Langmuir 1988,4,855-861 izabilities, e.g., He and Xe, respectively. Our experimental data with both gases are consistent with this theoretical treatment. Also, the relative orientation of the dominating components of the dipole moment has been estimated by comparing the values of the tangential absorptive index obtained with Xe and He. This model study shows the potential of photoacoustic FT-IR spectroscopy for obtaining at least semiquantitative data on the orientation of Langmuir-Blodgett films, and
855
we believe that it serves to illustrate the general applicability of the Rosencwaig-Gersho theory and the DignamRoth treatment of the photoacoustic effect.
Acknowledgment. We gratefully acknowledge the Army Research Office for supporting this research under Grant DAAG29-85-K-0225 and the CWRU-MRL under Grant DMR-8119425. Registry No. D-MAD, 114378-24-4;D-l8TH, 114378-25-5.
Precipitation of Calcium Carbonate in Aqueous Solutions in the Presence of Oxalate Anions Efthimios K. Giannimaras and Petros G. Koutsoukos* Department of Chemistry, Physical Chemistry Laboratory, Research Institute of Chemical Engineering and High Temperature Chemical Processes, University of Patras, P.O. Box 1239, GR-261 10 Patras, Greece Received December 23, 1987. I n Final Form: March 10, 1988 The precipitation of calcite, the thermodynamically most stable phase of calcium carbonate, was investigated at pH 8.50 and 25 OC in stable supersaturated calcium carbonate solutions seeded with calcite. The precipitation, both in the absence and in the presence of oxalate anions, was studied by maintaining solution supersaturation by the addition of titrant solutions with the stoichiometry of calcite. This methodology allows for very small rates to be accurately evaluated and the effect of inhibitors of crystal growth to be precisely assessed. The kinetics of crystal growth of calcite seed crystals pointed to a surface-controlled,spiral growth mechanism with an apparent order of 2. The presence of relatively low amounts of oxalate anions in the supersaturated solutions suppressed significantlythe rate of formation of calcite on synthetic calcite seed crystals, altered the apparent order of reaction from 2 to 4, and favored the formation and subsequent stabilization of calcium carbonate monohydrate. Application to a kinetic Langmuir-type model suggested that adsorption of oxalate at the active growth sites is responsible for the reduction in the crystal growth rates. Adsorption studies revealed that oxalate adsorbs on calcite, and the Freundlich equation was found to give a satisfactory fit of the data at ionic strengths of 0.01 and 0.1 mol dm-3. The ionic strength dependence suggested that the calciteoxalate interaction is mainly electrostatic in its nature. Adsorption of oxalate at the calcite/water interface resulted in negative electrokinetic charges of a magnitude practically independent of pH.
Introduction The formation of calcium carbonate is an important process in the natural environment, in wastewater treatment, and in physiological situations such as the formation of gallstones. As the formation of calcium carbonate takes place in the presence of a variety of constituents, including organic and inorganic anions and cations, their effect on the kinetics of calcium carbonate may be important for the control of the rate and of the extent of precipitation. Organic components in the natural aquatic environment are produced either directly or indirectly from photosynthesis or degradation. In all cases, the presence of the organic molecules not only has an inhibitory effect on the precipitation process1V2 but lowers the effective supersaturation with respect to the various calcium carbonate polymorphs by complexing calcium ions. Another effect that the organic solvents may have is the stabilization of one of the calcium carbonate crystalline phases, which include in order of decreasing solubility calcium carbonate hydrates (hexa and mono), vaterite,
aragonite, and calcite, which is the phase most commonly encountered. The nature of the solid forming in most cases is determined by kinetics rather than thermodynamic^.^ It is very likely that the formation of calcium carbonate both in the environment and in living organisms occurs heterogeneously on already existing surfaces. For this reason, overgrowth studies of calcium carbonate on wellcharacterized surfaces are very useful in understanding the mechanism of the formation of the various polymorphs and the factors governing the kinetics. Thus a number of studies of calcium carbonate formation have been done in solutions of relatively low supersaturations, seeded with well-characterized calcite These studies have shown that both in the absence and in the presence of additives the only phase forming is calcite. However, in most of these studies, the reaction was studied at constant pH, and changes in calcium concentration were measured (3)Finlayson, B.;Khan, S. R.; Hackett, R. L. Scanning Electron Microsc. 1984,1419. (4)Nancollas, G. H.; Reddy, M. M. J. Colloid Interface Sci. 1971,37,
-- -.(5) Reddy, M. M. J. Cryst. Growth 1977,41, 287. R3A
(1)Reddy, M.M.In Chemistry of Wastewater Technology; Rubin, A. J., Ed., Ann Arbor Scientific: Ann Arbor, 1978; p 31. (2)Nancollas, G. H.; Tomson, M. B.; Battaglia, G.; Wawrousek, H.; Zucherman, M. In Chemistry of Wastewater Technology; Rubin, A. J., Ed., Ann Arbor Scientific: Ann Arbor, 1978;p 17.
