Langmuir 1991, 7, 1742-1748
1742
Adsorption of Calcium Ions from Calcium Chloride Solutions onto Calcium Carbonate Particles? Yan C. Huang,$ Frederick M. Fowkes,ll Thomas B. Lloyd,$and Nigel D. Sanders**% Department of Chemistry, Lehigh University, Bethlehem, Pennsylvania 18015, and Specialty Minerals Group, Pfizer, Inc, Bethlehem, Pennsylvania 1801 7 Received November 26, 1990 The adsorption of calcium ions on the surface of precipitated calcium carbonate (calcite, prismatic, 0.7 pm diameter) has been studied in 20 w t 94 aqueous slurries by calcium chloride addition. The adsorption of the calcium ions was found to follow the Langmuir adsorption isotherm, and the adsorbed calcium ions were the {-potential-determiningions. The monolayer coverage was determined to be 1.58 X lV mol of Ca2+/m2 (105A2/Ca2+). Experimentally, this coverage was closely approached at the Ca2+concentration of 7.51 mM in the bulk solution after adsorption equilibrium and was about one-fifth of the lattice Ca2+ density. The cpotential, as determined in the high-solidsslurries by a Matec ESA-8000 instrument, varied with increasing Ca2+ concentration from +3.8 to +18.7 mV. As the {potential increased and the interparticle repulsion became greater, the sedimentation volume decreased, and the dispersion was slower to flocculate, a typical response to improved dispersion. The standard-state Gibbs free energy change for the adsorption process was calculated to be -28.3 kJ/mol (-6.8 kcal/mol) by using the experimentally determined adsorption equilibrium constant. The heat of adsorption measured by direct solution microcalorimetry was -6.9 kcal/mol of Ca2+, indicating that the adsorption is enthalpically driven. Introduction
Calcium carbonate is a slightly soluble oxysalt widely distributed in nature as limestone and marble and exhibiting a large variety of internal crystal structures and external crystal habits.' Natural and synthetic forms of calcium carbonate have achieved a position of prominence as industrial mineral products that are of utility not only as bulk fillers substituting for more expensive materials but as functional components of composite materials. Many of these functional applications require the calcium carbonate product to have controlled surface properties. Thus, the inherent surface and colloid chemistry of CaC03 is of great practical as well as fundamental interest. Surface chemical studies of calcium carbonates published to date have, however, occasionally yielded contradictory and inconclusive results, partly because of the heterogeneous nature of CaCOs surfaces which are produced by grinding of minerals2J and by low pH precipitation process. These studies have been directed primarily at understanding the colloidal properties of small calcite particles in aqueous suspension: properties important for industrial mineral separation (flotation) process control and for application of CaC03 in water-based coating formulations. Early work41semphasized the determination of surface charge by electrokinetic measurements as a function of Author to whom correspondence should be addressed. Presented at the 64th Colloid and Surface Science Symposium, sponsored by the Division of Colloid and Surface Chemistry of the American Chemical Society,Lehigh University, Bethlehem, PA, June 18-20,1990. t Lehigh University. 1 Deceased. Pfizer, Inc. (1) Palache, C.; Berman,H.; Frondel,C. Dana's Syetem of Mineralogy, 7th ed.;John Wiley and Son: New York, 1961;Vol. 11, p 141,p 182. (2)Goujon, G.;Mutaftschiev, B. J. Colloid Interface Sci. 1976,57, 148. (3)Gnmmage,R.B.;Gregg,9. J. J. ColloidhterfaceSci. 1972,38,118. (4)Somasundaran, P.; Agar, G. E. J. Colloid Interface Sci. 1967,24, 433. (5) Fuerstenau, M. C.; Gutierrez, G.; Elgillani, D. A. Trans. AIME 1968,241,319. t
0743-7463/91/2407-1742$02.50/0
pH and reported isoelectric points (IEP) over a wide range (8< IEP < 11). A paper by Siffert and Fimbels concluded that the electrokinetic properties of aqueous calcite suspensions are dependent on a number of factors including the solid concentration of the suspension. The first strong experimental evidence for a mechanism for the development of surface charge on calcite was presented by Foxall et ala7who found by microelectrophoresis that the Ca2+/COs2-pair formed the fundamental potential determining ion system for CaCOs. This discovery has been confirmed recently by Thompson and Pownall using the streaming potential technique.s The purpose of the present paper is to describe studies of the adsorption of Ca2+ion from aqueous solutions on the surface of a very well characterized prismatic calcite powder of syn*etic origin that demonstrate for the first time the adsorption thermodynamics which determines the surface potential of CaCOs in aqueous suspension.
