Langmuir 1997, 13, 4377-4382
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Kinetics of 18-Carbon Carboxylate Adsorption at the Fluorite Surface Michael L. Free† and J. D. Miller* Department of Metallurgical Engineering, University of Utah, Salt Lake City, Utah 84112 Received April 8, 1996. In Final Form: May 12, 1997X In this study the kinetics of oleate and linoleate adsorption at the fluorite surface have been examined, in-situ, by Fourier transform infrared internal reflection spectroscopy. Data obtained in this study show that the adsorption process is not controlled by convective diffusion. Other results from this study indicate that oleate adsorbs to approximately the same monolayer coverage level when given sufficient time over a variety of temperature and solution conditions. This study also shows that competitive anions such as OH-, CO32-, and F- can severely restrict the rate of oleate adsorption; furthermore, proper control of such species will result in maximum oleate adsorption kinetics.
Introduction Most hydrofluoric acid is derived from the mineral fluorite, CaF2, which is obtained by mining and subsequent froth flotation to separate fluorite particles from unwanted gangue minerals. Froth flotation of the mineral fluorite is made possible by the adsorption of carboxylate surfactants. Despite the importance of understanding the kinetics of carboxylate adsorption at the fluorite surface, few such studies have been performed. This research project was undertaken in order to provide a more complete analysis of carboxylate adsorption kinetics at the fluorite surface. The authors are aware of only two kinetic studies involving the adsorption of unsaturated 18-carbon carboxylates at the fluorite surface. One ex-situ study utilized IR transmission analysis of powdered fluorite dispersed in potassium bromide pellets to evaluate the adsorption kinetics of oleate from aqueous solutions.1 In the other study adsorption kinetics were inferred from the rate of fluorite dissolution, which decreased with increased oleate loading.2 Both studies have arrived at conclusions on the basis of substantial extrapolation and a limited amount of data. Also, the experimental procedures followed in these tests may have allowed significant amounts of calcium dioleate and/or oleic acid to form and adsorb, thus making the interpretation of the data and conclusions less certain. A study by Peck1 suggested that significant additions of F- (0.0625 M) as NaF resulted in higher fractions of physically adsorbed oleate than chemisorbed oleate. Other ions such as SO42- (added as Na2SO4) and CO32- (added as Na2CO3) did not result in high fractions of physically adsorbed oleate. The effect of ions such as F-, SO42-, and CO32- on oleate adsorption kinetics, however, was not examined by Peck.1 In the study of oleate adsorption at the fluorite surface at 18 °C and pH 5.8 by Sobieraj,2 no data are provided for adsorption times less than 300 min. For the data provided by Sobieraj, adsorption was approximately 75% complete before the first adsorption level was recorded. Also, the study by Sobieraj provides no kinetic model for adsorption. Instead, it provides the calculated coverage versus time † E-mail:
[email protected]. Telephone: 801-585-9798. * Author to whom correspondence should be addressed. X Abstract published in Advance ACS Abstracts, July 15, 1997.
(1) Peck, A. S. Ph.D. Dissertation, University of Utah, 1962. (2) Sobieraj, S. Royal Australian Chemistry Institute, 8th National Convention, Division of Surface Chemistry, University of New South Wales, August 24-28, 1987.
S0743-7463(96)00329-0 CCC: $14.00
Figure 1. Schematic diagram of the ATR or FTIR/IRS accessory used in this study.
