Surface electrical properties of calcite

Surface Electrical Properties of Calcite D U D L E Y W I L L I A M T H O M P S O N *'1 AND PAMELA G I L L I A N P O W N A L L ~ * School of Chemistry, University of Bristol, Bristol BS8 1Ts, England, and ~fCentralLaboratories, E.C,C. Ltd., Pentewan Road, St. Austell, England

Received June 27, 1988; accepted September 6, 1988 The electrical state at the calcite aqueous solution interface has been investigated using a streaming potential method applied to systemscontaining no gas phase. Experimentalresults have allowed a clear and unambiguous identification of the major surfaceions as Ca2÷and CO 2- speciesand have also shown that other solution species,including H ÷ and OH-, have no significant direct influence on the surface charge of calcite over the range pH 7-12, Induced changes in the pH of aqueous calcite dispersions in the directionsof increasingand of decreasingcalcitesolubilityhave been shownto result in the formation of a new solid phase at the calcitecrystalsurface.Formation of this surfacematerial leads to more positive values of the calcite ~'-potentialand it is suggestedthat this surfaceprecipitation phenomenon could be at least partly responsible for the diversity of experimental results which have been obtained from electrokinetic studies of aqueous calcite dispersions. © 1989AcademicPress,Inc. INTRODUCTION

lution interface have also led to conflicting opinions. Although results obtained from a n u m b e r of studies have indicated that Ca 2+ and CO32- are important surface ions (4, 9, 10) it has also been claimed that H ÷, O H - , H C O 3 , C a O H ÷, and C a H C O ~ play either major or m i n o r roles in the surface charging mechanism (4, 7, 9, 11 ). Most experimental studies carried out to date have been made using calcite dispersions open to atmospheric carbon dioxide (open system). With regard to the interpretation of experimental results, this system suffers the limitation that equilibrium concentrations of all calcium and carbonate ionic species having the same charge sign change in a similar m a n ner when the solution composition is changed by additions of acids and bases (Fig. 1 ) and also by additions of calcium and carbonate salts. Dispersions containing no gas phase ( d o s e d system), however, offer alternative systems which are more flexible with regard to chemical speciation (Fig. 2). In view of this, in this present work closed systems have been studied using a streaming potential method in an attempt to clarify the factors which control the electrical state at the calcite-aqueous solution interface.

The surface electrical properties of calcium carbonate, particularly those of calcite, dispersed in aqueous solutions have been studied using a n u m b e r of different electrokinetic and flotation techniques ( 1-11 ). Although these studies have provided some insight into the electrical state at this interface, they have also yielded m a n y inconsistent and apparently contradictory results. For example, widely diffeting ~'-potential values, including differences in sign, have been indicated for calcite crystals dispersed in solutions of similar composition (1, 4, 7, 8). Although attempts have been made to explain such variations in terms of the origins of different samples (2), no satisfactory correlation between sample source and surface electrical properties has been found. Furthermore, studies made on particular samples of calcite have shown that the magnitude and even the sign of the surface charge vary with different washing procedures ( 1 ) and also with the solid surface area/solution volu m e ratio ( 11 ). Investigations of the origin of the electrical charge at the calcite-aqueous so1To whom correspondence should be addressed. 74 0021-9797/89 $3.00 Copyright © 1989 by Academic Press, Inc. All fights of reproduction in any form reserved.

Journal of Colloid and Interface Science, Vol. 131, No. 1, August 1989

75

THE CALCITE S U R F A C E

"8 E ~

-2-0

~

-3.(]

"~/~A//oo.0o~//

/_

m

g m

\\

2 -5<

o -@0

"

,'

,,"V / A/

"+\\ { /X

k \//\

W\ X \

.to ,;o

p14

stored in sealed vessels under an atmosphere of CO2. Calcium analysis. Total concentrations of dissolved calcium were determined by EDTA titration (18). This analytical method was found to be satisfactory for calcium concentrations >~ ca. 10-5 mole dm-3. Carbonate analysis. Total concentrations of dissolved carbonate were determined by absorbing the CO2 liberated from known volumes of solution by reaction with phosphoric acid, in a standard solution of carbonate-free aqueous sodium hydroxide (18), and then titrating the latter with standard hydrochloric acid solution (19). This analysis was carried out under closed conditions in order to preclude absorption of atmospheric CO2 and was found to be satisfactory for total carbonate concentrations > ca. 10 -5 mole dm -3.

