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water interface
Journal of Colloid and Interface Science 313 (2007) 184–193 www.elsevier.com/locate/jcis
The calcite/water interface I. Surface charge in indifferent electrolyte media and the influence of low-molecular-weight polyelectrolyte Rasmus Eriksson a , Juha Merta b , Jarl B. Rosenholm a,∗ a Department of Physical Chemistry, Åbo Akademi University, Porthansgatan 3-5, FIN-20500 Åbo, Finland b Oy Keskuslaboratorio, Centrallaboratorium Ab, Tekniikantie 2, P.O. Box 70, FIN-02151 Espoo, Finland
Received 19 February 2007; accepted 13 April 2007 Available online 31 May 2007
Abstract Suspensions of calcium carbonate in water with an indifferent background electrolyte (NaCl) have been investigated using several techniques. Particular attention was paid to the dissolution of calcite at equilibrium and as a function of sodium polyacrylate (NaPA) concentration. Also of interest was how this affects the magnitude of the surface charge and the zeta potential. The development of the interfacial charge is discussed with respect to the dissolved species and with regard to the kinetics of dissolution. The partial pressure of CO2 in solution is believed to play a major role in determining the sign of the charge at equilibrium. In addition to effectively stabilizing calcite suspensions, NaPA was also found to act as a chelating agent at the calcite surface, enhancing the dissolution. The order of addition of NaPA to the suspensions was found to be important. © 2007 Elsevier Inc. All rights reserved. Keywords: Calcium carbonate; Sodium polyacrylate; Dissolution; Electrostatic stabilization; Surface charge; Zeta potential
1. Introduction Pigment-based coating is a common method for improving the printing properties of paper. Coating colors are typically of very high solids content, and thus usually contain some dispersing agent [1–3] along with the binding chemical(s). Typical pigments used in paper coating are kaolin and calcium carbonate. Traditionally, rheology has been an important field of study in the paper coating industry [1,4], but recently, more and more emphasis has been put on understanding colloid chemical phenomena in aqueous suspensions of pigments [5]. Calcium carbonate is one of the most commonly used pigments in paper coating, and the properties of aqueous suspensions of calcium carbonate are the focus of this study. Although the properties of calcium carbonate have been studied extensively, the charging mechanism at the surface in aqueous solutions is not yet fully understood. There has been a lot of debate about * Corresponding author. Fax: +358 2 215 4706.
E-mail address: [email protected] (J.B. Rosenholm). 0021-9797/$ – see front matter © 2007 Elsevier Inc. All rights reserved. doi:10.1016/j.jcis.2007.04.034
the potential-determining ions [6–9] and the sign of the surface charge [6–13]. Some authors have found an isoelectric point (IEP) [1,6,8,13,14], and others have found significant changes in the surface charge as a function of pCa [7,8,10,15–17] or pCO2− 3 [9–11,18]. Different origins of the calcite samples are reported as one reason for the discrepancies in surface charge measurements, and another reason reported is different measuring conditions (including specific adsorption of ionic species and partial pressure of CO2 in aqueous solutions) [5]. In particular, the partial pressure of CO2 may play a large role in the surface charge of calcite, which will be discussed in this work. However, the partial pressure of CO2 in solution is not easy to control, since it requires very pure water and measuring conditions completely isolated from the atmosphere. Also, the calcite pigment itself should be free of adsorbed CO2 , and this is not very likely to happen. The focus of the study is on determining the electrokinetic charge and the dissolution of the calcite surface as a function of the concentration of anionic polyelectrolyte in the presence of an indifferent electrolyte (NaCl). The charging mechanism of
R. Eriksson et al. / Journal of Colloid and Interface Science 313 (2007) 184–193
the calcite surface is discussed, and the partial charge of CO2 in solution is proposed to play a dominant role in the charging of the calcite surface. Anionic polyelectrolytes are commonly used as dispersing agents in calcium carbonate suspensions [2], and in this study, low-molecular-weight sodium polyacrylate (NaPA) was used as a dispersing agent. Particular attention is paid to the presence of ionic species in solution and their effect on the interaction between NaPA and the calcite surface. Measurements were made at several different solids contents in order to study how the available surface area affects NaPA adsorption onto the surface and the resulting electrokinetic charge.
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stirring before removal of the particles. The removal of particles was done after two-stage centrifugation. The first step was centrifugation at 2500 rpm for 40 min, after which the liquid phase, possibly still containing some small particles, was centrifuged at 25,000 rpm for 30 min. Due to the small size (Mw of NaPA) of the polyelectrolyte, it was assumed that any unbound polymer remained in the bulk liquid phase despite the high second centrifugation speed. The liquid phases were analyzed for elements using TOC (Shimadzu TOC-5050) and ICP-OES (Perkin–Elmer Optima 5300 DV). 2.4. Adsorption isotherm
2. Experimental 2.1. Materials The calcium carbonate from Omya Ltd. (HC-90) was used as received. Ninety percent (w/w) of the particles were less than 3 µm in size (average particle size 1.1 µm), determined by sedimentation velocity, and the BET surface area was 8.5 m2 /g (measured with Micromeritics ASAP 2010). An ESCA analysis of the surface revealed no major impurities. The sodium polyacrylate (NaPA) was purchased from BASF (Polysalz S) and was received as a 45% solution. The average molecular weight of the NaPA is 4000 g/mol. The NaCl (J.T. Baker, 100.1% purity) was used as received. Deionized water (Millipore A10) of high purity (conductivity of 0.054 µS/cm) was used in the dispersions. 2.2. General experimental procedure Dispersions of calcium carbonate of varying solids content were analyzed. The solids content was varied between 5 and 60 wt% with respect to water. The calcite was always added to an electrolyte solution during mixing. NaCl was used to vary the ionic strength, and the aquamolarity of NaCl was held between 0 and 0.5 M. When NaPA was included in the dispersions, the appropriate amount was added to the electrolyte solution prior to the addition of the calcite particles. The molar concentrations of NaPA at different solids contents and NaPA concentrations are summarized in Table 1. 2.3. Supernatant analyses With a few exceptions, the pH of all the dispersions was adjusted to 8.5 using either 1 M HCl or 1 M NaOH. All dispersions were allowed to equilibrate for an hour under vigorous
The adsorption of NaPA onto calcite was estimated from the equilibrium concentration of NaPA in solution after removal of the particles. The particles were removed from the suspension according to the method described above. The concentration of NaPA in solution was determined spectrophotometrically at 482 nm using a method developed jointly by Rohm & Haas and Hach Company. For the determination a polyacrylic acid determination apparatus kit (Prod. No. 2225700) and a polyacrylic acid reagent (Prod. No. 2225200) from Hach Company were used. The spectrophotometric analysis was done using a Hach DR/2000 spectrophotometer. 2.5. Zeta potential measurements The zeta potential measurements were performed as a function of solids content (5–60 wt%), NaPA concentration (0–2 wt% of calcite content, cf. Table 1), and electrolyte concentration (0–0.5 mol/dm3 ). For pH adjustments either 1 M HCl or 1 M NaOH was used. All zeta potential measurements were made with a Matec Acoustosizer instrument. The suspensions were stirred for about 45 min and the temperature was adjusted to 25 ◦ C before the first experimental point. After each NaPA addition, pH was adjusted to 8.5 ± 0.1, and the suspension was allowed to equilibrate for 15 min before pH was checked again and adjusted if necessary. 2.6. Net proton charge density titrations The net proton charge density (σH ) of the calcite surface was determined in solutions of varying ionic strength with potentiometric titrations. The solids content used in all measurements was 20 wt%; the electrolyte concentration was varied between 0 and 0.1 M NaCl. All titrations were made at room temperature using a 751 GPD Titrino automatic titrator. Solutions of 0.1 M NaOH and 0.1 M HCl were used to adjust pH.
Table 1 The aquamolarity of NaPA in solution at different loads of NaPA (mol/L) CaCO3 (wt%) 5 20 40 60
NaPA (wt% of calcite content) 0.1%
0.3%
0.5%
0.7%
1.0%
2.0%
5.85 × 10−5
1.75 × 10−4
2.92 × 10−4
4.09 × 10−4
5.85 × 10−4
7.41 × 10−4 1.67 × 10−3
2.22 × 10−3 5.00 × 10−3
3.70 × 10−3 8.33 × 10−3
5.19 × 10−3 1.17 × 10−2
7.41 × 10−3 1.67 × 10−2
1.17 × 10−3 5.56 × 10−3 1.48 × 10−2 3.33 × 10−2
2.78 × 10−4
8.33 × 10−4
1.39 × 10−3
1.94 × 10−3
2.78 × 10−3
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The net proton charge density is calculated by subtracting the net consumed amount of protons in the supernatant from the net consumed amount of protons in the suspension, σH =
sup F susp susp sup nH − nOH − nH − nOH , Asp ws
(1)
where F is Faradays constant, Asp is the specific surface area of the solid, ws is the mass of solid, and n is the amount, indicated in mol. The supernatants were acquired at the native pH (in every case between 9.1 and 9.4) using the same centrifugation procedure as described previously. 2.7. The calcite/water system The thermodynamics of calcite dissolution in aqueous solutions is fairly well known [5–7,19] and will be briefly described here. Calcite dissolves in water according to the following reaction: CaCO3 (s) ≡ Ca2+ (aq) + CO2− 3 (aq).
(2)
At equilibrium, the pH and distribution of different ionic species are dependent on the partial pressure of CO2 in solution. In other words, the species distribution can be controlled by modifying the pH and the concentration of CO2 in the solution. The governing reactions are H2 CO3 (aq) ≡ H2 O (l) + CO2 (g),
(3)
HCO− 3
(4)
(aq) + H2 O (l) ≡ H2 CO3 (aq) + OH− (aq),
Fig. 1. (Solid symbols) Total organic carbon in the liquid phase of suspensions of calcium carbonate at varying solids contents at pH 8.5. (Open symbols) Amount of organic carbon in solution per unit area of solid particles in mg/m2 .