on the basis of the rate of fluorite dissolution. Because no data is provided at relatively low coverages, it is difficult to fit Sobieraj’s data to any kinetic model. Consequently, the acquisition of data at low levels of coverages is needed in order to adequately evaluate oleate adsorption kinetics. In addition, the low pH value of 5.8 at which the experiments were done suggests that oleic acid formation should occur,3 further complicating the kinetic analysis. The lack of complete kinetic data with respect to both adsorption data, as well as solution speciation, prompted this study. Experimental Procedures All in-situ Fourier transform infrared internal reflection spectroscoy (FTIR/IRS) measurements were performed using a FTIR/IRS accessory obtained from Harrick Scientific Company shown in Figure 1 with deuterium oxide (99.9% pure) from Cambridge Isotope Laboratories (CIL). Sodium oleate (>99% purity) and sodium linoleate (>99% purity) were used as received from Sigma Chemical Co. All pH adjustments were made using DCl and NaOD (>99% pure), also from CIL. Fluorite internal reflection elements (IREs) from Optovac Inc. with parallelpiped geometry, dimensions (mm) of 52 × 9.5 × 1.68 and 52 × 10.0 × 2.03, and 73° ends were used in this study. IREs were cleaned by manual repolishing using finger cots from Thomas Scientific Co. with 0.05 µm γ-alumina polishing powder on Texmet polishing pads produced by Buehler. Spectra from repolished IREs were compared to plasma-cleaned IREs to determine the extent of organic contamination. These comparisons revealed that less than 5% of a monolayer of residual hydrocarbon contamination remained on the polished IRE (3) Free, M. L.; Miller, J. D. Int. J. Miner. Process. 1996, 48, 197.
© 1997 American Chemical Society
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surfaces relative to the plasma-cleaned IRE surfaces as determined by FTIR/IRS measurements. Glassware was cleaned by washing in 5 M KOH, rinsing in 18 MΩ‚cm deionized water, washing in chromic acid, and then rinsing in 18 MΩ‚cm deionized water. All spectra were obtained using a BioRad Digilab FTS-40 spectrometer with a liquid-nitrogen-cooled wide-band mercurycadmium-telluride detector (1024 scans, 8 cm-1 resolution). In each case tests were conducted using a multiple internal reflection FTIR/IRS edge-seal liquid cell obtained from Harrick Scientific Co. A rectangular foil aperture was utilized in front of the IRE to limit the incident infrared beam height to that of the sample chamber height, thus preventing portions of the incident beam from reaching the detector without having encountered the sample surface at which adsorption could occur. Significant errors usually result without control of the beam height due to improper application of Beer’s Law.4,5 Solutions were injected into the FTIR/IRS cell with an approximate volume of 1 mL using an Orion Model 352 syringe pump and a 100 mL glass syringe made by Popper and Sons Inc. Initially, 5 mL of D2O was injected and allowed to equilibrate with the IRE for 1 h before an absorbance spectrum of D2O versus the clean crystal background was taken. After the D2O spectrum was taken, 10 mL of D2O was injected at a rate of 1 mL/s to remove as much dissolved calcium as possible. Immediately following the injection of D2O, 10 mL of oleate solution was injected at a rate of 1 mL/s, followed by 10 mL/h for the first 2 h of adsorption and then 2 mL/h afterward until the end of the test. The D2O absorbance spectra were subtracted from solution absorbance spectra to produce the final spectra. All in-situ adsorption density calculations were performed using eq 1, which was initially derived and used by Sperline et al.6 and later modified for the parallelpiped geometry and applied to different systems by Miller and Kellar.7 In eq 1, A is the
Γ(µmol/m2) )
107(A - CbdeN) 2Nde dp
( )
(1)
integrated absorbance (cm-1), Cb is the bulk solution concentration (mol/L), is the molar absorptivity (L/mol‚cm2), de is the effective depth (cm), N is the number of internal reflections, and dp is the depth of penetration (cm). Additional information regarding the definitions of these terms and independent confirmation of the adsorption density equation is provided elsewhere in the study of transferred LangmuirBlodgett films by Jang and Miller.8 The values used in eq 1 were the same as those used by Miller and Kellar7 with the following exceptions that are believed to be more appropriate for this study: (1) refractive index of CaF2, 1.415;9,10 (2) refractive index of D2O, 1.24;11,12 (3) measured fraction of perpendicularly polarized light, 0.43;13 (4) measured fraction of parallel polarized light, 0.57;13 (5) measured oleate molar absorptivity, 39 500 (L/mol‚cm2) for a spectral integration range of 3030 to 2820 cm-1.13
Results and Discussion Mass Transport Effects. Because only two previous studies regarding the kinetics of oleate adsorption at a (4) Free, M. L.; Miller, J. D. Applied Spectroscopy, 1994, 48, 891. (5) Free, M. L.; Jang, W. H.; Miller, J. D. Colloids Surf. A 1994, 93, 127. (6) Sperline, R. P.; Muralidharan, S.; Freiser, H. Langmuir 1987, 3, 198. (7) Miller, J. D.; Kellar, J. J. In Challenges in Mineral Processing; Sastry, K. V., Fuerstenau, M. C., Eds.; SME: Littleton, 1989; p 109. (8) Jang, W. H.; Miller, J. D. Langmuir 1993, 9, 3159. (9) Browder, J. S.; Ballard, S. S.; Klocek, P. In Handbook of Infrared Optical Materials; Klocek, P., Ed.; Marcel Decker: New York, 1991; p 193. (10) Feldman, A.; Horowitz, D.; Waxler, R. M.; Dodge, M. J. NBS Technical Note 993; U.S. Department of Commerce: Springfield, VA, 1979. (11) Sethna, P. P.; Palmer, K. F.; Williams, D. J. Opt. Soc. Am. 1978, 66, 815. (12) Bertie, J. E.; Ahmed, M. K.; and Eysel, H. H. J. Phys. Chem. 1989, 93, 2210. (13) Free, M. L. Ph.D. Dissertation, University of Utah, 1994.
