**5.1 Effect of magnesium and sulphate ions**

It was shown that both sulphate and magnesium ions inhibit CaCO3 crystallization [50, 55, 56], even though some discrepancies on the effect of sulphate ions were revealed. For example, at fixed temperature and ionic strength, Karoui et al. [50] showed that, the addition of sulphate ions to the CaCO3 solution increased the induction time and decreased both the crystal growth rate and the amount of the precipitated CaCO3. Comparable results were found by Vavouraki et al. [57]. However, Tlili et al. [64] showed that sulphate ions increase the growth rate of CaCO3 precipitated electrochemically. Many works [51, 65, 66] have shown that sulphate ions inhibit vaterite transformation and enhance the calcite formation for low concentrations. At high concentrations, it was shown that sulphate ions

#### **Figure 4.**

*A tank containing iron-rich water. Iron ions can be present in the aqueous phase either from the corrosion of metallic parts or they may be present in groundwater at concentration levels depending on the mineral composition of the aquifer.*

*Effect of Operating Parameters and Foreign Ions on the Crystal Growth of Calcium Carbonate… DOI: http://dx.doi.org/10.5772/intechopen.94121*

(i) increased the nucleation time [50], (ii) decreased the precipitation rate of calcite formation [57] and (iii) promoted the formation of aragonite via a dissolutionprecipitation process after long reaction times [55, 59].

The temperature and the ratio of magnesium to calcium control the precipitation kinetics, the type of polymorph and the morphology of the CaCO3 precipitates [37, 64, 67]. For example, it was shown that when the Mg2+/Ca2+ ratio passes from 0 to 4, with a Ca2+concentration of 4 <sup>10</sup><sup>3</sup> mol L<sup>1</sup> , the scaling time increases from 50 to 400 min, respectively [64]. Mejri et al. [37] studied the effect of magnesium ions on the CaCO3 precipitation, by the CO2 repelling method, at different temperatures between 30°C and 60°C and with different magnesium to calcium molar ratios (R) in the range 2–5. They showed that, at a fixed temperature, the increase of Mg2+ concentration significantly increased the induction time (tn) and decreased both the initial crystal growth rate (Vi) and the amount precipitated (Rp). For example, at 40°C, tn passed from 7 to 30 min, Vi passed from 0.57 to 0.49 M min<sup>1</sup> and Rp decreased from 81 to 56% when R passed from 0 to 5, respectively. The increase in temperature from 30–60°C weakens the impact of magnesium ions on the retardation of CaCO3 precipitation. These results agree with previous works [38, 67–69], showing that the induction time significantly increased due to the presence of Mg2+. At high concentrations of magnesium, amorphous CaCO3 exhibited a prolonged stability, while it transformed instantly to calcite and vaterite in pure water [58]. In the presence of high amounts of Mg2+, CaCO3 exclusively precipitated as aragonite while calcite and vaterite formed when no Mg2+ ions are present [70]. At low temperature of 30°C, the addition of Mg2+ ions favored the crystallization in the bulk solution of secondary aragonite, resulted from the transformation of vaterite nuclei, while primary aragonite was favored, at high temperature of 60°C, with larger particles sizes than those obtained at 30°C [37].

#### **5.2 Effect of iron ions**

The results reported in literature on the effect of iron ions on CaCO3 precipitation are often contradictory. For example, Kelland [71] found that iron ions concentrations below 25 ppm did not affect the CaCO3 precipitation in high pressure dynamic tube blocking tests. Comparable results were found by Lorenzo et al. [72] who studied the effect of ferrous iron ions on the CaCO3 growth by the constant composition method. However, Herzog et al. [73] found that, in magnetic water treatment devices, an excess of ferrous ions (5.6 ppm) strongly inhibits both calcite growth and the transformation of aragonite to calcite. The inhibiting effectiveness depends strongly on the solution supersaturation [16, 17, 74–76]. Indeed, at high supersaturations, the inhibition effectiveness of iron on CaCO3 crystallization is small. In that case, the inhibition effectiveness can be improved by lowering the solution pH, increasing the iron concentration and/or by lowering the solution supersaturation [16, 17]. At relatively low supersaturations, the addition of iron ions retards the CaCO3 precipitation. Past the onset of nucleation, iron ions enhance the growth rate of CaCO3 and most of CaCO3 amounts precipitate in the bulk solution instead of on the cell walls minimizing, therefore, the risks of scaling. Scaling is defined as the precipitation on the cell walls [16]. These results agree with those of Katz et al. [74] and Takasaki et al. [76] showing that lower supersaturations require less iron for the same extent of CaCO3 growth inhibition.