FiG. I. The influence ofpH on concentrations of ionic species in saturated aqueous solutions of calcite open to atmospheric carbon dioxide (pC02 = 3.3 X 10 -4 arm). (Equilibrium constants taken from Refs. ( 1 2 ) - ( 16 ). -3"0 -

MATERIALS AND METHODS -4"0

Synthetic calcite. Samples of this material were prepared by exposing aqueous solutions ofCaCl/to commercial "ammonium carbonate" vapor using the method described by Gruzensky (17). This preparation yielded rhombohedral crystals of 1-2 mm edge length, which were subsequently crushed and sieved, and the size fraction 106-150 ~m used for streaming potential measurements. Ca(OH)2 solutions. Powdered Ca(OH)2 (BDH Analar grade) was decarbonated by heating in a glazed porcelain crucible for ca. 5 min. Saturated solutions of this material in CO2-free water or aqueous salt solutions were prepared and stored in sealed vessels under an atmosphere ofCO2-free N2. Prior to their use, these solutions were filtered through a porosity-2 sintered glass disk. H2CO~ solutions. These solutions were prepared by bubbling CO2 through water or aqueous salt solutions until the solution pH was in the region 3-4. The solutions were

\

E

g -5.°

\

_1

-6"0 w o_ z o

B. -2.0

o co u. o

%

-.3.0

z

o F-

-5"0 7

8

--6"0 - ( I

4"0

6-0

8-0

10"0

12,0

pH FiG. 2. The influence of pH on the concentrations of ionic species in saturated aqueous solutions of calcite in systems containing no gas phase. (A) pH controlled by additions of HC1/NaOH. (B) pH controlled by additions of H2CO~ / Ca ( OH )2. Journal of Colloid and Interface Science, V o L

131, N o . 1, A u g u s t 1989

76

THOMPSON AND POWNALL

Total concentrations of calcium and carbonate determined using these analytical methods were found to be in excellent areement with those predicted using simple solution equilibrium models (Fig. ! ). Streaming potential measurements. These measurements were made using a closed (no gas phase) circulatory system (19) in which solutions were pumped through a packing of calcite crystals contained between two perforated disk Ag, AgCI electrodes. The pressure difference across the calcite packing (AP) was measured using Hg manometers, and measurements of solution conductivity and pH were made using a conductivity cell and a combination pH electrode which were sealed into the circulatory system. The difference in electrical potential developed between the two Ag. AgC1 electrodes (Es) was recorded for a series of AP values and g'-potentials were determined from the slopes of graphical plots of Es against AP according to the Helmholtz equation =

d"

a/'.

[11

equilibrated with aqueous solutions were washed rapidly with water and then with methanol, and allowed to dry. The crystals were mounted on specimen supports using graphite paste and were coated with a thin layer of gold or carbon prior to their examination in a Jeol 100CX Temscan electron microscope. Examinations of crystals which had not been equilibrated with aqueous solutions showed that this method of specimen preparation did not produce artifactual features on the crystal surfaces. RESULTS AND DISCUSSION On initial setting up of the streaming potential apparatus with freshly prepared samples of calcite the solution pH and calcite ~'-potential attained constant values within 48 h of the start of solution circulation. When the settingup procedure was repeated using fresh samples of calcite and electrolyte solutions having the same composition, however, it was found that the equilibrium values of pH and of ~'-potential varied quite significantly, as illustrated by the data presented in Table I. Total carbonate analysis of these solutions provided clear evidence that the variations in equilibrium pH values could be attributed to different amounts of atmospheric carbon dioxide absorbed by the electrolyte solutions during their preparation. The variable initial ~'-potential values are more

In this equation, E, n, and X represent the permittivity, the viscosity coefficient, and the specific conductivity of the solution phase, respectively. Measurements were made with solution flow in each direction in order to account for asymmetry potentials. Analysis of errors associated with the measurements TABLEI showed that ~'-potential values determined using Ecl.[ 1] were subject to a maximum error Equilibrium pH Valuesand ~'-PotentialsObtainedon in the region of 10%. Initial Equilibration of Calcite with Aqueous Sodium Preliminary investigations of the rate of at- Chloride(5.0 × 10-3 moledm-3) tainment of equilibrium in the calcitepH ¢(mv) aqueous solution system showed that constant values of solution pH and of total concentra9.11 --14.20 9.33 --5.01 tions of dissolved calcium and carbonate were 9.35 --3.26 attained within 48 h of the start of liquid cir9.43 --23.12 culation in the streaming potential apparatus. 9.44 --18.10 Significantly shorter equilibration times (< ca. 9.47 -- 18.96 15 h) were found for subsequent additions of 9.59 +2.35 9.61 --15.71 acid and base solutions.