The equilibrium constants for reactions (2) and (7) are taken from [24], the equilibrium constant for reaction (3) is taken from [25], the equilibrium constants for reactions (4) and (5) are taken from [26] and the equilibrium constant for reaction (6) is taken from [19]. Under normal conditions, most of the calcium in solution is in the noncomplexated Ca2+ form. In alkaline pH, the concentration of CaHCO+ 3 may become significant, depending on the partial pressure of CO2 in solution. In all cases, CaOH+ may be ignored.
− − CO2− 3 (aq) + H2 O (l) ≡ HCO3 (aq) + OH (aq),
(5)
3. Results
Ca2+ (aq) + H2 O (l) ≡ Ca(OH)+ (aq) + H+ (aq),
(6)
+ Ca2+ (aq) + HCO− 3 (aq) ≡ CaHCO3 (aq).
(7)
Shown in Fig. 1 is the concentration of organic carbon in solution as a function of solids content in the absence of NaPA. It can be seen that the relative amount of organic carbon leached out increases as a function of the solids content. The surprisingly high concentration of organic carbon in solution is probably grinding aid remnants. The concentrations of calcium and inorganic carbon in supernatants extracted at various solids contents in the absence of NaPA are shown in Fig. 2. The supernatants were extracted at pH 8.5 (the native pH of all the prepared dispersions was between 9.1 and 9.4). Practically all of the inorganic carbon is believed to originate from CO2− 3 species dissolved from the calcite surface. An important conclusion in Fig. 1 is that the concentrations of calcium and inorganic carbon can be considered independent of solids content. Shown in Fig. 3 is the zeta potential of calcite as a function of pH. The surface charge of calcite was found to be positive over the pH range studied (7.5–11). A hysteresis was observed in the titration curve when going toward acidic pH. The net proton charge density of calcite as a function of pH and NaCl concentration is shown in Fig. 4. NaCl clearly functions mainly as an indifferent ion in the calcite/water system. The point of zero charge (pzc) is in all cases between pH 8 and 9. The charge curve at 0.001 M NaCl lies slightly below the other charge curves, and this is probably due to the fact that this titration was continued up to pH 11.3, while the others were continued up to pH 11. Since this method only
Ca(OH)2 also forms in aqueous solutions, but only at very high pH, and in such small quantities that it can be ignored. Assuming the activity coefficients of the species to be equal to 1, the corresponding equilibrium constants for reactions (2)–(7) are as follows: K1 =
[Ca2+ ][CO2− 3 ] = 10−8.48 , [CaCO3 ]
(8)
K2 =
PCO2 = 101.47 , [H2 CO3 ]
(9)
K3 =
[H2 CO3 ][OH− ] = 10−7.7 , [HCO− 3]
K4 =
K5 =
K6 =
− [HCO− 3 ][OH ]
[CO2− 3 ] [CaOH+ ][H+ ] [Ca2+ ] [CaHCO+ 3]
[Ca2+ ][HCO− 3]
(10)
= 10−3.7 ,
(11)
= 10−12.9 ,
(12)
= 101.1 .
(13)
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Fig. 2. Concentrations of calcium and inorganic carbon in the liquid phase of calcium carbonate suspensions of varying solids content, pH 8.5.
Fig. 3. Zeta potential of 20 wt% (8.7 vol%) calcium carbonate suspension as a function of pH.
considers H+ and OH− as potential-determining ions, the surface charge values presented in Fig. 4 are not representative of the actual surface charge [20]. The major potential-determining ions for calcite mentioned in most publications are Ca2+ and CO2− 3 and hydrated or hydroxylated species of these ions [5]. Therefore, the net proton charge titrations only give an indication of the concentration of positively charged (Ca2+ ) and negatively charged (CO2− 3 ) groups at the surface. Since the net proton charge density in Fig. 4 is negative over a wide pH range, it means in practice that the OH− consumption at the surface is larger than the H+ consumption. In other words, this is an indication of a Ca2+ -rich positively charged surface, which is confirmed by the zeta potential measurements (Fig. 3). Disregarding the 0.001 M NaCl curve in Fig. 4, a common intersection point (cip) is observed between pH 9.5 and 10.0. In the same curves there is also a slight shift in the pzc toward higher pH as a function of NaCl concentration. A shift in the cip toward the right (toward higher pH) with regard to the pzc
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Fig. 4. Net proton charge density of 20 wt% (8.7 vol%) calcium carbonate suspensions at various NaCl concentrations.
Fig. 5. Adsorption isotherm of NaPA on calcite. Solids content is 40 wt% (20.0 vol%) and pH is 8.5.
indicates specific adsorption of anions [27]. In this case Cl− ions seem to have a weak specificity to the calcite surface. The adsorption isotherm of NaPA onto calcite is shown in Fig. 5. The numbers in parentheses denote the wt% of NaPA with regard to the amount of calcite in the suspensions. Clearly, at low concentrations of NaPA, the isotherm is a typical highaffinity (Langmuir-type) isotherm, but after monolayer formation (at between 0.3 and 0.5% NaPA), there is further adsorption of NaPA onto calcite. With NaPA concentrations only a little higher than monolayer concentration, the adsorption layer is probably reorganized to allow a denser packing of NaPA molecules at the surface, and with higher concentrations of NaPA, multilayer adsorption is expected. From Figs. 6 and 7 it is clear that the NaPA molecules adsorb extensively on the calcite surface and reverse the surface charge to negative. In Fig. 6 the NaPA titration was started from 0% NaPA and in Fig. 7 the calcite was added to a 0.3% NaPA solution, after which the titration was started. The isoelectric point in Fig. 6 is at about 0.06% NaPA regardless of ionic strength.
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Fig. 6. Zeta potential of 40 wt% (20.0 vol%) calcium carbonate suspensions as a function of NaPA concentration at varying ionic strengths of the solution, pH 8.5.
In Fig. 7 are also shown single points measured at 0.6 and 1.6% NaPA. The zeta potential reaches a plateau value irrespective of ionic strength at ∼0.5% NaPA when the titration is started from 0% NaPA. When 0.3% NaPA is added to the dispersion prior to starting the titration (Fig. 7), there is a minimum in zeta potential at ∼0.5% NaPA, after which subsequent NaPA additions lead to a slight decrease in zeta potential. Generally, the influence of NaCl concentration on the magnitude of the charge on the calcite surface due to NaPA is small. In all titrations the zeta potential reaches a value of −40 mV or higher, which is enough to consider the particles to be electrostatically stable. This was confirmed with sedimentation tests [28]. However, when experimental points at the same electrolyte concentration and NaPA amount are compared in Figs. 6 and 7, the zeta potential values are consistently 5–10 mV higher in Fig. 7. No background correction was made in the zeta potential measurements, due to uncertainties in the (ionic) species composition of the supernatant. However, a NaPA/water (no calcite) titration gave very small electroacoustic signals (zeta potential −1.25 mV), and therefore the zeta potential values measured are considered, within instrument limitations, to be accurate. In addition, the time dependency on zeta potential (40 wt% calcite, no background electrolyte) was checked at a constant pH of 8.5 at two different NaPA concentrations (0.6 and 1.6%). The zeta potential did not change for 7 h, after which the experiments were terminated, since this is well beyond the time scale of interest in this work. Addition of NaPA to calcite dispersions leads to an increased dissolution of the calcite surface, which is demonstrated in Figs. 8 and 9. In other words, NaPA acts as a chelating agent at the surface [21], ripping out Ca2+ ions from the surface and binding them, while CO2− 3 ions diffuse to the bulk solution. This is clearly demonstrated in Fig. 8, where the amount of inorganic carbon increases dramatically in the liquid phase, while there is only a slight increase in the concentration of calcium. However, the amount of dissolved inorganic carbon relative to the amount of added NaPA decreases as a function of solids
Fig. 7. Zeta potential of 40 wt% (20.0 vol%) calcium carbonate suspensions as a function of NaPA concentration at varying ionic strengths of the solution. Open symbols indicate titrations where 0.3% NaPA was added to the solution prior to the start of the titration. Solid squares indicate single-point measurements. pH 8.5 in all titrations.
Fig. 8. Concentrations of calcium (open symbols) and inorganic carbon (solid symbols) in the liquid phase of calcium carbonate suspensions as a function of solids content at pH 8.5. (+) The amount of inorganic carbon per NaPA monomer.
content. Fig. 9 shows the total amount of inorganic carbon in solution as a function of added NaPA. Fig. 10 shows the amount of acid consumed in order to keep the pH at 8.5 ± 0.1 as a function of added NaPA. One titration was made with no calcite particles included (the NaPA amounts added simulated a 40 wt% dispersion) in order to verify that the proton consumption was not due to NaPA. The only other species that could bind protons are the CO2− 3 ions dissolved from the calcite surface due to NaPA addition. In fact, the acid consumption almost exactly matches the amount of inorganic carbon in solution (Fig. 11), which leads to two conclusions: (1) most of the inorganic carbon originates from CO2− 3 leached out from the surface, and (2) the CO2− ions react with water to 3 form bicarbonate ions according to reaction (5). An increasing ionic strength leads to a decrease in the overall acid consump-
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Fig. 9. Total inorganic carbon in the liquid phase of 20 wt% calcium carbonate suspensions as a function of NaPA concentration at pH 8.5.
Fig. 10. Consumed amount of hydrochloric acid in 40 wt% calcium carbonate suspensions as a function of NaPA concentration. (Solid square) NaPA-titration with no particles simulating a 40 wt% suspension.
tion (Fig. 10). This is most likely due to a lesser degree of dissociation of the charged groups on the polyelectrolyte at higher ionic strengths. Although an uncharged polymer also adsorbs onto the calcite surface, the ability to bind Ca2+ is diminished. Finally, in Figs. 12 and 13 are shown the effects of solids content on the zeta potential of calcite as a function of NaPA. In Fig. 12 the titration was started from 0% NaPA and in Fig. 13 the titration was started from 0.3% NaPA. It is immediately evident that there are large differences in the zeta potential values at different solids contents. With no or very little NaPA, i.e., when the zeta potential is still positive (Fig. 12), the zeta potential is independent of solids content, but soon after the zeta potential reaches negative values, differences start to appear. The main trend is that the zeta potential at a certain NaPA concentration is higher as a function of solids content. Also, similarly to what was observed in Figs. 6 and 7, the zeta potentials in Fig. 13 are consistently larger than the zeta potential
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Fig. 11. The proton consumption relative to the amount of inorganic carbon in the liquid phase of a 20 wt% calcium carbonate suspension at pH 8.5.