It was shown [16] that for both temperatures 28°C and 50°C, the addition of iron retarded the nucleation of CaCO3 and, most important, it decreased the CaCO3 precipitation in the cell walls which can be easily collected, reducing therefore scale phenomena in water treatment devices. At a fixed temperature of 50°C, the increase in the iron concentration in the solution to 0.5 mg/FeL did not significantly affect

the CaCO3 precipitation [17]. Comparable result was obtained by Macadam and Parsons [77] who showed that addition of 0.5 mg L<sup>1</sup> of iron did not affect the induction time of CaCO3 precipitation by magnetic stirring at 40°C. For higher iron concentrations, the solution supersaturation decreased, and the inhibition effectiveness of iron ions became more pronounced [17]. At 50°C, the addition of iron ions in the solution promoted aragonite formation rather than vaterite. Indeed, with the increase in the iron ions concentration, the solution supersaturation decreased and the dissolution of vaterite became, most likely, faster which promotes aragonite formation. However, at 28°C, the addition of iron ions did not affect the nature of the precipitated phases vaterite and aragonite, and vaterite with rough surfaces was the predominant polymorph. This was explained that iron ions substituted partially Ca2+, which inhibits vaterite dissolution and prevents the transformation of vaterite into aragonite [16, 17].

For a given iron concentration, increasing the initial pH causes the decrease in the induction time and the increase in the initial crystal growth rate. However, increasing iron ions concentration at a fixed pH in the range 7–9, results in the decrease in both the induction time and the crystal growth rate of CaCO3 [17]. Additionally, the increase in iron ions concentration at a fixed pH reduced the amount of the precipitates obtained and, however, enhanced the precipitation of CaCO3 in the bulk solution, reducing therefore the risks of scaling. Therefore, the CaCO3 precipitation on cell walls, which causes the major phenomenon of scaling, could be lowered by enhancing the initial solution pH or by controlling the concentration of iron ions in the solution. However, the increase of the initial solution pH leads to high amounts of CaCO3 precipitates and even when formed in the bulk solution, they can agglomerate and block up different parts of the water treatment devices such as pipes and conducts, resulting in the blockage of fluid cooling. This increases processing costs, decreases equipment life, and decreases the product water recovery.

At 28°C and initial pH 8, calcite and vaterite were obtained. The increase in iron concentration in the solution resulted in an increased formation of calcite. At initial pH 9 and iron concentration of 4 mg/FeL, calcite was the only phase detected. At 28°C, the increase of the initial solution pH resulted in the disappearance of aragonite and the precipitation of vaterite and calcite. Iron ions addition up to 4 mg/FeL accelerated the transformation of vaterite into calcite and only calcite with stepped surfaces (**Figure 5**) was obtained at initial pH 9 [17]. The stepped surfaces were explained by iron ions adsorption or to difficulties in incorporating building units into the surfaces of calcite crystals [78].

#### **5.3 Effect of chloride ions**

For high concentrations, chloride might affect the stability of calcium in the solution, and thus the precipitation of CaCO3. However, at relatively low temperatures between 15°C and 85°C, the equilibrium association constants of calcium ions with chloride ions in aqueous solution reported in literature are small. In fact, log*K* varies between 0.28 and 0.04 [79–83] which indicates, therefore, the instability of the ion pairs. Korchef and Touabi [17] showed that, at 28 and 50°C and for low concentrations of chloride below 0.18 mM, calcium chloride ion pairs either do not form or have stability fields too small to remarkably decrease the free calcium ion concentration in the solution and to reduce, therefore, the CaCO3 growth rate. Calcium chloride ion pairs form at higher temperatures, i.e., for temperatures in the range 100-360°C, the stability field for CaCl+ decreases with enhancing temperature, whereas that for *CaCl*<sup>0</sup> <sup>2</sup> increases significantly [84]. It should be noted that higher order calcium chloride ion pairs do not form or have small stability fields.

*Effect of Operating Parameters and Foreign Ions on the Crystal Growth of Calcium Carbonate… DOI: http://dx.doi.org/10.5772/intechopen.94121*

#### **Figure 5.**

*Calcite with stepped surfaces obtained at 28°C, initial solution pH 9, and iron ions concentration of 4 mg/FeL.*

In addition, chloride ions can be added as sodium chloride NaCl. Tai and Chen [45] studied the effect of ionic strength and additives concentration on the polymorphism of CaCO3 formed in a constant composition environment. They adjusted the ionic force by NaCl addition which is, they found, inactive in the polymorphic form. Moreover, Takita et al. [85] showed that low concentrations of NaCl in the range 0.1–0.5 M did not remarkably affect the solubility of CaCO3. However, when an excess of NaCl (2.5 M) was added, the solubility of CaCO3 increased which promoted the vaterite to calcite transformation. The resulting high concentration of CaCO3 favored crystal growth rather than nucleation.