Sample preparation for scanning electron microscopy. Calcite crystals which had been Journal of Colloid and Interface Science, Vol. 131,No. 1, August 1989

9.71 9.87

+ 1.90 --6.62

THE CALCITE SURFACE difficult to explain since they showed no systematic dependence on solution pH, dissolved calcium, or total dissolved carbonate concentrations. In addition, X-ray diffraction and electron probe microanalysis provided no evidence of impurity species or inhomogeneity in the calcite material. Attempts to establish uniform surface properties by treating calcite samples with gaseous carbon dioxide (20) and also by repeated washing with electrolyte solutions and with dilute mineral acids all failed to effect any substantial improvement in the reproducibility of initial ~'-potential values. It was found, however, that after the initial equilibration process, subsequent changes in solution composition result in the same relative changes in the f-potential regardless of the initial value of this potential. In view of this, streaming potential measurements were made on samples which had not been subjected to any surface-conditioning treatment. The experimental results shown in Fig. 3 were obtained by initially equilibrating calcite samples in electrolyte solutions containing sodium hydroxide and then lowering the solution pH by successive additions of hydrochloric acid. Qualitatively similar results were obtained when calcite was initially equilibrated with electrolyte solutions containing hydrochloric acid and the pH subsequently raised by additions of sodium hydroxide solutions, and also for systems in which the calcite was initially equilibrated in aqueous sodium chloride. Although the ~'-potential ranges encountered in these systems showed wide variations, it was found that, in general, ~--potentials developed in aqueous CaCI2 were significantly more positive than those developed in NaC1 and in NaC1/NaHCO3 solutions. While these systems yielded ~'-potentials which became increasingly negative with increasing solution pH, comparable results obtained for systems in which pH adjustments were made by additions of H2COs and Ca(OH)2 solutions exhibited well-defined minima in the ~'-potential/pH curves in the region pH I0. This is illustrated by the results presented in Fig. 4. The development of minima in these curves

77

10"0

5"0

0

> E

-5"0

~- -IO'O

< -lS.O

--20'0

-25.0

-30.0

I 7"0

I 8'0

I 9'0 pH

I 10"0

I 11"0

FIG.3. The influenceof solutionpH on the ~--potential of calcite dispersed in differentelectrolytesolutions. O, NaC1(5.0 X 10-3 moledm-3); D, NaC1(5.0 X 10-3mole dm-3)/NaHCO3 (1.0 X 10-3 moledin-3), A CaC12(5.0 X 10-4 mole dm-3), pH valuesadjustedby additionsof HC1/NaOH. could not be attributed to changes in the ionic strengths of the solution phases since over the range pH 7.5-11.5 the latter varied by no more than a factor of 1.5 in any of the experimental systems. Qualitative comparisons between the effects of pH on the magnitude of the ~'-potential developed at the calcite surface (Figs. 3 and 4) over the range of pH 7-10 and on the concentrations of dissolved ionic species (Fig. 2) clearly mitigate against HCO~ and CaOH + being important surface ions. Furthermore, in the case of systems in which pH adjustments were made by additions of H2CO3 and Ca (OH)2 solutions, a similar comparison for pH values > 10 provides a clear indication that O H - and CaliCO ~ species play no major direct roles in controlling the calcite surface charge. In fact, of the possible surface ions, only Ca 2+ and COl- species provide satisfactory qualitative correlations with changes in the f-potential over the range of solution pH 7-12. Further evidence of the potential controlling influence of the Ca2+/CO 2- couple is Journal of ColloM and Interface Science, Vol. I31, No. 1, August 1989

78

THOMPSON AND POWNALL Values for a d were determined using the standard G o u y - C h a p m a n equation, and values for K w e r e estimated from experimental values of dpCa/df, (Eq. 3).