Fig. 12. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration. NaPA titration started from 0% NaPA, pH 8.5.
values in Fig. 12 (at the same solids content and NaPA concentration). 4. Discussion The sign of the surface charge of calcite in aqueous solutions has been a matter of discussion for decades already, and is still ongoing [6,7,9,10,12,13,15]. In this work it was found to be positive in the pH range studied (Fig. 3), and it seems that a slight majority of publications are in agreement with this result [5]. According to species distribution diagrams, the 2− concentration of (HCO− 3 + CO3 ) exceeds the concentration 2+ of Ca in solutions open to atmospheric CO2 (other species concentrations are negligible), which implies that the surface of calcite is more rich in calcium ions than in carbonate ions at equilibrium. The equilibrium pH of calcite dispersions is very much dependent on the partial pressure of CO2 in solu-
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Fig. 13. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration at pH 8.5. NaPA titration started from 0.3% NaPA.
tion [5], and above a pCO2 of about 10−5.9 atm (corresponding to pHeq < ∼9.8), carbonate species are more abundant in solution than calcium species. This implies that the calcite surface is positive in aqueous suspensions unless the water is basically free of CO2 . The problem with many previous publications is the lack of sufficiently detailed descriptions of the experimental procedures, and therefore the CO2 content of the water used is not known. Douglas and Walker [10] pointed out the possibility that the surface charge of calcite is simply determined by the preferential desorption of Ca2+ and CO2− 3 from the surface. They measured a negative zeta potential of calcite. However, in their work they explicitly stated that the suspensions were prepared in CO2 -free water, and thus the negative charge is logical. Others, attaining a negative surface charge, have measured in water containing “no gas phase” [7], and solutions closed to atmospheric CO2 [8,9]. Also, some authors, having measured the zeta potential of calcite in systems “closed to atmospheric CO2 ,” acquired a positive zeta potential [16,22]. However, the equilibrium pH in these measurements was in fact less than 9.8, indicating enough CO2 in solution to make the surface positive. The equilibrium pH in all of the dispersions prepared in this work was slightly over 9, which implies that the partial pressure of CO2 in the dispersions is higher than 10−5.9 atm. Since the dispersions were prepared open to atmospheric CO2 , this is to be expected. The term “equilibrium pH” is rather ambiguous for calcite dispersions, since it is a very dynamic system. Somasundaran and Agar [6] studied pH changes in calcite dispersions as a function of time and came to the following conclusions: (1) upon addition of the calcite particles the change in pH is “immediate” and (2) after the initial change in pH, it may slowly drift (in one direction, maximum around one pH unit) for a period of several weeks. In our work, the change in pH is also “immediate” after addition of the calcite particles, and during 1 h of mixing, which is the maximum mixing time for the dispersions in this work (excluding titrations), the change in pH is insignificant. Therefore, the pH measured at the end of the mixing period is referred to as the equilibrium pH.
Since the partial pressure of CO2 in solution is presumably quite close to the atmospheric CO2 pressure, the dissolution of CO2− 3 from calcite is expected to be larger than the dissolution of Ca2+ from the surface and the zeta potential is expected to be positive. Figs. 1 and 3 confirm this, thus giving strong support to the theory that the partial pressure of CO2 in solution governs the surface charge of calcite at equilibrium. Simple calculations based on the amounts dissolved from the surface (Fig. 1) according to Eqs. (8)–(13) in order to determine a magnitude for the partial pressure of CO2 in solution do not, however, yield a unanimous figure for this. However, a value for the partial pressure slightly below atmospheric pressure (∼10−3.6 atm) fits perfectly with the measured value for inorganic carbon at pH 8.5. The measured value for the calcium in solution is in this case about six times higher than theory would predict. There are, however, some uncertainties involved. For instance, the pH was adjusted by adding a small amount of acid to the solutions. Addition of acid to calcite dispersions always yields an increase in Ca2+ concentration and a decrease in HCO− 3 and concentration. Another factor that probably leads to deCO2− 3 viations from ideality is the organic carbon found in solution, which is most likely grinding aid remnants. Common grinding aids include polyalcohols, carboxylic acids, and hydroxycarboxylic acids, which are negatively charged at high pH. Like NaPA, these could act as chelating agents at the calcite surface and form complexes with Ca2+ , and consequently, the calcium concentration in solution would be higher than predicted by thermodynamics. Assuming a pure calcite surface, the partial pressure of CO2 in solution seems to have a strong influence on the surface charge at equilibrium. Often the surface is not pure, however. Natural calcite samples in particular tend to contain impurities on the surface, which often are organic in nature, imparting a negative charge to the surface [8]. Specifically adsorbed ions, of course, also affect the surface charge. In previous works, much emphasis has been put on the surface charge of calcite as a function of pH (and whether there exists an IEP or not), but this is a rather confusing discussion, since neither H+ nor OH− are directly potential-determining ions (unlike for metal oxides). Rather, a change in pH changes the species distribution diagrams of the calcite–water system [5,6], and thus the concentration of the important potential− determining ions, i.e., Ca2+ and CO2− 3 (or HCO3 ). As can be seen in Fig. 3, the zeta potential of calcite varies very little as a function of pH, and many other authors have come to the same conclusion [8,10,12,15,17,18]. The slight decrease in zeta potential as a function of pH (Fig. 3) could be explained as an − increase in the concentration of CO2− 3 (as opposed to HCO3 ) in the Stern layer, i.e., an increase in the total negative counterion charge, while the surface charge remains positive. Another reason could be increased precipitation of neutral Ca(OH)2 onto the surface, which could also explain the slight hysteresis in the zeta potential curve (Fig. 3). The dissolution of neutral Ca(OH)2 would be a much slower process than the protonation of CO2− 3 , and thus the formation of a slowly dissolving species at the surface could very well lead to hysteresis in the surface charge curve when the experiment is relatively fast. The exact
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Fig. 14. Zeta potential of 40 wt% calcium carbonate suspensions as a function of NaPA concentration. Solid line—titration started from 0% NaPA, dashed line—titration started from 0.3% NaPA, dotted line—single points.
composition of the species distribution at the surface is, however, very difficult to determine, and at present, the influence of H+ and OH− on the surface charge of calcite remains a matter of speculation. NaPA is an efficient stabilizer of calcium carbonate suspensions (Figs. 6 and 7). Even with an electrolyte concentration as high as 0.5 M NaCl, the suspension is clearly electrostatically stabilized. Interesting to note are the differences in zeta potential in the stable region (CNaPA > ∼0.5%) depending on in which order the NaPA was added. At a certain NaPA concentration, a titration (or even a one-time addition) up to this point always yields a smaller zeta potential value than a dispersion prepared at this NaPA concentration where the particles are added to a solution containing the right amount of NaPA. Also, if a titration is started at some low NaPA concentration (0.3%, shown in Fig. 7), the zeta potential values are generally higher than in a titration that was started from 0% NaPA (Fig. 6), but still not as high as single points. Basically, this situation is illustrated in Fig. 14. The most likely explanation for this phenomenon can be found in the dissolution rate of calcite. Although calcium carbonate suspensions can be stabilized with NaPA in solutions of relatively high ionic strength (Figs. 6 and 7), there is nevertheless a small reduction in zeta potential at 0.1 and 0.5 M NaCl compared to lower electrolyte concentrations. The reason for this is a lower degree of dissociation of the carboxylic acid groups of NaPA, which is most evident in the reduced acid consumption at higher electrolyte concentrations (Fig. 10). Since most of the proton consumption is due to the formation of bicarbonate ions in the solution (Fig. 11), it is therefore a good approximation of the amount of CO2− 3 dissolved from the surface due to NaPA. Consequently, it is a measure of the chelating strength of NaPA. A large number of dissociated groups on NaPA means that more calcium ions bind to NaPA and accordingly, more CO2− 3 diffuse to the bulk solution. In other words, the higher the acid consumption, the larger the number of dissociated groups on NaPA that interact with the
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Fig. 15. Acid consumption in 40 wt% calcium carbonate suspensions as a function of NaPA concentration. Open symbols—titration started from 0% NaPA, solid symbols—titration started from 0.3% NaPA.