#### **5.4 Effect of additives and dissolved organic matter**

Ukrainczyk et al. [86] investigated the effect of salicylic acid derivatives on calcite precipitation kinetics and morphology. Their results showed that the adsorption of additive molecules lowered the CaCO3 growth rate by blocking the propagation of growth sites. This causes the formation of steps and jagged and discontinuous surfaces. Comparable results were obtained for magnesiumcontaining solutions, where calcite crystal edges were rough and growth steps were apparent [55]. Comparable effects were also found when iron was added [16, 17]. The presence of additives, such as monoethylene glycol (MEG), prevents the vaterite to calcite transformation and stabilizes the formation of the relatively unstable vaterite polymorph. For example, Natsi et al. [87] showed that, in the presence of low concentrations of MEG of (1020% v/v), vaterite was stabilized. At stable supersaturated solutions, seeded with quartz and calcite crystals, the growth of CaCO3 was significantly reduced with the increase of MEG concentrations. At 60% of MEG, the vaterite/aragonite- to-calcite transformation was prevented, whatever the synthesis temperature [88].

On the other hand, it was shown that dissolved organic matter (DOM) inhibits CaCO3 crystal growth [89]. The inhibition effect was related to the formation of calcium complexes and the occupation of the CaCO3 growth sites which resulted in CaCO3 rough surfaces. Dzacula et al. [90] studied in vitro precipitation of aragonite in artificial seawater at a high supersaturation of 11, and at a low supersaturation of 5.8. In either chemical systems, different concentrations of soluble organic matter (SOM), extracted from the symbiotic coral B. europea (SOM-Beu) and the asymbiotic one L. pruvoti (SOMLpr) were added in order to investigate their effect on crystal growth or nucleation processes. They showed that, at high supersaturation, the SOMs increased the induction time but did not change the growth rate, and they were incorporated within nanoparticles aggregates. At low supersaturation, the SOMs affected the overgrowing crystalline unit aggregation and did not substantially affect the growth rate.

#### **6. Possible mechanisms**

According to the Ostwald [91] law, the least stable phase, that has the highest solubility, precipitates first and subsequently transforms to the more stable one. This transformation can be either a solid-state transition that invokes internal rearrangements of atoms, ions or molecules [92, 93] or a solution mediated transformation that invokes the dissolution of the metastable phase in the solution and simultaneous precipitation of the stable phase [94, 95]. It was shown that the amorphous calcium carbonate is a precursor in spontaneous precipitation of CaCO3 at relatively high supersaturations. The amorphous phase, which consists of spherical particles with the diameter in the range 50–400 nm, dissolves rapidly in an undersaturated solution and it undergoes a rapid transformation to one of the more stable anhydrous forms (calcite, vaterite and aragonite) [4]. Therefore, the unstable amorphous calcium carbonate is firstly formed [36]. Then, it transforms into crystalline phases which are vaterite, calcite and aragonite. The vaterite can be transformed gradually into calcite or aragonite following two steps that are (i) the vaterite dissolution and (ii) the calcite or aragonite growth [71]. The second step controlled the overall rate of vaterite transformation to the most stable phase calcite or aragonite [96]. The two hydrated crystalline forms of calcium carbonate, calcium carbonate hexahydrate CaCO36H2O and calcium carbonate monohydrate CaCO3H2O, are more stable than the amorphous CaCO3. Calcium carbonate monohydrate crystallized in well-defined spherical crystals with diameters in the range 15–30 mm and calcium carbonate hexahydrate crystallizes in well-defined rhombohedral crystals in the size range between 10 and 40 mm. Generally, hydrated forms precipitate from supersaturated solutions before the more stable anhydrous forms calcite, vaterite and aragonite [4].