Experimental Techniques Materials. Precipitated calcium carbonate(prismaticCalcite, 99+ % CaCOs, Albaglos M) was obtained from Pfizer, Inc, and used without further purification. This material consistsof 0.7pm particles with a specificsurface area (BET Nd of 8.63 m*/g, about 20% of which is present in 20-and 27-Amicropores (Codter Omnisorp). Surface analyses by ESCA and ion scattering spectroscopy (ISS) show the surfaces to be free of major contaminants with the largest being a trace amount (1.8atom %) of magnesium. ISSsputtering depth profiie spectra indicate that the Mg is confined to a depth of 10-20 A. Doppler electrophoretic light-scattering measurements at low solids (0.17%; Coulter DELSA 440)yield a {potential of +18.3 f 0.8 mV at pH = 8 and pCa = 3.1 in 0.01 M KCl, which is in good agreement with the resulta of Foxall et al. (+21 mV).' Preparation of High Solid Content Dirperrionr. The calcium carbonatedispersionsusedin the study of the adsorption of calcium ion ontothe CaCOg pigments and ale0 used to measure [potential by the Matec instrument were prepared in 8-02 (237mL)wide-mouthclear-glassround bottleswithTeflon linedecrewcaps. Thirty grams of the calcium carbonate was weighed into (6) Siffert, B.,Fimbel, P. Colloids Surf. 1984, 377. (7)Foxall, T.;Peterson, G.C.; Rendall, H. M.; Smith,A. L. J. Chem. SOC.,Faraday Trans. 1979,1034. (8) Thompson, D. W.; Pownall, P. G. J. Colloid Interface Sci. 1909, 131, 74.
0 1991 American Chemical Society
Langmuir, Vol. 7, No.8, 1991 1743
Adsorption of Cas+ onto CaCOS Particles each of the bottles containing 120mL of calcium chloride (reagent grade CaClr2H~0,Aeaar) solutions or just plain distilled deionwater. To disperse the contents in each of the bottles, ized (DDI) they were vigorously sonicated 3 times, 20 minutes each time, using a 600-W ultrasonic probe. Between sonication treatments there was a period of 30 min to let the contents in the bottle equilibrate. The bottles were sealed tightly and then agitated at room temperature (23 "C) in a laboratory rolling mill for 5 days. Calcium carbonate dispersions used in sedimentation studies were prepared in heavy-duty 13-mLgraduated conicalcentrifuge tubes with flathead ground glass stoppers. For each dispersion, 2 g of the calcium carbonate was weighed into a 25-mL beaker containing 6 mL of a calcium chloride solution. The contents in the beaker were agitated in an ultrasonic bath at room temperature for 15 min and then transferred to a centrifuge tube. The beaker was rinsed out with 2 mL of the same CaCl2 solution. The 2 mL rinsing-off was also transferred into the centrifuge tube. These centrifuge tubes were shaken in a wrisbaction shaker (Burrell, Model 71,250 W) overnight at room temperature followed by 5 days of agitation in the laboratory rolling mill. t-Potential Determination with Matec. Twenty minutes before the Matec measurements, the contents in each of the 8-02 wide-mouth round bottles were sonicated vigorously for 15 min with the ultrasonic probe. The Matec electroacoustic detector was placed 1 in. deep (about half the total depth) into the dispersions. To ensure maximum electroacoustic signal, the frequency of the electrical field was optimized for each sample by using the Matec phase-calibrating program. Each l-potential reading was averaged from 10 measurements. Four or five replicates were determined for each initial CaC12 Concentration. The average particle size in the dispersions waa taken as 1 Mm, and the density of the calcium carbonate was determined by pycnometry to be 2.4421 f 0.0007 g/mL. Both the particle size and the density are parameters required by the Matec technique to convert the electrokinetic sonic amplitude (ESA) into [ potential. Sedimentation 0 bservations. After removal from the rolling mill, the 13-mL graduated centrifuge tubes were placed on a test-tube rack, and the sedimentation volumes were observed as a function of time. Flocculation Behavior in Dilute Suspensions. When no additional changes in sedimentation volume could be observed, the supernatant solution in each of the centrifuge tubes was taken out along with a tiny amount of the settled calcium carbonate particles from the top layer so as to prepare a dilute suspension. The particle concentrations in the dilute suspensions