surface, thus implying stronger adsorption. When the acid consumptions of the titrations in Figs. 6 and 7 are compared, the acid consumption is higher for the titrations where the titration was started at 0.3% NaPA (see Fig. 15) at the same background electrolyte concentration. Since both the acid consumptions and the zeta potentials are higher in the titrations started at 0.3% NaPA, a stronger interaction between NaPA and the calcite surface in these titrations compared to the titrations started from 0% NaPA is indicated. This could be due either to an increased amount of adsorbed NaPA or to an increased number of charged groups on the NaPA chain (or both), but the dominant effect is difficult to determine. Nevertheless, the conclusion is that the electrolyte concentration plays an important role in the adsorption of NaPA on calcite. Since a solution containing calcite particles always has some dissolved Ca2+ and CO2− 3 in the liquid phase, it becomes obvious why the adsorption of NaPA is dependent on the order of addition. In the case where a zeta potential titration is started from 0% NaPA (Fig. 6), neutralization of the NaPA molecules occurs after every addition by Ca2+ ions in solution, before adsorbing to the particles. When calcite particles are added to a solution already containing NaPA, however, the NaPA molecules adsorb onto the particles fast enough to prevent dissolution equilibrium from being reached [23]. This means that the NaPA molecules have a higher density of dissociated (charged) groups when they reach the calcite surface, thus leading to a small, but measurable, increase in zeta potential. Obviously 0.3% NaPA is not enough to reach an optimum amount of NaPA (from a stabilization point of view), since the single point at 0.6% NaPA has a higher zeta potential than any of the other corresponding ones in Fig. 6. According to the supplier, the optimum amount of NaPA to stabilize the calcite used in this work is about 0.55% NaPA. From the adsorption isotherm in Fig. 5, it is evident that this is beyond the first step of high-affinity adsorption, i.e., beyond the point where NaPA starts to appear in the solution. Since the zeta potential continues to increase (up to
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between 0.5 and 0.7% NaPA) after the first step of high-affinity adsorption, the charge density at the surface is further increased after the initial monolayer of NaPA. Restructuring of the adsorbed polyelectrolyte layer is the most probable cause for the increase in zeta potential after monolayer formation. Restructuring could include a more densely packed adsorption layer or some loops/tails pointing out from the particle surface, thus allowing a larger amount of NaPA to be adsorbed per unit area. Usually polyelectrolytes adsorb more or less flat on a surface, due to strong attraction between the charged groups and the surface. However, similarly to chelating complexes, stable loops and/or tails could possibly occur due to local neutralization by adsorbed calcium ions. The differences in zeta potential as a function of solids content (Figs. 12 and 13) are also a consequence of the electrolyte concentration, as discussed above. Since the calcium concentration at equilibrium is independent of solids content (Fig. 1), the amount of calcium in solution relative to the amount of added NaPA is a function of solids content. To achieve a certain concentration of NaPA (with respect to the amount of calcite), the amounts of NaPA added vary greatly as a function of solids content. Naturally the amount needed increases as a function of solids content. The Ca2+ concentration in solution at equilibrium is large enough to significantly neutralize NaPA at low solids contents (i.e., small amounts of NaPA added), while at larger solids contents the effect is not very pronounced. Fig. 16 shows the zeta potential curves from Fig. 13 with the amounts of calcium in solution relative to the total number of NaPA monomers shown at a few chosen experimental points. The underlined numbers are experimentally determined points, and in the other cases the calcium concentration is assumed to be the equilibrium concentration (9 × 10−4 M, taken as the average equilibrium concentration in Fig. 1). It is immediately obvious that there is a relationship between the calcium concentration relative to NaPA and the zeta potential. The general trend is that the zeta potential increases with decreasing calcium:NaPA monomer ratio. However, this seems to be individually specific for each of the solids contents. For example, at a calcium:NaPA monomer ratio of 0.09, a zeta potential of −40 mV is measured at 20 wt% calcite and a zeta potential of about −51 mV is measured at 40 wt% solids content (Fig. 16). However, since the calcium concentration in solution stays relatively constant (close to equilibrium concentration; see Fig. 17) at low NaPA concentrations, it means that initially the calcium:NaPA monomer ratio decreases and at a certain NaPA concentration (dependent on solids content) it levels out or starts to increase. For instance, a back calculation of the amount of NaPA needed to reach a Ca:NaPA monomer ratio of 0.09 (assuming equilibrium concentration of Ca in solution) for a 40 wt% dispersion gives a NaPA concentration of about 0.18%. From Fig. 12, this yields a zeta potential of about −30 mV. However, as was previously stated, titrations always yield a lower value than separately mixed dispersions, and therefore the “real” zeta potential value would probably be close to −40 mV, i.e., exactly the same as for the 20 wt% dispersion at a 0.09 Ca:NaPA monomer ratio.
Fig. 16. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration at pH 8.5. The numbers denote the amount of calcium relative to the amount of NaPA monomers. The underlined numbers are experimentally determined points and for the others the calcium concentration is assumed to be the average equilibrium concentration.
Fig. 17. Calcium concentration in liquid phase of 40 wt% calcium carbonate suspensions as a function of NaPA concentration, pH 8.5. Dashed horizontal line indicates the average equilibrium concentration of calcium taken from Fig. 1.
The NaPA concentration at the turning point is a function of the total available surface area and the background concentration of dissolved Ca2+ species. From the data in Fig. 17, a NaPA amount corresponding to 0.5% NaPA (at 40 wt% solids content) can be assumed to be the amount of NaPA at which the Ca:NaPA monomer ratio levels off or starts to increase again. This also correlates nicely with the maximum in zeta potential at 0.5% NaPA (Fig. 16). If the calcium concentration in solution stays constant up to this NaPA concentration independent of solids content, it can be used as a reference point for calculating the critical NaPA concentration for each solids content. The NaPA amount at 0.5% NaPA (40 wt% solids content) corresponds to roughly one NaPA molecule (4000 g/mol) for every three Ca2+ ions, which seems a realistic saturation concentration for the binding of Ca2+ to NaPA. The same amount
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of NaPA would correspond to 4.9% NaPA at 5 wt% solids content, 1.1% NaPA at 20 wt% solids content, and 0.3% NaPA at 60 wt% solids content. The shapes of the zeta potential curves in Fig. 16 follow fairly well the calculated values for the saturation concentration of NaPA. At 5 wt% solids content, the saturation concentration is not reached, and thus no maximum in zeta potential is observed. At 20 wt% solids content, the saturation concentration is 1.1% NaPA, which is the point where the zeta potential curve levels out, most likely caused by an increase in the Ca2+ concentration in solution. At 60 wt% solids content, the saturation concentration is 0.3% NaPA, which is before the maximum in zeta potential. However, 0.3% NaPA is most likely not enough to form complete coverage of the particle surface, and hence the surface potential still increases after the saturation point. The significant increase in Ca2+ in solution after the saturation point also explains the slight decrease in zeta potential after 0.5% NaPA at 40 and 60 wt% solids content (Fig. 16). At 40 wt% solids content the Ca:NaPA monomer ratio is doubled in going from 0.5% to 2% NaPA, and similar figures are expected for the 60 wt% dispersion. Since the decrease in zeta potential is relatively small, the reason for the decrease is probably mainly electrostatic screening rather than further neutralization of NaPA at the surface. 5. Conclusions The zeta potential of calcite was found to be positive in the pH region investigated (pH 7.5–11). The dissolved amounts of Ca2+ and CO2− 3 from the calcite surface were close to the expected values in an aqueous dispersion open to atmospheric CO2 . Preferential desorption of CO2− 3 from the surface in these systems is the reason why the calcite surface is positive. The species distribution in water is governed by the partial pressure of CO2 in water and is the most important factor determining the surface charge of pure calcite surfaces at equilibrium. Although calcite suspensions can be efficiently stabilized with NaPA, it also acts as a chelating agent at the surface. The optimum amount for stabilizing our calcite suspensions was found to be between 0.5 and 0.6% NaPA. The ability of NaPA to stabilize calcite suspensions is not very dependent on electrolyte concentration (up to at least 0.5 M NaCl), but at higher electrolyte concentrations the chelating effect is diminished, due to a smaller number of dissociated groups on the NaPA molecules. Ca2+ ions (and probably other multivalent cations) bind strongly to NaPA and thus reduce its effectiveness as a dispersing agent. For stabilizing purposes, this means that the order of addition of NaPA is important. NaPA added to a calcite suspension is always partly neutralized by Ca2+ ions in solution before adsorbing on to the particles. If, on the other hand, cal-
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cite particles are added to a solution containing NaPA, there is not enough time to reach dissolution equilibrium of calcite before the NaPA molecules are adsorbed, and thus the density of dissociated groups on NaPA is higher when the calcite surface is reached. The dependency on equilibrium Ca2+ concentration is also the reason that there are such large differences in zeta potentials at different solids contents. The larger the concentration of Ca2+ relative to NaPA, the lower the zeta potential. Acknowledgments The Technology Agency of Finland and KCL are gratefully acknowledged for financial support. References [1] B.W. Greene, A.S. Reder, Tappi J. 57 (5) (May 1974). [2] L. Järnström, Nord. Pulp Pap. Res. J. 1 (1993). [3] K.R. Rogan, A.C. Bentham, I.A. George, D.R. Skuse, Colloid Polym. Sci. 272 (1994) 1175. [4] L. Järnström, G. Ström, P. Stenius, Tappi J. 70 (9) (1987) 101. [5] L. Madsen, in: A.T. Hubbard (Ed.), Surface Charge of Calcite, Encyclopedia of Surface and Colloid Science, Dekker, New York, 2002. [6] P. Somasundaran, G.E. Agar, J. Colloid Interface Sci. 24 (1967) 433. [7] D.W. Thompson, P.G. Pownall, J. Colloid Interface Sci. 131 (1) (1989) 74. [8] D.S. Cicerone, A.E. Regazzoni, M.A. Blesa, J. Colloid Interface Sci. 154 (2) (1992) 423. [9] P. Moulin, H. Roques, J. Colloid Interface Sci. 261 (2003) 115. [10] H.W. Douglas, R.A. Walker, Trans. Faraday Soc. 46 (1950) 559. [11] P.V. Smallwood, Colloid Polym. Sci. 255 (1977) 881. [12] B. Siffert, P. Fimbel, Colloids Surf. 11 (1984) 377. [13] N. Vdovi´c, J. Biš´can, Colloids Surf. A 137 (1998) 7. [14] J.O. Amankonah, P. Somasundaran, Colloids Surf. 15 (1985) 335. [15] T. Foxall, G.C. Peterson, H.M. Rendall, A.L. Smith, J. Chem. Soc. Faraday Trans. 75 (1979) 1034. [16] Y.C. Huang, F.M. Fowkes, T.B. Lloyd, N.D. Sanders, Langmuir 7 (1991) 1742. [17] D.G. Kanellopoulou, P.G. Koutsoukos, Langmuir 19 (2003) 5691. [18] D.W. Fuerstenau, Pradip, R. Herrera-Urbina, Colloids Surf. 68 (1992) 95. [19] P. Somasundaran, J.O. Amankonah, K.P. Ananthapadmabhan, Colloids Surf. 15 (1985) 309. [20] G. Sposito, Environ. Sci. Technol. 32 (19) (1998) 2815. [21] C.N. Fredd, H.S. Fogler, J. Colloid Interface Sci. 204 (1998) 187. [22] T.S. Berlin, A.V. Khabakov, Geochemistry 3 (1961) 217. [23] M. Thompson, S.J. Wilkins, R.G. Compton, H.A. Viles, J. Colloid Interface Sci. 260 (2003) 204. [24] L.N. Plummer, E. Brusenberg, Geochim. Cosmochim. Acta 46 (1982) 1011. [25] W. Stumm, Chemistry of the Solid–Water Interface: Processes at the Mineral–Water and Particle–Water Interface in Natural Systems, Wiley, New York, 1992, pp. 46–59. [26] R.M. Garrels, C.L. Christ, Solutions, Minerals and Equilibria, Harper & Row, New York, 1965, pp. 74–93. [27] J. Lyklema, Pure Appl. Chem. 63 (6) (1991) 895. [28] R. Eriksson et al., unpublished results.