10-0

8'0

6,O

K=

t.-0

(Coexp(KA))

( 2.303kT/ ze)( dpX/ df) - e x p ( r A ) '

[3]

2.0

A = distance between the outer Helmholtz plane ( O H P ) and the hydrodynamic plane of shear. k = Boltzmann constant T = absolute temperature z -- charge n u m b e r of surface ion e = primary electronic charge X = surface ion Ca = differential capacity of the diffuse region of the EDL. For small values of f, Cd -- ~K where K represents the reciprocal Debye length.

~ -2.0 -4.0 -B'0 -8"0 -10"0 I 8.0

I 9.0

pH

I 10-0

/ 11"0

FIG. 4. The influence of solution pH on the ~'-potential of calcite dispersed in aqueous NaC1 (5.0 × 10-3 mole dm-3)/NaHCO3 (1.0 × 10-3 mole din-3). ©, pH values Although the range of ionic strengths used in adjusted by additions of H2CO3;R, pH values adjusted this present work was insufficient to satisfacby additions of Ca(OH)2. torily evaluate A, values of K calculated for different experimental systems having ionic provided by the linear relationships between f-potential and p C a which are shown in Fig. A. 10"0 5A. Although the effective isoelectric points ( p C a ( I E P ) ) indicated by these data show a ~ o wide variation, the gradients, df/dpCa, obtained for different experimental systems fall i -10,0 within the narrow range 10.1-12.7 m V (Table II). In order to relate the experimental dr/ -20"O dpCa values to those expected for simple Nernstian behavior of Ca 2+ and CO ]- ions at I I I I I It}, I I I I l I the calcite surface, values of the surface po200 tential (~bo) were estimated by assuming a linear reduction in potential between the solid surface and the hydrodynamic plane of shear ,~ -20'(J in accordance with

s

-L,D,

~bo = f

~d K '

-6c-

[21

i /*-8

ad = charge in the diffuse region of the electrical double layer ( E D L ) . K = integral capacity o f the inner region o f the EDL. Journal of Colloid and Interface Science, Vol. 131, No. 1, August 1989

i /*'/,

/.-0

3.6 pC,a

3-2

2.8

2-.~

FIG. 5. Changes in ~'-potential and surface potential (1/o) with the activity of Ca2+ ions in saturated aqueous solutions of calcite. Solution conditions defined in Table II.

79

T H E CALCITE SURFACE TABLE II Calcite Surface Electrical Data Obtained from Streaming Potential Measurements Symbol (Fig. 5)

Electrolyte medium

d~/dpCa (mY)

d~o/dpCa (mV)

pCa (PZC)

[] A

NaC1/HC1/NaOH NaCI/NaHCO3/HC1/NaOH NaC1/CaC12/HC1]NaOH NaC1/CaC12/HC1/NaOH NaC1/HaCO3 NaC1/NaHCO3/H2CO3 NaC1/NaHCO3/Ca(OH)2

- 12.7 - 10.2 - 11.7 - 10.1 - 10.8 - 12.9 - 12.2

-34.5 -29.8 -25.5 -27.1 -29.3 -26.0 -28.1

2.02 1.92 2.16 3.40 4.00 3.80 3.80

O

V • • •

strengths of 2 X 1 0 - 3 - 1 0 - 2 mole dm -3 were found to be within the narrow range 0.14 + 0.01 Fro-2 when coincidence between the OHP and the plane of shear was assumed (A = 0). Relationships between pCa and ~0 values calculated using this inner layer capacitance are shown in Fig. 5B. Values of d¢~o/ dpCa obtained from these data (Table II) were found to be remarkably close to the Nernstian value expected for bivalent surface ions (ca. 29 mV at 20°C). These results, together with the observed qualitative dependence of the calcite ~'-potential on solution pH in systems closed to the atmosphere, would seem to provide clear and unambiguous evidence that the

surface charge of calcite dispersed in aqueous solutions is fundamentally controlled by Ca 2+ and CO 2- surface ions. In addition, the results of this study have enabled the positive exclusion of the species H +, CaOH +, CaliCOS, OH-, and HCO 3 as having significant direct influences on the surface electrical properties of the calcite-aqueous solution interface.

Hysteres& Effects in the pH Dependence of the Calcite F-Potential Adjustments of solution pH made in the direction either of increasing or of decreasing pH value resulted in steady, systematic changes in the calcite ~'-potential. Reversing

A.

B.