(6)Kazmierczak, T.F.; Tomson, M. B.; Nancollas, G. H. J. Phys. Chem. 1982,186,103. (7) Giannimaras, E. K.; Koutsoukos, P. G. J. Colloid Interface Sci. 1987,116,423.
0743-746318812404-0855$01.50/0 0 1988 American Chemical Society
Giannimaras and Koutsoukos
856 Langmuir, Vol. 4, No. 4 , 1988
as the supersaturated solutions proceeded to equilibrium followingthe introduction of seed crystals. The application of the constant solution composition technique, however! to the study of a calcium carbonate system6 offered the possibility of more detailed studies on the formation of various crystalline phases. In the present work we have applied the constant solution supersaturation technique in order to study the effect of organic compounds on calcite, since it is believed that this type of compound, possibly by adsorption, forms interfaces with calcite of markedly reduced energy, thus catalyzing the nucleation of any of the calcium carbonate polymorph^.^ We have selected oxalate ions in our study both as a model for carboxylic acids encountered in wastewater and also because of their presence in biological fluids, in media where calcium carbonate formation is seen to take place. Experimental Section Crystal Growth Experiments. All experiments were done at 25 f 0.1 "C in a ca.0.250-dm3water-thermostated double-walled Pyrex vessel. Triply distilled, COz-free water wm used throughout the experiments. Stock solutions of calcium nitrate and pofassium nitrate were prepared from solid reagents (Fluka AG., Purum and Merck, reagent grade, respectively) that were recrystallized twice from triply distilled water and filtered through 0.22-bm membrane filters (Millipore, Bedford, MA). Calcium nitrate solutions were standardized by using atomic adsorption spectroscopy (Varian 1200), with spectrophotometric titrations and EDTA standard solutions, a t 520 nm (using nurexide as indicator) as well as by passing aliquots of solution through a cation-exchange resin (Dowex 50W-X8 BDH Chemicals, Ltd) in the hydrogen form and titrating the eluted acid with standard base. Potassium hydroxide solutions (Titrisol, Merck) were standardized against potassium hydrogen phthalate dried overnight at 105 "C by using phenolphthalein indicator. Stock sodium bicarbonate solutions were made fresh each time from the solid (Fluka AG., Puriss) dried a t 105 "C for 5 h. Supersaturated solutions were prepared in the double-walled reaction vessel by adding the appropriate volumes of stock calcium nitrate and sodium bicarbonate solutions. The volume of the supersaturated solutions employed was 0.200 dm3, so as to minimize the air above the working solution. pH was adjusted to 8.50 by the addition of standard 0.1 mol dm-3 potassium hydroxide solution. The high p H of the solution employed along with the small volume of air over the supersaturated solutions allowed us to consider the partial pressure of carbon dioxide as constanL6 The stability of the supersaturated solutions was ensured by the stability of the adjusted p H value over extended periods of time. Working solutions were stirred with a Teflon-coated magnetic stirrer, and pH was measured by using a pair of glass (Radiometer G 202C) and Hg/Hg2Cl2/C1- (Metrohm) electrodes in the cell GE/supersaturated solution/Hg/Hg2C12/C1The pair of electrodes used was standardized before and after each experiment by NBS buffer solutions a t pH 6.865 (0.025 m KH2POl + 0.025 m Na2HP04)and a t p H 9.810 (0.01 m Borax).lo Oxalic acid stock solutions (Titrisol, Merck) were standardized by standard potassium permanganate solution and by standard potassium hydroxide solution in the presence of phenolphthalein indicator. Precipitation reaction was initiated by the introduction in the supersaturated solution of a weighted amount (ca. 50 mg) of well-characterized calcite seed crystals. The seed crystals employed were prepared by the slow addition of sodium carbonate (Vioryl, Pro Analysi) in a calcium nitrate solution in a manner analogous to that described by Reddy and Nancollas." The aged solid was characterized by X-ray diffraction (Phillips XRG-3000 (8) Tomson, M. B.; Nancollas, G. H. Science (Washington,DC)1978, 200,1059. (9) Carlson, W. D. Rev. Mineral. 1984,11,191. (10)Bates, R. G.PH Determination, 2nd ed.; Wiley: New York, 1973. (11)Reddy, M.M.; Nancollas, G. H. J. Colloid Interface Sci. 1971,36,
166.