The calcite/water interface I. Surface charge in indifferent electrolyte media and the influence of low-molecular-weight polyelectrolyte Rasmus Eriksson a , Juha Merta b , Jarl B. Rosenholm a,∗ a Department of Physical Chemistry, Åbo Akademi University, Porthansgatan 3-5, FIN-20500 Åbo, Finland b Oy Keskuslaboratorio, Centrallaboratorium Ab, Tekniikantie 2, P.O. Box 70, FIN-02151 Espoo, Finland
Received 19 February 2007; accepted 13 April 2007 Available online 31 May 2007
Abstract Suspensions of calcium carbonate in water with an indifferent background electrolyte (NaCl) have been investigated using several techniques. Particular attention was paid to the dissolution of calcite at equilibrium and as a function of sodium polyacrylate (NaPA) concentration. Also of interest was how this affects the magnitude of the surface charge and the zeta potential. The development of the interfacial charge is discussed with respect to the dissolved species and with regard to the kinetics of dissolution. The partial pressure of CO2 in solution is believed to play a major role in determining the sign of the charge at equilibrium. In addition to effectively stabilizing calcite suspensions, NaPA was also found to act as a chelating agent at the calcite surface, enhancing the dissolution. The order of addition of NaPA to the suspensions was found to be important. © 2007 Elsevier Inc. All rights reserved. Keywords: Calcium carbonate; Sodium polyacrylate; Dissolution; Electrostatic stabilization; Surface charge; Zeta potential
1. Introduction Pigment-based coating is a common method for improving the printing properties of paper. Coating colors are typically of very high solids content, and thus usually contain some dispersing agent [1–3] along with the binding chemical(s). Typical pigments used in paper coating are kaolin and calcium carbonate. Traditionally, rheology has been an important field of study in the paper coating industry [1,4], but recently, more and more emphasis has been put on understanding colloid chemical phenomena in aqueous suspensions of pigments [5]. Calcium carbonate is one of the most commonly used pigments in paper coating, and the properties of aqueous suspensions of calcium carbonate are the focus of this study. Although the properties of calcium carbonate have been studied extensively, the charging mechanism at the surface in aqueous solutions is not yet fully understood. There has been a lot of debate about * Corresponding author. Fax: +358 2 215 4706.
E-mail address: [email protected] (J.B. Rosenholm). 0021-9797/$ – see front matter © 2007 Elsevier Inc. All rights reserved. doi:10.1016/j.jcis.2007.04.034
the potential-determining ions [6–9] and the sign of the surface charge [6–13]. Some authors have found an isoelectric point (IEP) [1,6,8,13,14], and others have found significant changes in the surface charge as a function of pCa [7,8,10,15–17] or pCO2− 3 [9–11,18]. Different origins of the calcite samples are reported as one reason for the discrepancies in surface charge measurements, and another reason reported is different measuring conditions (including specific adsorption of ionic species and partial pressure of CO2 in aqueous solutions) [5]. In particular, the partial pressure of CO2 may play a large role in the surface charge of calcite, which will be discussed in this work. However, the partial pressure of CO2 in solution is not easy to control, since it requires very pure water and measuring conditions completely isolated from the atmosphere. Also, the calcite pigment itself should be free of adsorbed CO2 , and this is not very likely to happen. The focus of the study is on determining the electrokinetic charge and the dissolution of the calcite surface as a function of the concentration of anionic polyelectrolyte in the presence of an indifferent electrolyte (NaCl). The charging mechanism of
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the calcite surface is discussed, and the partial charge of CO2 in solution is proposed to play a dominant role in the charging of the calcite surface. Anionic polyelectrolytes are commonly used as dispersing agents in calcium carbonate suspensions [2], and in this study, low-molecular-weight sodium polyacrylate (NaPA) was used as a dispersing agent. Particular attention is paid to the presence of ionic species in solution and their effect on the interaction between NaPA and the calcite surface. Measurements were made at several different solids contents in order to study how the available surface area affects NaPA adsorption onto the surface and the resulting electrokinetic charge.
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stirring before removal of the particles. The removal of particles was done after two-stage centrifugation. The first step was centrifugation at 2500 rpm for 40 min, after which the liquid phase, possibly still containing some small particles, was centrifuged at 25,000 rpm for 30 min. Due to the small size (Mw of NaPA) of the polyelectrolyte, it was assumed that any unbound polymer remained in the bulk liquid phase despite the high second centrifugation speed. The liquid phases were analyzed for elements using TOC (Shimadzu TOC-5050) and ICP-OES (Perkin–Elmer Optima 5300 DV). 2.4. Adsorption isotherm
2. Experimental 2.1. Materials The calcium carbonate from Omya Ltd. (HC-90) was used as received. Ninety percent (w/w) of the particles were less than 3 µm in size (average particle size 1.1 µm), determined by sedimentation velocity, and the BET surface area was 8.5 m2 /g (measured with Micromeritics ASAP 2010). An ESCA analysis of the surface revealed no major impurities. The sodium polyacrylate (NaPA) was purchased from BASF (Polysalz S) and was received as a 45% solution. The average molecular weight of the NaPA is 4000 g/mol. The NaCl (J.T. Baker, 100.1% purity) was used as received. Deionized water (Millipore A10) of high purity (conductivity of 0.054 µS/cm) was used in the dispersions. 2.2. General experimental procedure Dispersions of calcium carbonate of varying solids content were analyzed. The solids content was varied between 5 and 60 wt% with respect to water. The calcite was always added to an electrolyte solution during mixing. NaCl was used to vary the ionic strength, and the aquamolarity of NaCl was held between 0 and 0.5 M. When NaPA was included in the dispersions, the appropriate amount was added to the electrolyte solution prior to the addition of the calcite particles. The molar concentrations of NaPA at different solids contents and NaPA concentrations are summarized in Table 1. 2.3. Supernatant analyses With a few exceptions, the pH of all the dispersions was adjusted to 8.5 using either 1 M HCl or 1 M NaOH. All dispersions were allowed to equilibrate for an hour under vigorous
The adsorption of NaPA onto calcite was estimated from the equilibrium concentration of NaPA in solution after removal of the particles. The particles were removed from the suspension according to the method described above. The concentration of NaPA in solution was determined spectrophotometrically at 482 nm using a method developed jointly by Rohm & Haas and Hach Company. For the determination a polyacrylic acid determination apparatus kit (Prod. No. 2225700) and a polyacrylic acid reagent (Prod. No. 2225200) from Hach Company were used. The spectrophotometric analysis was done using a Hach DR/2000 spectrophotometer. 2.5. Zeta potential measurements The zeta potential measurements were performed as a function of solids content (5–60 wt%), NaPA concentration (0–2 wt% of calcite content, cf. Table 1), and electrolyte concentration (0–0.5 mol/dm3 ). For pH adjustments either 1 M HCl or 1 M NaOH was used. All zeta potential measurements were made with a Matec Acoustosizer instrument. The suspensions were stirred for about 45 min and the temperature was adjusted to 25 ◦ C before the first experimental point. After each NaPA addition, pH was adjusted to 8.5 ± 0.1, and the suspension was allowed to equilibrate for 15 min before pH was checked again and adjusted if necessary. 2.6. Net proton charge density titrations The net proton charge density (σH ) of the calcite surface was determined in solutions of varying ionic strength with potentiometric titrations. The solids content used in all measurements was 20 wt%; the electrolyte concentration was varied between 0 and 0.1 M NaCl. All titrations were made at room temperature using a 751 GPD Titrino automatic titrator. Solutions of 0.1 M NaOH and 0.1 M HCl were used to adjust pH.
Table 1 The aquamolarity of NaPA in solution at different loads of NaPA (mol/L) CaCO3 (wt%) 5 20 40 60
NaPA (wt% of calcite content) 0.1%
0.3%
0.5%
0.7%
1.0%
2.0%
5.85 × 10−5
1.75 × 10−4
2.92 × 10−4
4.09 × 10−4
5.85 × 10−4
7.41 × 10−4 1.67 × 10−3
2.22 × 10−3 5.00 × 10−3
3.70 × 10−3 8.33 × 10−3
5.19 × 10−3 1.17 × 10−2
7.41 × 10−3 1.67 × 10−2
1.17 × 10−3 5.56 × 10−3 1.48 × 10−2 3.33 × 10−2
2.78 × 10−4
8.33 × 10−4
1.39 × 10−3
1.94 × 10−3
2.78 × 10−3
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The net proton charge density is calculated by subtracting the net consumed amount of protons in the supernatant from the net consumed amount of protons in the suspension, σH =
sup F susp susp sup nH − nOH − nH − nOH , Asp ws
(1)
where F is Faradays constant, Asp is the specific surface area of the solid, ws is the mass of solid, and n is the amount, indicated in mol. The supernatants were acquired at the native pH (in every case between 9.1 and 9.4) using the same centrifugation procedure as described previously. 2.7. The calcite/water system The thermodynamics of calcite dissolution in aqueous solutions is fairly well known [5–7,19] and will be briefly described here. Calcite dissolves in water according to the following reaction: CaCO3 (s) ≡ Ca2+ (aq) + CO2− 3 (aq).
(2)
At equilibrium, the pH and distribution of different ionic species are dependent on the partial pressure of CO2 in solution. In other words, the species distribution can be controlled by modifying the pH and the concentration of CO2 in the solution. The governing reactions are H2 CO3 (aq) ≡ H2 O (l) + CO2 (g),
(3)
HCO− 3
(4)
(aq) + H2 O (l) ≡ H2 CO3 (aq) + OH− (aq),
Fig. 1. (Solid symbols) Total organic carbon in the liquid phase of suspensions of calcium carbonate at varying solids contents at pH 8.5. (Open symbols) Amount of organic carbon in solution per unit area of solid particles in mg/m2 .
The equilibrium constants for reactions (2) and (7) are taken from [24], the equilibrium constant for reaction (3) is taken from [25], the equilibrium constants for reactions (4) and (5) are taken from [26] and the equilibrium constant for reaction (6) is taken from [19]. Under normal conditions, most of the calcium in solution is in the noncomplexated Ca2+ form. In alkaline pH, the concentration of CaHCO+ 3 may become significant, depending on the partial pressure of CO2 in solution. In all cases, CaOH+ may be ignored.
