figure 7(c)**, CaSO4 solutions have been added to the sapphire basal plane to mimic the results from **Figure 4(c)**. Similar to the discussion with respect to **Figures 4(a)** and **7(a)**, it appears that the SHG is more comparable to the kaolinite system, since the decrease of the signal is rather showing either an increase in positive charge or a decrease in negative charge. The latter happens in **Figure 4(a)** for kaolinite, but none of the options would be applicable for gibbsite. Also the strong change at the lower concentration and the quick plateauing occurs for the kaolinite (**Figure 4(c)**) and for the sapphire basal plane (**Figure 7(c)**). While this cannot exclude a role played by sulfate that is very apparent for gibbsite (**Figure 4(c)**), the predominant role of calcium is obvious. So the role of the calcium sulfate ion pair remains elusive also based on these results, while the small but discernible increase in the signal in the CaSO4 system indicates sulfate interaction at concentrations above 5 mM (**Figure 7(c)**).

To further investigate the role of these ion pairs, temperature-dependent experiments were carried out. The results will be presented and discussed in the final section.

#### **4.5 Second harmonic generation experiments at variable temperature**

Preliminary speciation calculations using various data bases were performed to gain knowledge about the extent of ion-pair formation when the temperature is changed. Calculations were done for 15 mM solutions of divalent ions, starting from 20°C. **Figure 8(a)** shows that with decreasing temperature the extent of ion-pair formation decreases (for CaCl2) or remains approximately constant (for Na2SO4). For both systems the dominant ion pairs make up less than 10%. This is different for the CaSO4 system, where (as could be expected) a decrease in the extent of ion-pair formation occurs with decreasing temperature. Clearly also the ion pair makes up more than 30% as discussed in previous sections. The concomitant solutions were put into contact with sapphire basal planes and studied as a function of temperature by second harmonic generation. The setup used in this case has been extensively described [28, 29]. **Figure 8(b)** shows results of second harmonic generation (SHG) experiments, starting with 15 mM solutions of the divalent ions at pH 6 and room temperature with decreasing temperature. In water the SHG signal is changing little. This is similar for the sodium sulfate system, which is also similar to water. In agreement with the results shown in **Figure 7(a)**, where only small changes were observed in that system. Actually data not shown for about 15 mM sulfate showed on average 1.06 times the water signal. With the average water signal at 0.535 at 20°C, we obtain a value of 0.567 in the presence of 15 mM sulfate, which is within the range of signals collected in the temperature dependence curve. The black curve also suggests a slight decrease of the signal for the sulfate signal compared to water at 4°C, but even here the scatter remains too large to draw ultimate conclusions. With decreasing temperature the sulfate adsorption should increase, but with respect to **Figure 7**, it was concluded that Ca was the more important ion in this system. The more interesting systems in this temperature range in the SHG experiments with the c-cut of sapphire are clearly also the Ca systems. The quite strong decrease in the signal for both Ca systems is observed. The stronger extent in CaCl2 solutions at 15 mM is also observed (giving 0.6 of the water signal, data now shown). Most interestingly, a trend with temperature occurs in the chloride system

**57**

**Figure 8.**

been explored to the extent possible.

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

to lower signal with lower temperature, while the CaSO4 system remains constant. If a decrease in the signal is associated with stronger positive charge, it would suggest that there is a role of sulfate in the CaSO4 system, which may be triggered by the decreased ion-pair formation, making more sulfate available with decreasing temperature. Note that overall much more free Ca is available in the CaCl2 system, even if it is slightly decreasing with decreasing temperature. So decreasing temperature would favor at the same time enhanced sulfate interaction (anions adsorb more strongly on oxides with decreasing temperature) and make more free ions available, which would also favor uptake. Based on the results, the effects of temperature on sulfate would then outcompete those on calcium, but the role of the ion-pair formation could become visible. More experiments with different concentrations would be helpful in gaining more insight. Additional experiments with more soluble MgSO4 are also planned. The formation of the ion pairs can also be an issue in freezing processes on surfaces in the presence of ions like Ca and sulfate, which can be expected in aerosols due to their presence in sea water. This could be an interesting link between the aqueous chemistry and the atmospheric chemistry which has not

*Comparison of Na2SO4 (black) CaCl2 (green) CaSO4 (yellow) solutions (15 mM of divalent ions at 20°C) between 4 and 20°C with respect to speciation (a) and second harmonic signal (b). (b) also includes two runs* 

*with water only, one before and one after the experiments with the salt solutions.*