Cu K a radiation with Ni filter), infrared spectroscopy (PerkinElmer grating spectrophotometer Model 177),and scaning electron microscopy (Jeol JSM 35). The BET specific surface area as determined by a multiple-point analysis (Perkin-Elmer sorptometer, Model 212D) was found to be 3.26 m2 8-l. Following the addition of seed crystals precipitation started immediately, without an induction period, resulting in decreasing pH of the working solution. Decrease of pH, measured by a pH meter (Radiometer 26C), as small as 0.5 mV triggered the addition of titrant solutions from two mechanically coupled burets of an appropriately modified pH-stat (Radiometer TT 11, ABU Ib). The buret piston was mechanically coupled with a recorder (Radiometer SBR 2c), and the volume of titrant added as a function of time was thus recorded. The titrants had the stoichiometry of the solid phase forming, that is, calcite (Ca:C03= 1:l). The concentrations of titrants were buret 1: [Ca(N03)z]= (r + 2)[Ca], mol dm-3
(1)
[Na2C03]+ [NaHC03] = ( x + 2)[NaHC03], mol dm-3 (2)
buret 2:
In eq 1 and 2 brackets denote concentrations, s stands for the working solution, and x is a number 21. In this work x = 10 throughout the experiments. In addition, potassium nitrate with total concentration [KNOB]= 2x[CaIB3was added as inert electrolyte in the working solution to maintain constant ionic strength, since changes in activity coefficients could also trigger the titrant addition." During the course of the reaction samples were withdrawn so as to the keep the total volume constant and were filtered through 0.22-pm membrane fdten (Millipore, Bedford, MA). The filtrates were analyzed for calcium by spectrophotometric titrations or atomic adsorption spectrometry in order to verify the constancy of the composition. The solid residues were examined by the physicochemical methods mentioned before. In experiments in which the effect of oxalate anion was investigated, the additive was introduced both in the working solution and in buret no. 2 so t u to avoid dilution due to the addition of titrants during the course of the precipitation reaction. Both in the absence and in the presence of oxalate, potassium hydroxide was added in buret no. 2 as required for the pH adjustment of the working solution a t the desired value. Inclusion of the necessary oxalate concentration to avoid dilution helped maintain the concentration of the additive constant in solution, as the kinetics of growth were faster than the kinetics of adsorption. The precipitation rates were obtained from the straight lines obtained from the recordings of the volume of titrant added as a function of time, corrected for the total surface area of the seed crystals. The rates were taken for the addition of the first 0.002 dm3 of titrants, which correspond to extent of growth, such that the BET specific surface area of the seed crystals is not changed due to crystallite perfection. Moreover, the rate a t very low extents of growth was preferred for the sake of comparison with experiments in which the overall extent of growth was small due to suppression of the crystallization process. Adsorption Experiments. Adsorption experiments were done in 0.015-cm3 stoppered glass centrifuge tubes, as well as in 0.030-dm3screw capped polyethylene tubes. In both cases, in each tube, a carefully weighed amount of calcite was introduced corresponding to a total surface area of approximately 0.08 m*. Next, a 0.010-dm3 volume of potassium oxalate solution, of concentrations ranging from 1 X to 5 x mol dm-3, was introduced, and the solid was dispersed by vigorous shaking. The pH of the suspension was a function of the potassium oxalate concentration. The ionic strength of the oxalate solutions was adjusted by potassium nitrate. Adsorption experiments were done at ionic strengths of 0.01 and 0.1 mol dm-3. The pH was measured in each tube by a combination glass/calomel electrode (Metrohm). Subsequently, the tubes were rotated end over end for a period of 24 h at 25 f 0.5 "C. Following the equilibration period, the pH was measured again and the suspensions were centrifuged at 3000g (MSE centrifuge). The aqueous phase was separated from ~
~~~
(12) Koutsoukos,P. G.; Amjad, Z.; Nancollas, G. H. J . Colloid Interface Sci. 1981,83,599.