− − CO2− 3 (aq) + H2 O (l) ≡ HCO3 (aq) + OH (aq),
(5)
3. Results
Ca2+ (aq) + H2 O (l) ≡ Ca(OH)+ (aq) + H+ (aq),
(6)
+ Ca2+ (aq) + HCO− 3 (aq) ≡ CaHCO3 (aq).
(7)
Shown in Fig. 1 is the concentration of organic carbon in solution as a function of solids content in the absence of NaPA. It can be seen that the relative amount of organic carbon leached out increases as a function of the solids content. The surprisingly high concentration of organic carbon in solution is probably grinding aid remnants. The concentrations of calcium and inorganic carbon in supernatants extracted at various solids contents in the absence of NaPA are shown in Fig. 2. The supernatants were extracted at pH 8.5 (the native pH of all the prepared dispersions was between 9.1 and 9.4). Practically all of the inorganic carbon is believed to originate from CO2− 3 species dissolved from the calcite surface. An important conclusion in Fig. 1 is that the concentrations of calcium and inorganic carbon can be considered independent of solids content. Shown in Fig. 3 is the zeta potential of calcite as a function of pH. The surface charge of calcite was found to be positive over the pH range studied (7.5–11). A hysteresis was observed in the titration curve when going toward acidic pH. The net proton charge density of calcite as a function of pH and NaCl concentration is shown in Fig. 4. NaCl clearly functions mainly as an indifferent ion in the calcite/water system. The point of zero charge (pzc) is in all cases between pH 8 and 9. The charge curve at 0.001 M NaCl lies slightly below the other charge curves, and this is probably due to the fact that this titration was continued up to pH 11.3, while the others were continued up to pH 11. Since this method only
Ca(OH)2 also forms in aqueous solutions, but only at very high pH, and in such small quantities that it can be ignored. Assuming the activity coefficients of the species to be equal to 1, the corresponding equilibrium constants for reactions (2)–(7) are as follows: K1 =
[Ca2+ ][CO2− 3 ] = 10−8.48 , [CaCO3 ]
(8)
K2 =
PCO2 = 101.47 , [H2 CO3 ]
(9)
K3 =
[H2 CO3 ][OH− ] = 10−7.7 , [HCO− 3]
K4 =
K5 =
K6 =
− [HCO− 3 ][OH ]
[CO2− 3 ] [CaOH+ ][H+ ] [Ca2+ ] [CaHCO+ 3]
[Ca2+ ][HCO− 3]
(10)
= 10−3.7 ,
(11)
= 10−12.9 ,
(12)
= 101.1 .
(13)
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Fig. 2. Concentrations of calcium and inorganic carbon in the liquid phase of calcium carbonate suspensions of varying solids content, pH 8.5.
Fig. 3. Zeta potential of 20 wt% (8.7 vol%) calcium carbonate suspension as a function of pH.
considers H+ and OH− as potential-determining ions, the surface charge values presented in Fig. 4 are not representative of the actual surface charge [20]. The major potential-determining ions for calcite mentioned in most publications are Ca2+ and CO2− 3 and hydrated or hydroxylated species of these ions [5]. Therefore, the net proton charge titrations only give an indication of the concentration of positively charged (Ca2+ ) and negatively charged (CO2− 3 ) groups at the surface. Since the net proton charge density in Fig. 4 is negative over a wide pH range, it means in practice that the OH− consumption at the surface is larger than the H+ consumption. In other words, this is an indication of a Ca2+ -rich positively charged surface, which is confirmed by the zeta potential measurements (Fig. 3). Disregarding the 0.001 M NaCl curve in Fig. 4, a common intersection point (cip) is observed between pH 9.5 and 10.0. In the same curves there is also a slight shift in the pzc toward higher pH as a function of NaCl concentration. A shift in the cip toward the right (toward higher pH) with regard to the pzc
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Fig. 4. Net proton charge density of 20 wt% (8.7 vol%) calcium carbonate suspensions at various NaCl concentrations.
Fig. 5. Adsorption isotherm of NaPA on calcite. Solids content is 40 wt% (20.0 vol%) and pH is 8.5.
indicates specific adsorption of anions [27]. In this case Cl− ions seem to have a weak specificity to the calcite surface. The adsorption isotherm of NaPA onto calcite is shown in Fig. 5. The numbers in parentheses denote the wt% of NaPA with regard to the amount of calcite in the suspensions. Clearly, at low concentrations of NaPA, the isotherm is a typical highaffinity (Langmuir-type) isotherm, but after monolayer formation (at between 0.3 and 0.5% NaPA), there is further adsorption of NaPA onto calcite. With NaPA concentrations only a little higher than monolayer concentration, the adsorption layer is probably reorganized to allow a denser packing of NaPA molecules at the surface, and with higher concentrations of NaPA, multilayer adsorption is expected. From Figs. 6 and 7 it is clear that the NaPA molecules adsorb extensively on the calcite surface and reverse the surface charge to negative. In Fig. 6 the NaPA titration was started from 0% NaPA and in Fig. 7 the calcite was added to a 0.3% NaPA solution, after which the titration was started. The isoelectric point in Fig. 6 is at about 0.06% NaPA regardless of ionic strength.
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Fig. 6. Zeta potential of 40 wt% (20.0 vol%) calcium carbonate suspensions as a function of NaPA concentration at varying ionic strengths of the solution, pH 8.5.
In Fig. 7 are also shown single points measured at 0.6 and 1.6% NaPA. The zeta potential reaches a plateau value irrespective of ionic strength at ∼0.5% NaPA when the titration is started from 0% NaPA. When 0.3% NaPA is added to the dispersion prior to starting the titration (Fig. 7), there is a minimum in zeta potential at ∼0.5% NaPA, after which subsequent NaPA additions lead to a slight decrease in zeta potential. Generally, the influence of NaCl concentration on the magnitude of the charge on the calcite surface due to NaPA is small. In all titrations the zeta potential reaches a value of −40 mV or higher, which is enough to consider the particles to be electrostatically stable. This was confirmed with sedimentation tests [28]. However, when experimental points at the same electrolyte concentration and NaPA amount are compared in Figs. 6 and 7, the zeta potential values are consistently 5–10 mV higher in Fig. 7. No background correction was made in the zeta potential measurements, due to uncertainties in the (ionic) species composition of the supernatant. However, a NaPA/water (no calcite) titration gave very small electroacoustic signals (zeta potential −1.25 mV), and therefore the zeta potential values measured are considered, within instrument limitations, to be accurate. In addition, the time dependency on zeta potential (40 wt% calcite, no background electrolyte) was checked at a constant pH of 8.5 at two different NaPA concentrations (0.6 and 1.6%). The zeta potential did not change for 7 h, after which the experiments were terminated, since this is well beyond the time scale of interest in this work. Addition of NaPA to calcite dispersions leads to an increased dissolution of the calcite surface, which is demonstrated in Figs. 8 and 9. In other words, NaPA acts as a chelating agent at the surface [21], ripping out Ca2+ ions from the surface and binding them, while CO2− 3 ions diffuse to the bulk solution. This is clearly demonstrated in Fig. 8, where the amount of inorganic carbon increases dramatically in the liquid phase, while there is only a slight increase in the concentration of calcium. However, the amount of dissolved inorganic carbon relative to the amount of added NaPA decreases as a function of solids
Fig. 7. Zeta potential of 40 wt% (20.0 vol%) calcium carbonate suspensions as a function of NaPA concentration at varying ionic strengths of the solution. Open symbols indicate titrations where 0.3% NaPA was added to the solution prior to the start of the titration. Solid squares indicate single-point measurements. pH 8.5 in all titrations.
Fig. 8. Concentrations of calcium (open symbols) and inorganic carbon (solid symbols) in the liquid phase of calcium carbonate suspensions as a function of solids content at pH 8.5. (+) The amount of inorganic carbon per NaPA monomer.
content. Fig. 9 shows the total amount of inorganic carbon in solution as a function of added NaPA. Fig. 10 shows the amount of acid consumed in order to keep the pH at 8.5 ± 0.1 as a function of added NaPA. One titration was made with no calcite particles included (the NaPA amounts added simulated a 40 wt% dispersion) in order to verify that the proton consumption was not due to NaPA. The only other species that could bind protons are the CO2− 3 ions dissolved from the calcite surface due to NaPA addition. In fact, the acid consumption almost exactly matches the amount of inorganic carbon in solution (Fig. 11), which leads to two conclusions: (1) most of the inorganic carbon originates from CO2− 3 leached out from the surface, and (2) the CO2− ions react with water to 3 form bicarbonate ions according to reaction (5). An increasing ionic strength leads to a decrease in the overall acid consump-
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Fig. 9. Total inorganic carbon in the liquid phase of 20 wt% calcium carbonate suspensions as a function of NaPA concentration at pH 8.5.
Fig. 10. Consumed amount of hydrochloric acid in 40 wt% calcium carbonate suspensions as a function of NaPA concentration. (Solid square) NaPA-titration with no particles simulating a 40 wt% suspension.
tion (Fig. 10). This is most likely due to a lesser degree of dissociation of the charged groups on the polyelectrolyte at higher ionic strengths. Although an uncharged polymer also adsorbs onto the calcite surface, the ability to bind Ca2+ is diminished. Finally, in Figs. 12 and 13 are shown the effects of solids content on the zeta potential of calcite as a function of NaPA. In Fig. 12 the titration was started from 0% NaPA and in Fig. 13 the titration was started from 0.3% NaPA. It is immediately evident that there are large differences in the zeta potential values at different solids contents. With no or very little NaPA, i.e., when the zeta potential is still positive (Fig. 12), the zeta potential is independent of solids content, but soon after the zeta potential reaches negative values, differences start to appear. The main trend is that the zeta potential at a certain NaPA concentration is higher as a function of solids content. Also, similarly to what was observed in Figs. 6 and 7, the zeta potentials in Fig. 13 are consistently larger than the zeta potential
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Fig. 11. The proton consumption relative to the amount of inorganic carbon in the liquid phase of a 20 wt% calcium carbonate suspension at pH 8.5.