were adjusted to 600-700 (w/w) ppm so that their photon counts per second were within the range from 5 X 1V to 1 X 108, which is a suitable particle concentration range for the Coulter N4 submicrometer particle analyzer,then time-dependent particle-sizing was performed on these dilute suspensions. Microcalorimetric Titration. Heats of adsorption of Ca2+ on CaCOs were determined by directly measuring the heat of titration of a 1.0 M CaClz solution into a 15 wt ?6 slurry with a Tronac 1260batch-type solution microcalorimeter running in an ieoperibol mode and subtracting the background heat of dilution of the CaClz adsorbate. Adsorption Determination. The studies of the adsorption of the calcium ions from the calcium chloride solutions onto the calcium carbonate pigments were conducted at room temperature, which is 23 0.5 "C in our laboratory. After removal from the rolling mill, the 8-02 wide-mouth round bottles stood on the laboratory bench until the calcium carbonate particles settled, then the supernatant solutions were pipetted into beakers. The adsorption of Caz+ ion onto the calcium carbonate pigments was determined by measuring the changes in concentration in solution before and after adsorption, using a Ca2+ion selectiveelectrode (ISEOrion, Mode193-20) connected to a pH/ISE potential analyzer (Orion, Model E A 940). The initial molarity of Gag+ ion ranged from 0 to 10 mM in seven stock adsorbate solutions, and four or five replicates were determined for each of the initial concentrations. The ion analyzer was precalibrated in the same [Ca*+]range as that in the adsorbate solutions and the linearity of the calibration was found satisfactory.
*
The pH of the solutions was determined also with the Orion ion analyzer. To measure the pH of the supernatant solutione, the electrode was swirled with its tip right at the boundary between the settled CaCOa and the supernatant. Adsorption-DataAnalysis. Since CaCOa is a slightlysoluble salt, the following consecutive chemical equilibria, together with the associationequilibrium of water, are taken into consideration. These equilibria and their equilibrium cOnstantseJOat 25 OC are as follows:
+ CO,"(aq); Ksp = 4.96 x lo4 M' Cog* + HZO G OH- + HCOC Kb,J = 2.25 X IO-' M HCO, + H,O F? OH- + H,C03; Kb,2 = 2.13 X lo* M
CaCO,(s) a Ca'+(aq)
(1)
(2) (3)
and
H,O + H,O
H30++ OH-; K, = 1.00 X lo-" Mz
(4) The concentrations in the equilibrium constants for eqs 1-4 are all in molarity (M). As can be seen, the amount of Caz+due to the dissociation of the dissolved CaCO3, [Ca2+]di,,,must equal the equilibrium concentration of COS*, [CO3*],, plus the amount of COS'- consumed in the dissociation reactions with water (eqs 2 and 3), [COS*]&; moreover, [COs*]di,, is equal to the sum of the equilibrium concentrations of HCOa- and HzCOs. Thus, we have F?
+[ C O , ~ ] ~
[c~'+I* = [co~*J,
(5)
and [CO:-l, = [HCO;], + [HzC031m (6) where the subscripts eq and diss stand for equilibrium and dissociation, respectively. The three equilibrium concentrations in eqs 5 and 6 can be evaluated from the equilibrium constants for eqs 1-3, respectively, and eq 5 can then be rewritten as
The depletion of calcium ion concentration due to adsorption, [Caz+]d, can be written as where the subscript ini stands for the initial Caz+ adsorbate concentration before adsorption. The amount of the calcium ion adsorbed onto the surface can be calculated by multiplying [Caa+]d by the volume of the adsorbate solution. The chemical equation for the adsorption process can be written as Ca2+(bulk)+ H20(surface) F! Ca'+(surface)
+ HzO(bulk)(g)
The linear form of the Langmuir adsorption isotherm can now be written aa in which r is the surface concentration of the adsorbed calcium ion (mol/m2), rmis the surface concentration of adsorption sites at monolayer saturation, and KC,as shown below, is the equilibrium constant for the adsorption processll n
where the term
(rm- l')
is the concentration of the surface
(9) Weast,R. C.,Aetle, M.J.,Beyer, W. H.,Eds. Handbookof Chemiutry and Phyuic8, 19th ed.; CRC Preas: Boca Raton, FL, 1988. (10) Skoog,D.A.; W&,D. M.~ndcrmontabofAnolyticalChemistry, 4th ed.; Saundera College Publishing: Philadelphia, PA, 1982. (11) h e n , M.J. Surfactant8 and Interfacial Phenomena; Wiley: New York, 1978.