Langmuir, Vol. 4 , No. 4, 1988 a57
Precipitation of Calcium Carbonate
exP
C1D C2A
Table I. Crystallization of Calcite on Calcite Seed Crystalsa [Cat], 10-3(moldm-3) AGM AGV AGA AGC 1.26 +0.34 +0.28 -0.60 -1.17 1.59 1.87 2.00
C3A C4A
-0.13 -0.46 -0.60
-0.19
-1.06 -1.39 -1.53
-0.52 -0.65
R, W ( m o 1 m-2 min-') 2.76 8.67 10.22 14.58
-1.63 -1.96 -2.10
O25 OC, pH 8.50;total calcium [Cat] = total carbonate [C,]; AG values are in kJ mol-'.
the solid residue and was analyzed for oxalate with ferric 5-nitrosalicylate complex spectrophotometrically a t 492 nm. The accuracy of the method employed was 0.5%.13 The calcite residues were repeatedly washed with potassium nitrate solution of the same concentration as the ionic strength of the adsorption medium and were dried overnight at 105 "C. The dry crystals were analyzed by infrared spectroscopy by the KBr pellet method and by powder X-ray diffraction. Electrokinetic Measurements. The electrophoretic mobilities of calcite particles suspended in 0.01 mol dm-3 potassium nitrate in the presence of oxalate were determined in a Rank Brothers MKII apparatus using a four-electrode capillary cylindrical cell. The velocities of at least 20 particles (in both directions of the electric field) in each of the two stationary layers were measured.
Results and Discussion The concentrations of ionic species in the supersaturated solutions were calculated from the experimental pH values and mass balance expressions for total calcium (CaJ, total carbonate (C,), and total oxalate (0,)by successive approximations for the ionic strength as described previ0us1y.~ Activity coefficients, yz, of z-valent ions were calculated by using the modified Debye-Huckel equation proposed by Davies.14 The calcium carbonate system is complicated by the relatively slow gas/solution equilibrium: C02 + H2O = H+ + HCOc
(3)
In the present work, this problem was avoided by keeping to a minimum the volume of the gas space above the solution in the crystallization cell and because of the high value of pH (8.50) of the supersaturated solutions. On this account, the system was assumed to be closed. The effective isolation of the system was verified by the constancy of the pH of the supersaturated solutions for periods of hours in the absence of inoculating seed crystals. The degree of supersaturation, 3, is defined as 3=
(Ca2+)t(C032-)t (Ca2+),(C032-),
PH
Figure 1. Solubility isotherms of calcite, aragonite, vaterite, and calcium carbonate monohydrate. A experimental conditions. Shaded area shows lower limit for spontaneous precipitation.
(4)
In eq 4 parentheses denote activities at time t and at equilibrium, m. The driving force for the crystal growth process, which is the change in Gibbs free energy for transfer from the supersaturated to saturated solution, is given by AG = -(RT/2) In Q (5) The rate of crystal growth was measured directly from the volume of titrant solutions added. The experimental conditions are summarized in Table I. As can be seen from Figure 1 also, we have worked in supersaturated solutions stable over long periods of time (stable at least for a week). The degree of supersaturation was >1for all calcium carbonate polymorphs (region A in Figure 1). The solubility isotherm for the calcium carbonate monohydrate was calculated from literature solubility data,'5 from which (13) Lee, K. S.; Lee, D. W.; Hwang, J. Y. Anal. Chem. 1968,40,2049. (14)Davies, C. W. Ion Association; Butterworths: London, 1962. (15)Dalas, E.;Koutsoukos, P. G. J. Cryst. Growth, in press.
-6.0-
o 0 -O.Orn~l.drn-~ tOt=3.5r10-3 r n ~ l . d r n - ~
-6.5-
-7.0
- 4%
I
-4.5
I
- 4.4
I
I
- 4.3 -4.2
I
-4.0
- 44
log{./Eziq5
-
-3.9
fi I
Figure 2. Kinetics of crystal growth of calcite seed crystals, both and in the presence ( 0 )of oxalate anions. Plot in the absence (0) of the logarithm of the rate of crvstal growth as a function of the logarithk of the supersaturation* exprksion: [ (Caz')t(C0,z-)t]1/2 - [ (Ca2+),(CO~-),]1/z.