-5'0 -2-0

> -1"0.0 E

-4"0

~- -15"0 -6"0

o

o.. -20"0

tlJ N

-25'(

-8.0

-304

-10.0

6-0

I

I

I

7"0

8,0

9"0 pH

I

10"0

I

I

I

11"0 12-0

9'0

10-0

I

I

11.0

12.0

pH

FIG. 6. Hysteresis effects in ~--potential-pH cuives resulting from pH reversal (A) p H values adjusted by additions of HC1/NaOH; (B) p H values adjusted by additions of H2CO3/Ca(OH)2. A,, Initial equilibrium values. Journal of Colloid and Interface Science, Vol. 131, No. 1, August 1989

80

THOMPSON

Journal of Colloid and Interface Science, Vol. 131,No. 1, August 1989

AND POWNALL

THE CALCITE SURFACE

the direction of the pH change during the course of a particular series of measurements, however, invariably resulted in a marked hysteresis in the ~-potential-pH curve (Fig. 6). This effect, which always resulted in more positive ~'-potential values after pH reversal, was most evident in cases where the pH reversal was in a direction consistent with decreasing calcite solubility. It was found, however, that a similar, albeit smaller hysteresis effect occurred in some systems where the pH reversal resulted in an increase in calcite solubility. Chemical analysis showed that all pH adjustments made to these systems resulted in total dissolved calcium and carbonate concentrations which were in excellent agreement with those expected for saturated solutions of calcite at the respective pH values. In this case it seems improbable that hysteresis in the ~'potential-pH curves can be attributed to nonequilibrium solution conditions, but is more likely to arise from changes in the nature of the calcite crystal surface. In an attempt to clarify this situation the effects of different solutions on the physical state of the calcite crystal surface were investigated using scanning electron microscopy (SEM). Examination of calcite crystals prior to their contact with aqueous solutions showed the crystal surfaces to be reasonably smooth, with widely spaced linear growth steps on uncleaved crystal faces. Equilibration of these crystals with aqueous solutions under closed system conditions resuited in the appearance of small protrusive growths on the crystal surfaces (Fig. 7A). Additions of HC1 or H2CO3 to these systems re-

81

suited in a certain amount of etching of the crystal surfaces together with a small but significant increase in the population density of surface growths (Fig. 7B). Subsequent additions of NaOH or Ca(OH)2 solutions led to a large increase in the surface growth population density (Figs. 7C and 7D). Continued additions of NaOH resulted in an increase in the size and surface coverage of the needle-like growths. At high concentrations of added Ca(OH)2, however, the surface material became more crystalline in appearance and, eventually, well-formed rhombohedral crystals appeared on the calcite surfaces (Figs. 7E and 7F). Powder XRD patterns obtained from crystals having needle-like surface growths exhibited strong reflections characteristic of the calcite structure together with very much weaker lines characteristic of aragonite. Previous work has indicated that the aragonite surface carries a more positive charge than does calcite in solutions of similar composition (2, 3, 9 ). Thus the formation of a surface aragonite phase would seem to be consistent with the observed development of more positive ~'potentials under conditions where a profusion of surface growths occurred. It must be conceded, however, that identification of the needle-like surface material as aragonite is somewhat tenuous owing to rather small differences between the intensities of aragonite diffraction lines exhibited by samples containing these surface growths and by samples of untreated calcite crystals. The rhombohedral morphology of the surface particles formed in Ca(OH)2 solutions (Fig. 7F) strongly suggests that these

FIG. 7. Scanning electron micrographs of calcite crystal surfaces after exposure to aqueous solutions. Supporting electrolyte

pH adjustments

None

A. NaC1 (5 X 10 -3 mole d m -3) HCI --~ pH 7.5 UCl --~ pH 7.5

B. CaC12 (5 X 10-4 mole dm -3) C. CaC12 (5 x 10-4 mole dm -3) D. NaCI (5 X 10 -3 mole d m -3) + NaHCO3 (10 -3 mole dm -3) E. NaC1 (5 X 10 -3 mole d m -3) + NaHCO3 (10 -a mole dm -3) F. NaC1 (5 × 10-3 mole dm -3)