Fig. 12. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration. NaPA titration started from 0% NaPA, pH 8.5.
values in Fig. 12 (at the same solids content and NaPA concentration). 4. Discussion The sign of the surface charge of calcite in aqueous solutions has been a matter of discussion for decades already, and is still ongoing [6,7,9,10,12,13,15]. In this work it was found to be positive in the pH range studied (Fig. 3), and it seems that a slight majority of publications are in agreement with this result [5]. According to species distribution diagrams, the 2− concentration of (HCO− 3 + CO3 ) exceeds the concentration 2+ of Ca in solutions open to atmospheric CO2 (other species concentrations are negligible), which implies that the surface of calcite is more rich in calcium ions than in carbonate ions at equilibrium. The equilibrium pH of calcite dispersions is very much dependent on the partial pressure of CO2 in solu-
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Fig. 13. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration at pH 8.5. NaPA titration started from 0.3% NaPA.
tion [5], and above a pCO2 of about 10−5.9 atm (corresponding to pHeq < ∼9.8), carbonate species are more abundant in solution than calcium species. This implies that the calcite surface is positive in aqueous suspensions unless the water is basically free of CO2 . The problem with many previous publications is the lack of sufficiently detailed descriptions of the experimental procedures, and therefore the CO2 content of the water used is not known. Douglas and Walker [10] pointed out the possibility that the surface charge of calcite is simply determined by the preferential desorption of Ca2+ and CO2− 3 from the surface. They measured a negative zeta potential of calcite. However, in their work they explicitly stated that the suspensions were prepared in CO2 -free water, and thus the negative charge is logical. Others, attaining a negative surface charge, have measured in water containing “no gas phase” [7], and solutions closed to atmospheric CO2 [8,9]. Also, some authors, having measured the zeta potential of calcite in systems “closed to atmospheric CO2 ,” acquired a positive zeta potential [16,22]. However, the equilibrium pH in these measurements was in fact less than 9.8, indicating enough CO2 in solution to make the surface positive. The equilibrium pH in all of the dispersions prepared in this work was slightly over 9, which implies that the partial pressure of CO2 in the dispersions is higher than 10−5.9 atm. Since the dispersions were prepared open to atmospheric CO2 , this is to be expected. The term “equilibrium pH” is rather ambiguous for calcite dispersions, since it is a very dynamic system. Somasundaran and Agar [6] studied pH changes in calcite dispersions as a function of time and came to the following conclusions: (1) upon addition of the calcite particles the change in pH is “immediate” and (2) after the initial change in pH, it may slowly drift (in one direction, maximum around one pH unit) for a period of several weeks. In our work, the change in pH is also “immediate” after addition of the calcite particles, and during 1 h of mixing, which is the maximum mixing time for the dispersions in this work (excluding titrations), the change in pH is insignificant. Therefore, the pH measured at the end of the mixing period is referred to as the equilibrium pH.
Since the partial pressure of CO2 in solution is presumably quite close to the atmospheric CO2 pressure, the dissolution of CO2− 3 from calcite is expected to be larger than the dissolution of Ca2+ from the surface and the zeta potential is expected to be positive. Figs. 1 and 3 confirm this, thus giving strong support to the theory that the partial pressure of CO2 in solution governs the surface charge of calcite at equilibrium. Simple calculations based on the amounts dissolved from the surface (Fig. 1) according to Eqs. (8)–(13) in order to determine a magnitude for the partial pressure of CO2 in solution do not, however, yield a unanimous figure for this. However, a value for the partial pressure slightly below atmospheric pressure (∼10−3.6 atm) fits perfectly with the measured value for inorganic carbon at pH 8.5. The measured value for the calcium in solution is in this case about six times higher than theory would predict. There are, however, some uncertainties involved. For instance, the pH was adjusted by adding a small amount of acid to the solutions. Addition of acid to calcite dispersions always yields an increase in Ca2+ concentration and a decrease in HCO− 3 and concentration. Another factor that probably leads to deCO2− 3 viations from ideality is the organic carbon found in solution, which is most likely grinding aid remnants. Common grinding aids include polyalcohols, carboxylic acids, and hydroxycarboxylic acids, which are negatively charged at high pH. Like NaPA, these could act as chelating agents at the calcite surface and form complexes with Ca2+ , and consequently, the calcium concentration in solution would be higher than predicted by thermodynamics. Assuming a pure calcite surface, the partial pressure of CO2 in solution seems to have a strong influence on the surface charge at equilibrium. Often the surface is not pure, however. Natural calcite samples in particular tend to contain impurities on the surface, which often are organic in nature, imparting a negative charge to the surface [8]. Specifically adsorbed ions, of course, also affect the surface charge. In previous works, much emphasis has been put on the surface charge of calcite as a function of pH (and whether there exists an IEP or not), but this is a rather confusing discussion, since neither H+ nor OH− are directly potential-determining ions (unlike for metal oxides). Rather, a change in pH changes the species distribution diagrams of the calcite–water system [5,6], and thus the concentration of the important potential− determining ions, i.e., Ca2+ and CO2− 3 (or HCO3 ). As can be seen in Fig. 3, the zeta potential of calcite varies very little as a function of pH, and many other authors have come to the same conclusion [8,10,12,15,17,18]. The slight decrease in zeta potential as a function of pH (Fig. 3) could be explained as an − increase in the concentration of CO2− 3 (as opposed to HCO3 ) in the Stern layer, i.e., an increase in the total negative counterion charge, while the surface charge remains positive. Another reason could be increased precipitation of neutral Ca(OH)2 onto the surface, which could also explain the slight hysteresis in the zeta potential curve (Fig. 3). The dissolution of neutral Ca(OH)2 would be a much slower process than the protonation of CO2− 3 , and thus the formation of a slowly dissolving species at the surface could very well lead to hysteresis in the surface charge curve when the experiment is relatively fast. The exact
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Fig. 14. Zeta potential of 40 wt% calcium carbonate suspensions as a function of NaPA concentration. Solid line—titration started from 0% NaPA, dashed line—titration started from 0.3% NaPA, dotted line—single points.
composition of the species distribution at the surface is, however, very difficult to determine, and at present, the influence of H+ and OH− on the surface charge of calcite remains a matter of speculation. NaPA is an efficient stabilizer of calcium carbonate suspensions (Figs. 6 and 7). Even with an electrolyte concentration as high as 0.5 M NaCl, the suspension is clearly electrostatically stabilized. Interesting to note are the differences in zeta potential in the stable region (CNaPA > ∼0.5%) depending on in which order the NaPA was added. At a certain NaPA concentration, a titration (or even a one-time addition) up to this point always yields a smaller zeta potential value than a dispersion prepared at this NaPA concentration where the particles are added to a solution containing the right amount of NaPA. Also, if a titration is started at some low NaPA concentration (0.3%, shown in Fig. 7), the zeta potential values are generally higher than in a titration that was started from 0% NaPA (Fig. 6), but still not as high as single points. Basically, this situation is illustrated in Fig. 14. The most likely explanation for this phenomenon can be found in the dissolution rate of calcite. Although calcium carbonate suspensions can be stabilized with NaPA in solutions of relatively high ionic strength (Figs. 6 and 7), there is nevertheless a small reduction in zeta potential at 0.1 and 0.5 M NaCl compared to lower electrolyte concentrations. The reason for this is a lower degree of dissociation of the carboxylic acid groups of NaPA, which is most evident in the reduced acid consumption at higher electrolyte concentrations (Fig. 10). Since most of the proton consumption is due to the formation of bicarbonate ions in the solution (Fig. 11), it is therefore a good approximation of the amount of CO2− 3 dissolved from the surface due to NaPA. Consequently, it is a measure of the chelating strength of NaPA. A large number of dissociated groups on NaPA means that more calcium ions bind to NaPA and accordingly, more CO2− 3 diffuse to the bulk solution. In other words, the higher the acid consumption, the larger the number of dissociated groups on NaPA that interact with the
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Fig. 15. Acid consumption in 40 wt% calcium carbonate suspensions as a function of NaPA concentration. Open symbols—titration started from 0% NaPA, solid symbols—titration started from 0.3% NaPA.