a thermodynamic solubility product, K O , , = 1.279 X (mol dm-3)2,was calculated after taking into account ion pairing. The kinetics data were fitted to the semiempirical rate equation:
R = -(d[Ca,]/dt) = ks([(Ca2f)t(C032-)t]1/2 - [(Ca2'),(C0,2-),]1/2)n (6) In eq 6 k is the rate constant for the crystal growth reaction, s is proportional to the total number of available growth sites on the added seed crystals of calcite, and n is the apparent order of the crystal growth reaction. Logarithmic plots of the rate, R , as a function of the right-hand side of eq 6 yielded a straight line, shown in Figure,2, from which an apparent order of n = 2 was obtained. This result is in agreement with previous studies for calcium carbonate and other 2:2 sparingly soluble salts.l@ This order has been interpreted as suggesting a surface diffusion controlled reaction. By use of the con-
858 Langmuir, Vol. 4, No. 4, 1988
Giannimaras and Koutsoukos
Table 11. Crystallization of Calcium Carbonate Monohydrate on Calcite Seed Crystals in the Presence of Oxalate (0,)" exp [ O , ] ,104(mol dm-? AGM AGv AGA ACC AGCOM R, 104(mol m-2min-') C3A C70A C80A C60A C90A
" [Ca,] = [C,]
-0.46 -0.45 -0.44 -0.44 -0.41
2.50 3.50 5.00 10.00
= 1.87
mol dm-3; 25
X
OC,
-0.52 -0.51 -0.50 -0.49 -0.47
-1.39 -1.38 -1.38 -1.37 -1.34
-1.96 -1.95 -1.95 -1.94 -1.92
10.22 8.18 3.84 0.77 0.29
+8.92 +8.51 +8.07 +7.23
pH 8.50; AG values are in kJ mol-'.
Table 111. Crystallization of Calcium Carbonate Monohydrate on Calcite Seed Crystals at Constant Supersaturation in the Presence of Oxalate Ions" extent of time, [Calt, crystallization, 70 min
10-3(moldm-3)
of original seed
0 5 50 90
1.87 1.88 1.88 1.89
1.2 17.2 40.8
r
XRD
1
Calcite Seeds
"Experiment no. C80A: total calcium [Cat] = total carbonate mol [C,] = 1.87 X mol dm-3; total oxalate [O,] = 3.5 X dm-3; pH 8.50, 25 "C. stant-composition method it was possible to grow extensive amounts of calcite, more than twice the amount of the seed crystals. It was thus confirmed that under our experimental conditions calcite was the only phase forming, although the formation of aragonite, vaterite, or monohydro calcium carbonate could be thermodynamically feasible. Furthermore, the experiments of crystal growth of calcium carbonate on calcite seed crystals were done in the presence of oxalate anions. The initial conditions are summarized in Table 11. I t may be seen that relatively low amounts of oxalate suppressed significantly the rate of calcium carbonate overgrowth on the calcite seed crystals. In all cases the solution composition was kept constant, as may be seen from the typical results of a crystal growth experiment of crystallization of calcium carbonate in the presence of oxalate ions, summarized in Table 111. It should be noted that, in the presence of oxalate, the formation of a number of calcium oxalate solid phases is possible. These phases include, in order of increasing solubility, whewellite (CaC2O4.H20,COM), weddellite (CaC204-2H20,COD), and calcium oxalate trihydrate (CaC204-3H20, COT). Our experimental conditions were such that in no case was the solubility product of any of the calcium oxalate phases exceeded, as may be seen in Table 11. The calculations in this table have taken into account calcium oxalate ion-pair formation. Kinetics plots, however, shown in Figure 2, suggested that the presence of oxalate anions in solution altered the apparent order of the precipitation reaction from n = 2 to n = 4. Further examination of the solid phases formed, revealed that the presence of oxalate favored the formation and subsequent stabilization of calcium carbonate monohydrate. The formation of this calcium carbonate polymorph was spectroscopically confirmed, since by using the constant solution composition technique we were able to grow extensive amounts of the overgrowing phase, sufficient for an adequate characterization. In Figure 3 the powder XRD spectra of calcite and of the calcium carbonate monohydrate overgrowth are shown. As can be seen, the new peaks appear at d = 6.02, 5.86, 3.66, and 2.03 A. The strongest peak is that corresponding to d = 2.03 A. The appearance of the new peaks in the powder X-ray diffractogram rules out the possibility of incorporation of the oxalate ions into the calcite lattice. It should also be noted that the same powder X-ray diffractogram was obtained
50
45
40
35
30 25 go/dcg.
20
15
10
5
Figure 3. Powder X-ray diffraction spectra of calcite, calcite grown on calcite seed crystals, and calcium carbonate mono-
hydrate.