As shown above, both sulphate and magnesium ions inhibit CaCO3 crystallization [50, 55, 56]. The inhibition effect of magnesium ions was explained by (i) the increase in calcite solubility induced by magnesium ions insertion into the calcite crystals by substituting calcium ions [97]. This substitution depends on the Mg2+, Ca2+ and *SO*<sup>2</sup> <sup>4</sup> ions concentrations [50], (ii) accumulated strain provoked by smaller magnesium ions incorporation [50, 56] and (iii) adsorption of magnesium ions on calcite surfaces [56]. Indeed, based on the calculation of defect-forming energies in the crystalline lattice of aragonite, Mandakis et al. [98] showed that Mg2+ ions can substitute Ca2+ ions. This substitution may take place in the volume of the orthorhombic lattice of aragonite or on the crystal surfaces. Assessing the mechanism by which Mg2+ ions act to inhibit the crystalline growth of CaCO3 requires an understanding of the changes that must occur when an ion is transferred from the solution to its location in the crystalline lattice. In the solution, Mg2+ ions are surrounded by six molecules of H2O [99]. The incorporation of Mg2+ ions implies therefore the replacement of the H2O molecules by six carbonate groups of the CaCO3 lattice. This is done by partial dehydration (replacement of three H2O molecules by oxygen from

#### *Effect of Operating Parameters and Foreign Ions on the Crystal Growth of Calcium Carbonate… DOI: http://dx.doi.org/10.5772/intechopen.94121*

the calcite lattice) and then diffusion to a growth site and total dehydration (replacement of the three remaining H2O molecules) [99]. According to Folk [100], the adsorption and incorporation of Mg2+ ions into the calcite lattice leads to the distortion of the crystalline lattice because of the lower ion radius of Mg2+ than that of Ca2+ (**Figure 6**). Thus distorted, the crystalline structure of the calcite can no longer accept the following Ca2+, which comes to seek its position in the CaCO3 lattice. Therefore, the crystalline growth of calcite is inhibited. This results in a solubility of calcite containing Mg2+ ions greater than that of pure calcite (without Mg2+) [101].

Sulphate ions can be incorporated into the CaCO3 lattice structure. Kitano et al. [97] demonstrated that this incorporation is easier in calcite than in aragonite and that with the increase in NaCl concentration in the solution, the sulphate ions content in calcite decreases significantly while in aragonite it remains constant. A simple model that describes the incorporation of sulphate into the calcite lattice was proposed by Kontrec et al. [102]. In this model, a carbonate group can be replaced by sulphate. Three oxygen atoms of the tetrahedral sulphate structure, which occupies an area of 256.4 Å<sup>2</sup> , are accommodated in place of the planar carbonate group occupying a smaller surface area of 212.1 Å2 . Sulfur and the fourth oxygen of the sulphate group are outside the carbonate plane. As a result, this substitution causes the crystalline lattice of calcite to be distorted. The distortion on the c-axis is more important than on the other two axes (a and b). A comparable model was proposed for the incorporation of sulphate ions into the orthorhombic structure of aragonite [50, 51]. In addition, sulphate ions can be adsorbed on the CaCO3 surfaces [50]. This adsorption was explained by a mechanism involving a change in the surface charge of CaCO3 and a contribution of the Ca2 ions to the release of carbonate ions from the surface of CaCO3. Indeed, Strand et al. [103] showed that the negatively charged carbonate ions are released from the surface of CaCO3 (desorption) which leads to a positively charged surface. It should be noted that the causes of this desorption have not been mentioned but it was shown that, using a measure of zeta potential, sulphate and calcium ions generate the properties of the CaCO3/water interface. The sulphate ions are adsorbed on this surface and the released carbonate ions react with the Ca2+ ions to form CaCO3 on this same surface (**Figure 7**). This adsorption increases with temperature and/or Ca2+ ion concentration [103].

The inhibitory effectiveness of iron ions on CaCO3 nucleation, and growth was generally attributed to that (i) iron ions are blocking the growth sites of CaCO3 by incorporating the growing layers, which leads to a strained lattice where calcium ions no longer fit [72] (ii) the formation of colloidal iron oxides, i.e., Fe2O3 [104] (iii) iron precipitates mainly in the form of siderite, which acts as heterogeneous sites for CaCO3 precipitation and improves the scale inhibition effectiveness.