surface, thus implying stronger adsorption. When the acid consumptions of the titrations in Figs. 6 and 7 are compared, the acid consumption is higher for the titrations where the titration was started at 0.3% NaPA (see Fig. 15) at the same background electrolyte concentration. Since both the acid consumptions and the zeta potentials are higher in the titrations started at 0.3% NaPA, a stronger interaction between NaPA and the calcite surface in these titrations compared to the titrations started from 0% NaPA is indicated. This could be due either to an increased amount of adsorbed NaPA or to an increased number of charged groups on the NaPA chain (or both), but the dominant effect is difficult to determine. Nevertheless, the conclusion is that the electrolyte concentration plays an important role in the adsorption of NaPA on calcite. Since a solution containing calcite particles always has some dissolved Ca2+ and CO2− 3 in the liquid phase, it becomes obvious why the adsorption of NaPA is dependent on the order of addition. In the case where a zeta potential titration is started from 0% NaPA (Fig. 6), neutralization of the NaPA molecules occurs after every addition by Ca2+ ions in solution, before adsorbing to the particles. When calcite particles are added to a solution already containing NaPA, however, the NaPA molecules adsorb onto the particles fast enough to prevent dissolution equilibrium from being reached [23]. This means that the NaPA molecules have a higher density of dissociated (charged) groups when they reach the calcite surface, thus leading to a small, but measurable, increase in zeta potential. Obviously 0.3% NaPA is not enough to reach an optimum amount of NaPA (from a stabilization point of view), since the single point at 0.6% NaPA has a higher zeta potential than any of the other corresponding ones in Fig. 6. According to the supplier, the optimum amount of NaPA to stabilize the calcite used in this work is about 0.55% NaPA. From the adsorption isotherm in Fig. 5, it is evident that this is beyond the first step of high-affinity adsorption, i.e., beyond the point where NaPA starts to appear in the solution. Since the zeta potential continues to increase (up to
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between 0.5 and 0.7% NaPA) after the first step of high-affinity adsorption, the charge density at the surface is further increased after the initial monolayer of NaPA. Restructuring of the adsorbed polyelectrolyte layer is the most probable cause for the increase in zeta potential after monolayer formation. Restructuring could include a more densely packed adsorption layer or some loops/tails pointing out from the particle surface, thus allowing a larger amount of NaPA to be adsorbed per unit area. Usually polyelectrolytes adsorb more or less flat on a surface, due to strong attraction between the charged groups and the surface. However, similarly to chelating complexes, stable loops and/or tails could possibly occur due to local neutralization by adsorbed calcium ions. The differences in zeta potential as a function of solids content (Figs. 12 and 13) are also a consequence of the electrolyte concentration, as discussed above. Since the calcium concentration at equilibrium is independent of solids content (Fig. 1), the amount of calcium in solution relative to the amount of added NaPA is a function of solids content. To achieve a certain concentration of NaPA (with respect to the amount of calcite), the amounts of NaPA added vary greatly as a function of solids content. Naturally the amount needed increases as a function of solids content. The Ca2+ concentration in solution at equilibrium is large enough to significantly neutralize NaPA at low solids contents (i.e., small amounts of NaPA added), while at larger solids contents the effect is not very pronounced. Fig. 16 shows the zeta potential curves from Fig. 13 with the amounts of calcium in solution relative to the total number of NaPA monomers shown at a few chosen experimental points. The underlined numbers are experimentally determined points, and in the other cases the calcium concentration is assumed to be the equilibrium concentration (9 × 10−4 M, taken as the average equilibrium concentration in Fig. 1). It is immediately obvious that there is a relationship between the calcium concentration relative to NaPA and the zeta potential. The general trend is that the zeta potential increases with decreasing calcium:NaPA monomer ratio. However, this seems to be individually specific for each of the solids contents. For example, at a calcium:NaPA monomer ratio of 0.09, a zeta potential of −40 mV is measured at 20 wt% calcite and a zeta potential of about −51 mV is measured at 40 wt% solids content (Fig. 16). However, since the calcium concentration in solution stays relatively constant (close to equilibrium concentration; see Fig. 17) at low NaPA concentrations, it means that initially the calcium:NaPA monomer ratio decreases and at a certain NaPA concentration (dependent on solids content) it levels out or starts to increase. For instance, a back calculation of the amount of NaPA needed to reach a Ca:NaPA monomer ratio of 0.09 (assuming equilibrium concentration of Ca in solution) for a 40 wt% dispersion gives a NaPA concentration of about 0.18%. From Fig. 12, this yields a zeta potential of about −30 mV. However, as was previously stated, titrations always yield a lower value than separately mixed dispersions, and therefore the “real” zeta potential value would probably be close to −40 mV, i.e., exactly the same as for the 20 wt% dispersion at a 0.09 Ca:NaPA monomer ratio.
Fig. 16. Zeta potential of different solids content calcium carbonate suspensions as a function of NaPA concentration at pH 8.5. The numbers denote the amount of calcium relative to the amount of NaPA monomers. The underlined numbers are experimentally determined points and for the others the calcium concentration is assumed to be the average equilibrium concentration.
Fig. 17. Calcium concentration in liquid phase of 40 wt% calcium carbonate suspensions as a function of NaPA concentration, pH 8.5. Dashed horizontal line indicates the average equilibrium concentration of calcium taken from Fig. 1.
The NaPA concentration at the turning point is a function of the total available surface area and the background concentration of dissolved Ca2+ species. From the data in Fig. 17, a NaPA amount corresponding to 0.5% NaPA (at 40 wt% solids content) can be assumed to be the amount of NaPA at which the Ca:NaPA monomer ratio levels off or starts to increase again. This also correlates nicely with the maximum in zeta potential at 0.5% NaPA (Fig. 16). If the calcium concentration in solution stays constant up to this NaPA concentration independent of solids content, it can be used as a reference point for calculating the critical NaPA concentration for each solids content. The NaPA amount at 0.5% NaPA (40 wt% solids content) corresponds to roughly one NaPA molecule (4000 g/mol) for every three Ca2+ ions, which seems a realistic saturation concentration for the binding of Ca2+ to NaPA. The same amount
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of NaPA would correspond to 4.9% NaPA at 5 wt% solids content, 1.1% NaPA at 20 wt% solids content, and 0.3% NaPA at 60 wt% solids content. The shapes of the zeta potential curves in Fig. 16 follow fairly well the calculated values for the saturation concentration of NaPA. At 5 wt% solids content, the saturation concentration is not reached, and thus no maximum in zeta potential is observed. At 20 wt% solids content, the saturation concentration is 1.1% NaPA, which is the point where the zeta potential curve levels out, most likely caused by an increase in the Ca2+ concentration in solution. At 60 wt% solids content, the saturation concentration is 0.3% NaPA, which is before the maximum in zeta potential. However, 0.3% NaPA is most likely not enough to form complete coverage of the particle surface, and hence the surface potential still increases after the saturation point. The significant increase in Ca2+ in solution after the saturation point also explains the slight decrease in zeta potential after 0.5% NaPA at 40 and 60 wt% solids content (Fig. 16). At 40 wt% solids content the Ca:NaPA monomer ratio is doubled in going from 0.5% to 2% NaPA, and similar figures are expected for the 60 wt% dispersion. Since the decrease in zeta potential is relatively small, the reason for the decrease is probably mainly electrostatic screening rather than further neutralization of NaPA at the surface. 5. Conclusions The zeta potential of calcite was found to be positive in the pH region investigated (pH 7.5–11). The dissolved amounts of Ca2+ and CO2− 3 from the calcite surface were close to the expected values in an aqueous dispersion open to atmospheric CO2 . Preferential desorption of CO2− 3 from the surface in these systems is the reason why the calcite surface is positive. The species distribution in water is governed by the partial pressure of CO2 in water and is the most important factor determining the surface charge of pure calcite surfaces at equilibrium. Although calcite suspensions can be efficiently stabilized with NaPA, it also acts as a chelating agent at the surface. The optimum amount for stabilizing our calcite suspensions was found to be between 0.5 and 0.6% NaPA. The ability of NaPA to stabilize calcite suspensions is not very dependent on electrolyte concentration (up to at least 0.5 M NaCl), but at higher electrolyte concentrations the chelating effect is diminished, due to a smaller number of dissociated groups on the NaPA molecules. Ca2+ ions (and probably other multivalent cations) bind strongly to NaPA and thus reduce its effectiveness as a dispersing agent. For stabilizing purposes, this means that the order of addition of NaPA is important. NaPA added to a calcite suspension is always partly neutralized by Ca2+ ions in solution before adsorbing on to the particles. If, on the other hand, cal-
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cite particles are added to a solution containing NaPA, there is not enough time to reach dissolution equilibrium of calcite before the NaPA molecules are adsorbed, and thus the density of dissociated groups on NaPA is higher when the calcite surface is reached. The dependency on equilibrium Ca2+ concentration is also the reason that there are such large differences in zeta potentials at different solids contents. The larger the concentration of Ca2+ relative to NaPA, the lower the zeta potential. Acknowledgments The Technology Agency of Finland and KCL are gratefully acknowledged for financial support. References [1] B.W. Greene, A.S. Reder, Tappi J. 57 (5) (May 1974). [2] L. Järnström, Nord. Pulp Pap. Res. J. 1 (1993). [3] K.R. Rogan, A.C. Bentham, I.A. George, D.R. Skuse, Colloid Polym. Sci. 272 (1994) 1175. [4] L. Järnström, G. Ström, P. Stenius, Tappi J. 70 (9) (1987) 101. [5] L. Madsen, in: A.T. Hubbard (Ed.), Surface Charge of Calcite, Encyclopedia of Surface and Colloid Science, Dekker, New York, 2002. [6] P. Somasundaran, G.E. Agar, J. Colloid Interface Sci. 24 (1967) 433. [7] D.W. Thompson, P.G. Pownall, J. Colloid Interface Sci. 131 (1) (1989) 74. [8] D.S. Cicerone, A.E. Regazzoni, M.A. Blesa, J. Colloid Interface Sci. 154 (2) (1992) 423. [9] P. Moulin, H. Roques, J. Colloid Interface Sci. 261 (2003) 115. [10] H.W. Douglas, R.A. Walker, Trans. Faraday Soc. 46 (1950) 559. [11] P.V. Smallwood, Colloid Polym. Sci. 255 (1977) 881. [12] B. Siffert, P. Fimbel, Colloids Surf. 11 (1984) 377. [13] N. Vdovi´c, J. Biš´can, Colloids Surf. A 137 (1998) 7. [14] J.O. Amankonah, P. Somasundaran, Colloids Surf. 15 (1985) 335. [15] T. Foxall, G.C. Peterson, H.M. Rendall, A.L. Smith, J. Chem. Soc. Faraday Trans. 75 (1979) 1034. [16] Y.C. Huang, F.M. Fowkes, T.B. Lloyd, N.D. Sanders, Langmuir 7 (1991) 1742. [17] D.G. Kanellopoulou, P.G. Koutsoukos, Langmuir 19 (2003) 5691. [18] D.W. Fuerstenau, Pradip, R. Herrera-Urbina, Colloids Surf. 68 (1992) 95. [19] P. Somasundaran, J.O. Amankonah, K.P. Ananthapadmabhan, Colloids Surf. 15 (1985) 309. [20] G. Sposito, Environ. Sci. Technol. 32 (19) (1998) 2815. [21] C.N. Fredd, H.S. Fogler, J. Colloid Interface Sci. 204 (1998) 187. [22] T.S. Berlin, A.V. Khabakov, Geochemistry 3 (1961) 217. [23] M. Thompson, S.J. Wilkins, R.G. Compton, H.A. Viles, J. Colloid Interface Sci. 260 (2003) 204. [24] L.N. Plummer, E. Brusenberg, Geochim. Cosmochim. Acta 46 (1982) 1011. [25] W. Stumm, Chemistry of the Solid–Water Interface: Processes at the Mineral–Water and Particle–Water Interface in Natural Systems, Wiley, New York, 1992, pp. 46–59. [26] R.M. Garrels, C.L. Christ, Solutions, Minerals and Equilibria, Harper & Row, New York, 1965, pp. 74–93. [27] J. Lyklema, Pure Appl. Chem. 63 (6) (1991) 895. [28] R. Eriksson et al., unpublished results.