BO
100
120
140
160
Temperature / 'C
Figure 4. DSC analysis results for calcite (a) and for the calcium carbonate monohydrate grown in the presence of oxalate anions (b).
for the growth of calcium carbonate monohydrate on polymeric surfaces by heterogeneous surface nuc1eati0n.l~ Moreover, differential scanning calorimetric and thermogravimetric analyses have also confirmed the formation of calcium carbonate monohydrate, as can be seen in Figure 4. The loss of water at about 120 "C is characteristic for calcium carbonate monohydrate.16 Calculations from thermogravimetric analysis showed that one water molecule is lost at this temperature. The formation of any higher calcium carbonate hydrate is thus precluded since the differential scanning calorimetric and thermogravimetric analysis data would be markedly different. So far, the formation of calcium carbonate hydrates has been observed only in precipitation from highly supersaturated calcium carbonate s01utions.l~ The scanning electron micrographs in Figure 5 show the morphology of calcium (16) Hull, H.; Turnbull, A. G. Geochim.Cosmochim.Acta 1973, 37, 685. (17) Brooks, R.; Clark, L. M.; Thurston, E. F. Philos. Trans. R. SOC., London Ser. A 1950,243, 145.
Precipitation of Calcium Carbonate
Langmuir, Vol. 4, No. 4, 1988 859
C;
IO-' r n o ~ . d r n - ~
Figure 6. Langmuir-type kineties plot for the crystal growth of calcite seed crystals in the presence of oxalate anions. R, is the rate of -tal growth in the abaence of Ci mol dm+ oxdate anions in the solution and Ri in its presence. pursued the investigation of the adsorption of oxalate on calcite at 25 "C. In order to avoid undesirable changes in ionic strength we avoided adjusting the pH of the solutions in which adsorption took place. Thus the adsorption pH was in all samples 6.1 0.3. This initial pH remained unchanged during the course of the adsorption process even though oxalate was extensively adsorbed on calcite. It should he noted that oxalate was adsorbed significantly on the surface of the tubes we used, either glass or polyethylene. Thus in each run we had to run blank samples in order to evaluate the true adsorption on calcite. The extent of adsorption was comparable both for glass and polyethylene, but in the former case the results of the adsorption were very irreproducible and a t the same time showed significant scattering. Thus we have done the experiments in polyethylene vials. A kinetic analysis of the adsorption showed that it was complete within 5 h. After this time period, oxalate concentration in solution reached a steady level, and pH did not show any change. The adsorption isotherms at 0.01 and 0.1 mol dm" KNO, are shown in Figure 7. Adsorption data gave a satisfactory fit to the Freundlich isotherm: log r = log (kr,) + (i/n) log c , (7)
*
I,
Figure 5. Scanning electron micrographs of calcite (a) and calcium carbonate monohydrate crystalsgrown in the presence of oxalate anions (b). carbonate monohydrate growing in the presence of oxalate ions. It is believed that the surface of the calcite seed crystals is modified energetically, probably by adsorption of oxalate anions, thus stabilizing the formation of the monohydrate even a t very low degrees of solution supersaturation. Reevaluation of the kinetics data, considering the rates as the initial rates of calcium carbonate monohydrate, yielded an apparent order of approximately 2 for this salt. Application of the kinetics results, on a simple kinetic Langmuir-type mode1,'sJ9 yielded the straight line shown in Figure 6. Application of the Langmuir model in this case is an oversimplification, because the prerequisites for the application of such a model are not valid. However, plots like that of Figure 6 are useful for comparative purposes, since the slope of such plots is a measure for the affinity of a certain inhibitor for the Darticular solid substrate.I2 From the slope of the line in Figure 6 a value of 47.70 X 103 was obtained for the ratio of the specific rate constants. Since kinetia results suggested adsorption of the oxalate anions at the active growth centers of calcite, we further (18)Danes, C. W.;Nanmh, G. H.%urns. F o r d a y Soc. 1955.51.818. (19) Christofferasn, J.; Christoffemn, M. R.; Christensen, S. B.; NmcoUns, G. H.J. Cryat. Growth 1983.62.254,
In eq 7, C, is the equilibrium oxalate concentration, r is the amount of oxalate adsorbed per unit surface area of calcite, r, is the adsorption capacity, n is a constant usually greater than one, and k is a constant related to the free energy of adsorption, A(?*&, by k = exp(-AG,,/RT) (8) where T is the absolute temperature and R the gas constant. Adsorption was found to depend on the ionic strength, thus suggesting electrostatic interaction of the adsorbate with calcite. Further experimentation at higher pH values did not show any significant dependence of adsorption on pH. This is in agreement with extensive studies of the adsorption of various anions on a number of substrates, including calcite, which have shown that adsorption of anions shows a maximum at the point of their complete dissociation.20 The pH in our experiments had by far exceeded that corresponding to the complete dissociation of oxalate. In Fieure 8. the infrared sDectra I
(20) Parfitt,R. L.Ad". Agron. 1978,30, 1.