fluorite surface have been conducted (neither of which considered diffusion), it was felt that this study should begin with an examination of the effect of diffusion in 18-carbon carboxylate adsorption kinetics. The effect of mass transport in this study was evaluated using the equation derived by Levich14 for convective diffusion to the surface of a flat plate through a solution flowing over the flat plate under laminar flow conditions as given in eq 2 below:14,15
J)
0.34D(Cb - Cs)xU ν D xνb
()
1/3
(2)
where J is the diffusional flux (mol/cm2‚s), D is the diffusivity (cm2/s), Cb is the bulk solution concentration (mol/L), Cs is the solution concentration at the adsorbing surface (mol/L), U is the velocity of fluid upstream from the leading edge of the plate (cm/s), ν is the kinematic viscosity (cm2/s), and b is the distance from the leading edge of the plate to the point of measurement (cm). The flat plate fluid flow geometry was selected to best represent the geometry of the IREs used in the attenuated total reflection (ATR) solution contact accessory. Describing the solution chamber with a characteristic dimension (width and depth) of 3 mm, the Reynold’s number at each of the flow velocities selected for this study was calculated to be less than 15; thus the laminar flow conditions that are necessary for proper application of eq 2 were achieved. According to eq 2, the diffusional flux should be proportional to the square root of the flow velocity for diffusion-controlled adsorption. To evaluate the effect of mass transport in this study, the dependency of oleate flux on flow velocity was examined by performing three experiments with 1 × 10-5 M sodium oleate in D2O at pH 9.4 and 25 °Cseach with a different flow velocity. The selected experimental flow velocities were 10, 31, and 91 mL/h. Because the diffusional flux is directly proportional to the square root of the flow velocity in eq 2, the oleate flux should increase by approximately a factor of 3 over this range of selected flow rates if the adsorption reaction is diffusion-controlled. In order to determine the effect of flow velocity on the rate of adsorption, the level of adsorption as a function of time was analyzed in the three flow velocity experiments using a linear regression of adsorption density (µmol/m2) versus time (min) over the first 15 min of adsorption. The slopes of the lines, which correlate with the rate of adsorption, were 0.190, 0.192, and 0.194 µmol/m2‚min for the 10, 31, and 91 mL/h flow rates, respectively. All R2 values for the regressions were greater than 0.96. These results, which show that the rate of adsorption is independent of flow rate, suggest that mass transport did not play a major role in the kinetics of carboxylate adsorption under the conditions analyzed. In addition, using an estimated ionic diffusivity value of 1 × 10-5 cm2/s16,17 for oleate (based upon values for other carboxylic acids as well as an estimated value for large molecules), an average measurement of 2.0 cm as the distance from the leading edge of the plate, a concentration difference of 1 × 10-5 mol/L, and a fluid (14) Levich, V. G. Physicochemical Hydrodynamics, Prentice-Hall, Inc.: Englewood Cliffs, 1962. (15) Strickland, P. H.; Lawson, F. International Symposium on Hydrometallurgy; Evans, D. J. I.; Shoemaker, R. S., Eds.; AIME: Boston, 1973. (16) Perry, R. H., Ed. Perry’s Chemical Engineers Handbook, 6th ed.; McGraw-Hill Book Co.: New York, 1984. (17) Perry, R. H.; Chilton, C. H., Eds. Chemical Engineer’s Handbook, 5th ed; McGraw-Hill Book Co.: New York, 1973.