1744 Langmuir, Vol. 7, No. 8, 1991
Huang et al.
Table I. Adsorption of Ca*+ Ion from CaCls Solutions onto CaCO: at 23 O C
r [Caz+]i,i, mol/dm3
0 1.07 X 1.07 X 1-07x 1.61 x 5.02 X 1.00 x
1od lo-' 10-3 10-3 10-2
[Ca2+],, m o l / d d [H+I,* PH (6.19f 0.04)X 1od 9.88 f 0.13 (9.76f 0.18)X 106 9.91 f 0.14 (1.04 f 0.17) X (2.46f 0.02)x (4.82f 0.05)X (2.77f 0.02)X (7.51f 0.01)X
lo-' lo-' lo-'
109 10-9
9.92 f 0.12 9.72 f 0.22 9.73 f 0.14 9.68 f 0.21 9.34 f 0.23
pmol/g of CaC03
f potential, mV 3.8 f 0.8 3.3 f 0.8 3.5 1.0 6.4 f 0.9 7.8 f 1.0 15.7f 1.4 18.7 f 1.5
mol/m2 [Caz+],/r, mol/mS/mol/m2 1.45 X lO-' 4.28 X 1@ 1.67 X 108 5.85 X 10-8 1.02 x 108 1.02 x 10-7 5.40 x 10-7 4.56 x 1@ 6.83 X 7.05 X l@ 1.31 X 10" 2.12 x 108 1.47X 10" 5.11 X 108
1.25 5.05 X 10-' t m x 10-1 4.66 5.90 11.3 12.7
Table 11. Various Concentration Terms (mol/dm*) Involved in Calculating the Adsorbed Ca*+ Ions [OH-], 7.59 x 10-6 8.13 X 10" 8.32x 5.25 x 5.37 x 10" 4.79 x l(r6 2.19 X
[HCOS-I, lo-' lo-' lo-' 1od
2.31 X 1.37 X 1.26 X 8.41 X 4.20 X 8.19 X 6.61 X
1od
lv 10-8
[HzCOSI~ 6.80 X 10-8 3.76 X 10-8 3.37 x 10-8 3.58 X 10-8 1.74 X 10-8 3.82 X 1O-g 6.74 X l@
adsorption sites occupied by the solvent molecules (H20). The units of the two surface concentration terms are cancelled as long as the units of r and rmare consistent; therefore, Kc has a unit of the reciprocalof the molar concentrationunit of [Caz+l,. As can be seen from eq 10, a plot of [Ca2+], versus [Ca2+l,/r should bear a straight line with its slope equal to the reciprocal of rm and the intercept equal to the reciprocal of the product of KCand rm;in other words rm = l/slope (12) and Kc = slope/intercept. (13) The standard state Gibbs free energy change of the adsorption process at the studied temperature T,ACo(T), is then given by ACo(r) = -RT In Kw
(14)
where R is the gas constant. K, is the equilibrium constant for the adsorption process, in which all the concentrations are expressed in terms of activities. At low concentrations, mole fractions can be substituted for the activities, so that Kq = x:/xfx: (15) where the x values are the mole fractions. The subscripts 1and 2 denotethe solvent (HzO) and the adsorbate (Ca2+), respectively. The superscripts S and B are abbreviations for the surface and bulk, respectively. In dilute adsorbate solutions,K, and KCare interconvertible,for the molar concentration of [Ca2+], in KC can be converted to the dimensionless mole fraction xzB in Kw It is obvious that K, is a dimensionless quantity. Results and Discussion Adsorption Isotherm. The adsorption data are tabulated in Table I with the directly measured quantities (four columns) on the left-hand side and data for the adsorption isotherm on the right-hand side. To calculate [Ca2+],/I', the unit of [Ca2+], (the second column) was converted from molarity to millimolarity (identical with mol/ms), which matches the unit of the surface concentration and the adsorbed Ca2+ ions (mol/m2). Various concentration terms involved in calculating the adsorbed Ca2+ions (eqs 5-8) are given in Table 11. It can be seen that as the bulk Ca2+ concentration increased, the amount of the Ca2+ions adsorbed onto the CaC03 0')increased smoothly, accompanied by an increase in the l potential. The exception occurred at the zero initial Ca2+ concentration where the equilibrium Ca2+ concentration was very close to (K,p)1/2= 7.04 X 10". A comparison of [Ca2+], to [Ca2+]i,i shows that for the first two initial Ca2+concentrations the CaC03 surface disso-