The formation of iron oxides and their interactions with CaCO3 remain controversial. For example, experiments were conducted by Herzog et al. [73] on various forms of ferric hydroxide to test if a form is effective for the heterogeneous nucleation of CaCO3. They found that the tested forms of iron hydroxides or

**Figure 6.** *Calcite lattice distortion due to the incorporation of Mg2+.*

**Figure 7.** *Schematic illustration of the adsorption of sulphate ions on the surfaces of CaCO3.*

hydrated iron oxides did not have any effect on CaCO3 nucleation and growth. This disagrees with the results of Mejri et al. [105]. In fact, they studied the CaCO3 precipitation in 2 mM Ca(HCO3)2 solution in the presence of iron ions concentrations between 5 and 20 mg L<sup>1</sup> . They showed that iron ions addition promotes CaCO3 precipitation. This was attributed to the formation of iron hydroxide before the onset of CaCO3 precipitation which plays a role of seed and initiates the CaCO3 nucleation. Coetzee et al*.* [104] showed that the addition of iron had a small effect on the crystal morphology but reduced the induction times. The enhanced nucleation rates were explained by the formation of hematite Fe2O3 colloids that act as seed crystals and increase the heterogeneous nucleation process. What is more, it was shown that the presence of Fe2+ ions in contact with calcite surfaces accelerates calcite dissolution sites during its oxidation and co-precipitation onto calcite surfaces as ferric iron hydroxide, which was rapidly transformed into FeOOH nanoparticles [61]. Thus, the water being treated can be softened by iron precipitation and the final product can be reused as mineral filler powders or as a pigment for industrial applications.

The formation of siderite FeCO3 on the CaCO3 surfaces blocks growth sites and/ or forms a protection layer. This layer retards the CaCO3 precipitation and decreased the corrosion rate of mild steel in simulated saline aquifer environments [106]. The formation of siderite, before the onset CaCO3 precipitation, was experimentally observed by scanning electron microscopy [16]. FeCO3 was formed at relatively low supersaurations and serves as a seed for heteronucleation of vaterite. At high supersaturations, however, FeCO3 either does not form or has stability fields too small to remarkably affect the CaCO3 precipitation. In fact, it was shown that CaCO3 crystallizes in higher amounts than FeCO3 in a solution supersaturated with respect to both FeCO3 and CaCO3 [107], and FeCO3 precipitation is not probable, at high supersaturations, because calcium ions increase the solubility of FeCO3 decreasing, therefore, its precipitation rate [107, 108]. It was shown that, when the carbonate system presents relatively high concentrations, iron solubility is controlled by FeCO3 rather than by Fe(OH)2 in CO2 saturated solutions [109]. This was confirmed by Korchef's work [16]. Indeed, according to Korchef's mechanism [16], FeCO3 forms first from amorphous iron hydroxide, i.e., consistent with the socalled nonclassical nucleation theory [110] and then serves as a template for heteronucleation of CaCO3. When the concentration of the iron ions increases, the amount of FeCO3 formed increases, which increases the number of heterogeneous sites and enhances the inhibition effectiveness of scale. Then, FeCO3 could be incorporated in a solid solution FexCa1xCO3 (0 ≤ x ≤ 1) by a solid–solid transition (**Figure 8**). Indeed, the solids CaCO3 and FeCO3 are isostructural, and their constituent cations (Ca2+ and Fe2+) can coexist, as a mixed metal carbonate, in the

## *Effect of Operating Parameters and Foreign Ions on the Crystal Growth of Calcium Carbonate… DOI: http://dx.doi.org/10.5772/intechopen.94121*

substitutional solid solution FexCa1xCO3 [108]. The FeCO3 dissolution was recently reported by Rizzo et al. [111] who studied the effect of CaCO3 precipitation on the corrosion of carbon steel, covered by a protective FeCO3 layer. The authors showed that CaCO3 precipitation lead to an undersaturated solution with respect to FeCO3 followed by dissolution of the protective FeCO3 layer. The partial substitution of Ca2+ by Fe2+ suppresses the transformation of vaterite into aragonite which results in the decrease of the aragonite amount with increasing iron ions

#### **Figure 8.**

*Korchef's mechanism [16]: In iron-rich water, siderite (FeCO3) form first from amorphous iron hydroxide, consistent with the so-called nonclassical nucleation theory and then serves as a template for heteronucleation of CaCO3. Next, siderite is incorporated in FexCa1 xCO3 (0* ≤ *x* ≤ *1) solid solution by a solid–solid transition.*

concentration. Comparable results were found for Fe(II), [75] Fe(III), [74, 76] and Zn(II) [112]. In fact, Zn(II) precipitated mainly as ZnCO3 which, like siderite, acts as heterogeneous sites in the bulk of solution and decreased the number of crystals on the surface. As a result, scale is reduced [112].

It is worth noticing that, since FeCO3 precipitates at lower pH values than CaCO3 [16], maintaining the solution pH at relatively low values (<7) allowed obtaining only FeCO3. FeCO3 can form protective film on different steel parts of industrial processes which, reduces or completely prevents their corrosion. This reduces remarkably the treatment costs and increases the equipment life.