Giannimaras and Koutsoukos
860 Langmuir, Vol. 4, No. 4, 1988
I r /rnol.rn-Z
U
a
PH
-050-
\
-100b-1%-
-200
-
pigure 9. Electrophoretic ability of calcite particles: (0) bare; ( 0 )calcite covered with 11.75 pmol of oxalate m-2; (A)calcite covered with 11.75 mmol of oxalate m-2. el~ewhere.~ The electrokinetic results of both bare and oxalate-covered calcite particles are shown in Figure 9. The potential-determining ions of calcite are Ca2+and C032-, while pH changes would affect the speciation equilibria. The charging mechanism in the presence of oxalate can be considered as the result of the following equilibria: Ca2++ HC0,-
S=
CaHC03+
Ca2++ OH- -+CaOH+ lom4
10-6
Figure 7. Adsorption isotherms of oxalate anions on calcite at 25 O C : (0) 0.01 mol dm-3; (0) 0.1 mol dm-3.
+
Ca2+ C032-S= CaC02
+
+ OHH2C03+ OH-
C032- H 2 0 ; i HC03HCOC + H 2 0 =+
Ca2++ HC204- + CaHC204+ Ca2++ 2HC204-+ Ca(HC204)2 Ca2++ C z O z - * CaC2040 Ca2++ 2C2042-+ Ca(C204)22H2C2O4 + OH- + HC204-+ H 2 0 HC204-+ OH- ==C204'-
4obO
1
I
3500
3000
2500
I
l
I
20to
1803
1600
I
l
I
I
14W 1200 1000 800
WAVENUMBER
(
CM-'
1
600
I
400
200
>
Figure 8. Infrared spectra of (A) calcite and calcium carbonate grown on calcite seed crystals in the absence (B) and presence (C) of oxalate anions. of calcite and of calcite solid samples following adsorption are shown. The new bands at 1600 and 1300 cm-', which appeared after adsorption, belong to the oxalate present on the surface of calcite. Examination of the X-ray diffraction spectra of calcite samples following adsorption confirmed the absence of any coprecipitated calcium oxalate phase as well. The electrophoretic mobilities of oxalate-covered calcite particles were converted into electrokinetic charges by the procedure of Wiersema et a1.21taking into account relaxation effects. In the present case, however, these phenomena are not important for reasons described in detail (21) Wiersema, P. H.; Loeb, A. L.; Overbeek, J. Th.G. J. Colloid Interface Sci. 1966, 22, 78.
+ H20
Adsorption of oxalate would result in reduction of the Ca2+ sites on the surface of calcite, rendering them positive (CaHC204+),neutral (CaC202),or negative (Ca(C2O4)2-). Under our experimental conditions, where the C2042species is predominant, it is expected that adsorption of oxalate would render the calcite surface negative. Indeed, as may be seen in Figure 9, this effect was very pronounced, even reversing the electrokinetic charge at lower pH values. It is interesting to note that the relative amount of oxalate adsorbed on the calcite surface did not have any significant effect either on the sign or on the magnitude of the electrokinetic charge of calcite, as was the case with phosphate adsorption on calcite.'
Conclusions The crystallization of calcite on synthetic calcite seed crystals proceeds without the formation of any precursor phase under conditions of sustained supersaturation. The apparent order of 2 for the crystallization reaction is consistent with a surface-controlled process. The presence of oxalate in the supersaturated solutions favored the formation of calcium carbonate monohydrate as a precursor phase, converting to the thermodynamically more stable calcite. The verification of the formation of this precursor phase has been possible only by maintaining the
Langmuir 1988,4,861-867
86 1
initial supersaturation. Oxalate ions strongly adsorb onto calcite, their adsorption conforming with a Freundlich isotherm. As a consequence of their adsorption, oxalate ions render the calcite surface negative, causing an electrokinetic charge reversal at low pH values.
TGA analyses, Peter Busch of the State University of New York at Buffalo, and Lambros Kombotiatis for the scanning electron microscopy. A grant from the European Communities (NREN3G-0036 GR) in support of this work is gratefully acknowledged.
Acknowledgment. We would like to express our thanks to Anthony Margaritis for assistance with DSC and
Registry No. CaC03,471-34-1;CaC03.H20,32825-96-0; C204%, 338-70-5; calcite, 13397-26-7.