Kinetics of 18-Carbon Carboxylate Adsorption
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flow velocity of 0.046 cm/s, the molar flux, J, in eq 2 for diffusion was calculated to be 4.7 × 10-13 mol/cm2‚h or 4.7 × 10-3 µmol/m2‚s, which is significantly higher than the typically measured initial adsorption rate of 3.3 × 10-3 µmol/m2‚s. Thus, the potential mass transport flux is greater than the observed flux, suggesting that the adsorption reaction is not mass-transport limited. The calculated value of 4.7 × 10-3 µmol/m2‚s, however, is only 42% higher than the measured value, suggesting that, at significantly lower oleate concentrations, mass transport may affect the rate of adsorption. The diffusional flux value is, nevertheless, significantly higher than the measured adsorption rate, despite the fact that it was determined at the lowest initial injection speed and a relatively low-concentration value. It should be recognized that this calculated value assumes an idealistic flow geometry that is slightly different from the actual one; thus this calculated value is only an approximation. The conclusion from these experimental results and the accompanying theoretical analysis is that oleate adsorption (research1 indicates that oleate, rather than sodium oleate, is adsorbed) is not diffusion-controlled in this study at oleate concentrations of 1 × 10-5 M or greater. Surface Reaction Considerations. One common type of kinetic model used to describe adsorption rates is the Langmuir model,18 where the adsorption reaction is described as
B + V S B‚V
(3)
where B is the adsorbing species, V is a vacant surface site, and B‚V is the adsorbed species at a surface site. The rate of adsorption for the reaction described in eq 3 can be expressed as
Rads ) kf[B][V] - kb[B‚V]
(4)
in which Rads represents the rate of adsorption, kf represents the forward reaction rate constant, and kb represents the backward reaction rate constant. From eq 4 it is clear that after adsorption reaches equilibrium the rate of adsorption drops to zero (i.e., the forward and backward reaction rates become equal). Consequently, the significance of the backward reaction rate can be determined by simply observing the extent of desorption after reducing the solution concentration of the adsorbing species by several orders of magnitude after equilibrium is reached. If very little desorption occurs, the backward reaction rate can be ignored, leading to
Rads ) kf[B][V]
(5)
To verify that the backward reaction rate is small in this study, several adsorption tests were conducted for 20 h in sodium oleate and sodium linoleate solutions (1 × 10-5 M) at pH 9.4 in D2O. The adsorption was followed by injecting 10 mL of pure D2O at pH 9.4 to remove most of the oleate or linoleate molecules. After the system was rinsed, it was allowed to equilibrate for approximately 1 h, and then the adsorption density was determined and compared to the adsorption density value prior to rinsing. The results are tabulated in Table 1 and show that no consistent desorption is observed. Most importantly, the rate of change in the adsorption density is a small fraction of the initial adsorption rate, which increases rapidly with increasing temperature. The overall average change in adsorption density was -1.4%, further suggesting that (18) Fogler, H. S. Elements of Chemical Reaction Engineering; Prentice-Hall, Inc.: Englewood Cliffs, 1986.
Table 1. Comparison of FTIR/IRS In-Situ Adsorption Dataa ads/rinse temp (°C) 25 37 50 50 25 37 50
compound
ads density before rinsing
ads density after rinsing
ads change (%)
oleate oleate oleate oleate linoleate linoleate linoleate
7.9 (µmol/m2) 8.5 (µmol/m2) 7.6 (µmol/m2) 7.5 (µmol/m2) 6.8 (µmol/m2) 7.2 (µmol/m2) 8.9 (µmol/m2)
8.1 (µmol/m2) 7.9 (µmol/m2) 7.9 (µmol/m2) 8.9 (µmol/m2) 6.6 (µmol/m2) 6.6 (µmol/m2) 7.4 (µmol/m2)
+2.5 -7.1 +4.0 +18.7 -2.9 -8.3 -16.9
a Before and after rinsing for 1 h with D O at pH 9.4 after 20 h 2 of prior oleate adsorption at the given temperature.