[Ca2+]& 3.12 X lo-' 1.88 X lo4 1.73 X lo-' 1.04 x 10-4 5.23 X lod 9.99 x 1o-e 7.28 X lW
[COsZ-Iw 8.01 X lo4 5.08 X lo4 4.77x 104 2.02 x 10-5 1.03 X 1W6 1.79 X 1o-e 6.61 x 10-7
~C~32-ldiU
2.31 X 1.37 X 1.26 X 8.42 x 4.20 X 8.20 X 6.62 X
lo-' lo-' lo-' 10" 10"
lv 10"
2.50 X 1.01 x 1.77 X 9.31 X 1.18 X 2.26 X 2.54 x
lo-' lo-' lo-' lo-'
109 109
109
"E "E
33 2 2
i u X
- 5
k \
N t
7
- 4
p
- 3
2
- 2
tQ
v)
a
u
Y
- kfl Ordinate M#hC Ordinate
0.0 0
'
I
I
I
I
1
1
2
3
4
5
6
- I
u
10 7
6
Equilibrium Calcium Ion Millimolarity
Figure 1. Langmuir isotherm for the adsorption of Ca2+ions from CaClz solutions onto CaCOa particles at 23 O C . ciation was prominent. In Figure 1, the surface Ca2+ concentration in micromoles per gram of CaCOs is plotted against the bulk Ca2+ equilibrium concentration (the dotted line). The point of zero initial Ca2+concentration was omitted from Figure 1. It is obvious that the adsorption isotherm approached a plateau at a sufficiency of the adsorbate. The solid line shows a linear relationship (eq 10)of the four points a t higher adsorbate concentrations. A fit of the adsorption data to the linear form of the Langmuir adsorption isotherm gives rm = 1.58 X lo4 mol/m2 (i.e., the surface Ca2+ concentration is 105 A2/ Ca2+ion or 13.6 pmol/g CaCOs) and Kc = 1.80 ms/mol ( K , = 9.99 X W). The goodness of the least-squares fitting is excellent, as indicated by a correlation coefficient of 0.9998. By a kinetic method, Compton and Pritchard have recently found that the adsorption of divalent copper ion on calcite is also Langmuirian.12 A comparison of rm with the I' values listed in Table I suggests that the adsorption achieved at the highest initial CaClz concentration (10.0 mM or the equilibrium concentration of 7.51 mM) was very close to the monolayer (12)Compton, R.G.;Pritchard, K.L.J. Chem. Soc., Faraday TTOM. 1990,86,129.
Langmuir, Vol. 7, No. 8, 1991 1745
Adsorption of Ca* onto CaCOs Particles 1000
E 7
.d
2 (d
0.1
2r u
3 0.01
0.001 0
10
20
30
40
50
80
70
80
90
100
Surface Coverage (%) Figure 2. Bulk Ca2+concentration as a function of the surface Cap+coverage in the CaCOs/CaCl*aqueous solution dispersions.
saturation (J?/Fm= 93%). At the monolayer saturation, rmaverages one Ca2+ion every 105A2of the CaCOssurface. Supposing the Ca2+ion sits at the center of the square, the distance from the center to the corner of the square is 7 A, which is plausible to be considered as in the ionic range. The surface density of Ca2+or Cos2- ions on the (1011) cleavage plane of calcite is 4.97 ions/ 100 A2 as calculated from crystallographic datal3 and we may therefore conclude that just 20% of the surface area is accessible to the adsorbing Ca2+ions. Since the surface monolayer contains only two components
(