Light-Initiated Surface Modification of Oxide Semiconductors with Organic Dyes M . A. Ryan Department of Chemistry, Mount Holyoke College, South Hadley, Massachusetts 01075
M. T. Spitler* The Polaroid Corporation, 103 Fourth Avenue, Waltham, Massachusetts 02254 Received September 21, 1987. In Final Form: March 7, 1988 Photoaffinity labels were used as reagents for a light-initiated modification of oxide semiconductor surfaces with rhodamine dyes. Illuminating a dye-label complex with UV light resulted in a reactive nitrene which formed an N-0 bond between dye and semiconductor stable for up to 1 year. Intrinsic excitation of the semiconductor was also found to bond rhodamine and xanthene dyes through an ester bridge with the carboxyl functions of the dyes. It was demonstrated that dyes could be attached to electrode surfaces in patterns of 2 5 - ~ mstripes. Work in chemical modification of solid surfaces has demonstrated that extensive control may be exercised over the characteristic reactivity of metal and semiconductor electrodes.1-12 Electrodes have been modified to improve stability against corrosion, to extend the range of their chemical reactivity, and to improve their catalytic prope r t i e ~ . ~Chemical J reagents have been attached to electrode surfaces in order to accelerate interfacial redox reactions and to control diffusion of species to electrode surface^.^ In addition, chemically modified electrode surfaces have been widely used in light-assisted redox reaction~~5 and in photoelectrochemical and electrochromic devi~es.~.' Sensitizing dyes form an important class of molecules used as surface modification agents, primarily at semiconductor electrodes, where a photocurrent can be produced through oxidation of the excited dye. They have (1)Murray, R. W. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Decker: New York, 1984;Vol. 13,p 191. (2)Chidsey, C. E.D., Murray, R. W. Science (Washington, 00 1986, 231, 25. (3)Wrighton, M. W. Science (Washington, DC) 1986,231, 32. (4)Abruna, H. D.;Calvert, J. M.; Denisevich, P.; Ellis, C. D.; Meyer, T. J.; Murphy, W. R.; Murray, R. W.; Walsh, J. L. ACS Symp. Ser. 1982, 192. 133. ---. (5) Murray, R . . Acc. Chem. Res. 1980,13,135.
(6)Reichman, B.;Fan, F. F.; Bard, A. J. J. Electrochem. Soc. 1980, 127,333. (7)Viehbeck, A.; DeBerry, D. J. Electrochem. SOC.1985,132,1369. (8) Moses, P.R.; Murray, R. W. J. Am. Chem. SOC.1976,98,7435. (9)H a m , D.;Armstrong, N. R. J. Phys. Chem. 1978,2,1288. (10)Fox, M. A.; Nobs, F. J.; Voynick, T. A.; J. Am. Chem. SOC.1980, 102,4036. (11)Fujihira, M.; Ohishi, N.; Osa, T. Nature (London) 1977,268,226. (12)Fujihira, M.; Kubota, T.;Osa, T. J.Ekctroanul. Chem. 1981,119, 379.
0743-7463/88/2404-0861$01.50/0
been attached to wide band gap metal oxides such as Ti02, Sn02,and SrTi03.8-12Early work in modification of these electrode surfaces with organic sensitizing dyes generally involved attaching an alkylsilane layer to the surface in a silyl ether linkage and then attaching the photosensitizer to the silyl layer by way of an alkyl chain. The oxidation of the excited dye leading to photocurrent under these conditions was, however, significantly lower than the photocurrent produced with dye adsorbed from an aqueous solution. Fujihira et a1."J2 attributed this difference to the increased distance of the attached dye molecule from the semiconductorsurface and to the inhibition of electron transfer by the nonconductive,linking layer on the surface. Other authors have also discussed the extent to which the silyl layer hinders electron transfer to the underlying semiconductor.1° In an effort to improve the sensitization efficiency of the attached dye, Fujihira et al.ll developed a method of attaching rhodamine B directly to the semiconductor electrode surface through condensation of the carboxyl function on the dye with the hydroxylated semiconductor surface. With this ester linkage the sensitized photocurrent was found to be equal to that of adsorbed dye. A similar approach has been taken in the attachment of an inorganic sensitizer, ruthenium tris(bipyridyl), to metal oxide semiconductor electrode surfaces.13 However, esters are subject to hydrolysis in aqueous solution, and this ester linkage is hydrolyzed within hours of its formation. In this work we present a new method for electrode modification that leads to a stronger link between surface (13)Anderson, S.;Constable, E.; Dare-Edwards, M.; Goodenough, J.; Hamnett, A.; Seddon, K.; Wright, R. Nature (London) 1971,280,571.
0 1988 American Chemical Society