*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 8.**

*Advanced Sorption Process Applications*

for gibbsite might suggest charge reversal at higher Ca concentrations, which in turn would agree with the observation of Siretanu et al. [20] who directly observed

the specific interaction of Ca with gibbsite basal planes. As for **Figure 7(a)**, the results for the basal plane from **Figure 7(b)** concur with the data for the particles from **Figure 4(b)**, which could suggest the relevance of the basal planes in the reactions of the particles. Finally in **Figure 7(c)**, CaSO4 solutions have been added to the sapphire basal plane to mimic the results from **Figure 4(c)**. Similar to the discussion with respect to **Figures 4(a)** and **7(a)**, it appears that the SHG is more comparable to the kaolinite system, since the decrease of the signal is rather showing either an increase in positive charge or a decrease in negative charge. The latter happens in **Figure 4(a)** for kaolinite, but none of the options would be applicable for gibbsite. Also the strong change at the lower concentration and the quick plateauing occurs for the kaolinite (**Figure 4(c)**) and for the sapphire basal plane (**Figure 7(c)**). While this cannot exclude a role played by sulfate that is very apparent for gibbsite (**Figure 4(c)**), the predominant role of calcium is obvious. So the role of the calcium sulfate ion pair remains elusive also based on these results, while the small but discernible increase in the signal in the CaSO4 system indicates

sulfate interaction at concentrations above 5 mM (**Figure 7(c)**).

To further investigate the role of these ion pairs, temperature-dependent experi-

ments were carried out. The results will be presented and discussed in the final

Preliminary speciation calculations using various data bases were performed to gain knowledge about the extent of ion-pair formation when the temperature is changed. Calculations were done for 15 mM solutions of divalent ions, starting from 20°C. **Figure 8(a)** shows that with decreasing temperature the extent of ion-pair formation decreases (for CaCl2) or remains approximately constant (for Na2SO4). For both systems the dominant ion pairs make up less than 10%. This is different for the CaSO4 system, where (as could be expected) a decrease in the extent of ion-pair formation occurs with decreasing temperature. Clearly also the ion pair makes up more than 30% as discussed in previous sections. The concomitant solutions were put into contact with sapphire basal planes and studied as a function of temperature by second harmonic generation. The setup used in this case has been extensively described [28, 29]. **Figure 8(b)** shows results of second harmonic generation (SHG) experiments, starting with 15 mM solutions of the divalent ions at pH 6 and room temperature with decreasing temperature. In water the SHG signal is changing little. This is similar for the sodium sulfate system, which is also similar to water. In agreement with the results shown in **Figure 7(a)**, where only small changes were observed in that system. Actually data not shown for about 15 mM sulfate showed on average 1.06 times the water signal. With the average water signal at 0.535 at 20°C, we obtain a value of 0.567 in the presence of 15 mM sulfate, which is within the range of signals collected in the temperature dependence curve. The black curve also suggests a slight decrease of the signal for the sulfate signal compared to water at 4°C, but even here the scatter remains too large to draw ultimate conclusions. With decreasing temperature the sulfate adsorption should increase, but with respect to **Figure 7**, it was concluded that Ca was the more important ion in this system. The more interesting systems in this temperature range in the SHG experiments with the c-cut of sapphire are clearly also the Ca systems. The quite strong decrease in the signal for both Ca systems is observed. The stronger extent in CaCl2 solutions at 15 mM is also observed (giving 0.6 of the water signal, data now shown). Most interestingly, a trend with temperature occurs in the chloride system

**4.5 Second harmonic generation experiments at variable temperature**

**56**

section.

*Comparison of Na2SO4 (black) CaCl2 (green) CaSO4 (yellow) solutions (15 mM of divalent ions at 20°C) between 4 and 20°C with respect to speciation (a) and second harmonic signal (b). (b) also includes two runs with water only, one before and one after the experiments with the salt solutions.*

to lower signal with lower temperature, while the CaSO4 system remains constant. If a decrease in the signal is associated with stronger positive charge, it would suggest that there is a role of sulfate in the CaSO4 system, which may be triggered by the decreased ion-pair formation, making more sulfate available with decreasing temperature. Note that overall much more free Ca is available in the CaCl2 system, even if it is slightly decreasing with decreasing temperature. So decreasing temperature would favor at the same time enhanced sulfate interaction (anions adsorb more strongly on oxides with decreasing temperature) and make more free ions available, which would also favor uptake. Based on the results, the effects of temperature on sulfate would then outcompete those on calcium, but the role of the ion-pair formation could become visible. More experiments with different concentrations would be helpful in gaining more insight. Additional experiments with more soluble MgSO4 are also planned. The formation of the ion pairs can also be an issue in freezing processes on surfaces in the presence of ions like Ca and sulfate, which can be expected in aerosols due to their presence in sea water. This could be an interesting link between the aqueous chemistry and the atmospheric chemistry which has not been explored to the extent possible.