H2CO3 H2CO3

NaOH ~" p H 9 . 9

pH 7.3

c~(o~ pH 8.1

pH 7.2

c~(o~2 pH 8.3

Ca(OH)2

pH ll.6

Journal of Colloid and Interface Science, Vol. 131,No. I, August 1989

82

THOMPSON AND POWNALL

are particles of calcite, and this is certainly consistent with results obtained from X R D analysis of samples containing large amounts of this surface material. While precipitation in some form or another is to be expected in calcite dispersions under conditions of decreasing calcite solubility, resuits obtained in this work have shown quite clearly that precipitation can also occur under conditions where overall dissolution of calcite occurs. The reason for this is not entirely clear; however, it seems quite possible that this phen o m e n o n is an example of the subsurface precipitation process which has been theoretically predicted and shown to occur during the dissolution in protonic acids, of ionic solids which effect pH-buffering of aqueous solutions (22, 23). While the results obtained from this present study have not allowed positive identification of the surface phase it is clear that this material significantly modifies the surface electrical properties of calcite crystals. In view of this, variations in surface coverage by this material are almost certainly responsible, to a large extent, for the marked variations in ~'potential ranges and p C a ( I E P ) values obtained in this present work (Table II). CONCLUSIONS

Results obtained from streaming potential measurements made on calcite-aqueous solution systems containing no gas phase have shown that the major calcite surface ions are Ca 2+ and CO 2- species. No evidence has been found for participation of the complex species C a O H ÷ and C a l i C O ~ in the development of electrical charge at this interface. The well-established dependence of the calcite ~'-potential on solution p H almost certainly arises from the influence of p H on the solution concentrations of the lattice ions. The formation of new surface material which modifies the physical and electrical properties of calcite crystal surfaces would seem to be a c o m m o n feature of aqueous calcite dispersions. In view of this it is clear that in studies of the calcite-aqueous solution interface careful attention must be paid to experimental conditions and to the Journal of Colloidand InterfaceScience, Vol.131,No. I, August1989

immediate history of the calcite surface. It is possible that this surface precipitation phen o m e n o n could be at least partly responsible for the widely differing results which have been obtained from electrokinetic studies of aqueous calcium carbonate systems. REFERENCES 1. Douglas, H. W., and Walker, R. A., Trans. Faraday Soc. 46, 559 (1950). 2. Berlin, T. S., and Khabakov, A. V., Geochemistry 3,

217(1961). 3. De Groot, K., and Duyvis, E. M., Nature (London) 272, 183 (1966). 4. Somasundaran, P., and Agar, G. E., J. Colloid Interface Sci. 24, 433 (1967). 5. Fuerstenau, M. C., Gutierrez, G., and Elgillant, D. A., Trans. Amer. lnst. Min. Eng. 241, 319 (1968). 6. Yarar, B., and Kitchener, J. A., lnst. Min. Metall. Trans. Sect. C 79, 23 (1970). 7. Ney, P., "Zeta Potentials and flotierbarkeityon mineralen." Springer-Verlag,New York, 1973. 8. Weigl,J., and Huggenberger,L., Wochenbl. Papierfabr. 23, 886 (1974). 9. Smallwood,P. V., Colloid and Polym. Sci, 255, 881 (1977). 10. Foxall,T., Peterson, G. C., RendaU,H. M., and Smith, A. L., J. Chem. Soc. Faraday Trans. 175, 1034 (1979). 11. Siffert,B., and Fimbel, P., Colloids andSurf 11, 377 (1984). 12. Harned, H. S., and Davies,J. R., J. Amer. Chem. Soc. 65, 2030 (1943). 13. Harned, H. S., and Scholes, S. R., J. Amer. Chem. Soc. 63, 1706 ( 1941). 14. Greenwald, I., J. Biol. Chem. 41, 789 ( 1941). 15. Amankonah, J. O., Somasundaran, P., and Ananthapadmabhan, K. P., Colloids and Surf 15, 309 (1985). 16. Parsons, R., "Handbook of Electrochemical Constants." Butterworths, London, 1959. 17. Gruzensky, P. M., J. Phys. Chem. Solids Suppl. 1, 365 (1967). 18. Vogel,A. I., "A Textbook of Quantitative Inorganic Analysis," 4th Ed. Longmans, Green, New York, 1978. 19. Pownall, P. G., "The Surface Electrical Properties of Calcium Carbonate." Ph.D. thesis, University of Bristol, 1987. 20. Goujon, G., and Mutaftschiev, B., J. Colloid Interface Sci. 57, 148 (1976). 21. Smith, A. L., J. ColloidlnterfaceSci. 55, 525 (1976). 22. Cussler, E. L., and Featherstone, D. B., Science 213, 1018 (1981). 23. Leaist,D. G., J. Colloid Interface Sci. 118, 262 (1987).