the rate of desorption is slow and within the general experimental error of 5%. Further evidence for a limited backward reaction can be inferred from Table 1 by comparing the level of adsorption which on average is maintained at a relatively constant value over a significant change in temperature. Because little change in the average adsorption density is observed as the temperature rises, it is evident that the adsorbate is bonded to the surface with substantial bond strength, making significant desorption unlikely in the temperature range studied. This implies that the species are more likely chemisorbed than physically adsorbed and confirms previous findings that show oleate predominantly chemisorbs at the fluorite surface unless significant amounts of calcium are available for calcium dioleate formation.3 Because the change in adsorption density after injecting D2O without surfactant is small, it is difficult to measure precisely the degree of desorption; therefore, accurate analysis of the free energy of adsorption, which implicitly relies upon accurate forward and backward reaction rate constants or true equilibrium constants, was not possible. Only the forward reaction rate constant could be accurately obtained. Because the backward reaction rate is very small in this study, it is neglected in the present adsorption rate analysis; therefore, eq 5 is used to express the rate of adsorption. For a first-order reaction described by eq 5, it can be shown that
dθ ) kf[B][1 - θ] dt
(6)
where, θ represents the fractional coverage, kf represents the forward reaction rate constant, t represents reaction time, and [B] represents the adsorbing species concentration. Integration of eq 6 leads to
(1 -1 θ) ) k [B]t
ln
f
(7)
assuming no coverage at time zero. Thus, for constant solution concentration, plotting ln[1/(1 - θ)] versus time should yield a slope of kf[B] and a zero intercept if the kinetics follow eq 7, assuming a first-order reaction with respect to the adsorption density at constant bulk concentration. Using spectral data from in-situ adsorption experiments, such as those presented in Figure 2, the resulting adsorption densities were determined. Figure 3 illustrates the resulting plot of ln[1/(1 - θ)] versus time for the adsorption of 1 × 10-5 M oleate in which the maximum monolayer level adsorption density was selected between the realistic monolayer coverage values of 5.8 and 6.8 µmol/m2 (on the basis of Langmuir-Blodgett results8) such that the best fit of the data was obtained as determined by maximizing the R2 values. As predicted
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Figure 2. In-situ FTIR/IRS spectra of oleate adsorption at the fluorite surface as a function of adsorption time as indicated. Adsorption occurred at 25 °C in D2O with 1 × 10-5 M oleate. (Note that spectral base lines were altered for presentation purposes.)
Free and Miller
Figure 4. Plot of experimental data points together with the predicted response represented by the solid line as determined using eq 7 and the original data given in Figure 3. Table 2. Kinetic Data from in-Situ FTIR/IRS Adsorption Experimentsa temp slope kf (°C) compound (L/min) (L‚min/mol) 16.0 25.0 37.0 50.0 50.0 16.5 25.0 50.0 a
oleate oleate oleate oleate oleate linoleate linoleate linoleate
0.0189 0.0470 0.0711 0.1295 0.1437 0.0308 0.0515 0.1186
1890 4700 7110 12 950 14 370 3080 5150 11 860
R2 0.967 0.993 0.955 0.989 0.974 0.968 0.972 0.980
θmax time (µmol/m2) (min) 5.91 6.45 6.80 5.95 6.39 6.80 6.80 6.27
210 53 64 29 29 61 50 29
Solution concentration was 1 × 10-5 M and the pH was 9.4.
Activation Energy. The adsorption activation energy can be determined using the Arrhenius equation, which can be expressed as:18 Figure 3. Plot of ln[1/(1 - θ)] versus time for oleate adsorption at a fluorite surface as measured from in-situ FTIR/IRS at 25 °C. The oleate concentration was 1 × 10-5 M. The maximum adsorption density was set at 6.45, and the linear regression was performed using all data points between 0 and 53 min.
by eq 7, a reasonably linear relationship exists between ln[1/(1 - θ)] and time for the initial rate data obtained in this study as illustrated in Figure 3. Also, Figure 3 shows that the linear relationship is accompanied by an intercept near zero, though it should be noted that for the regression analysis the line was forced through zero. The data indicate first-order kinetics are suitable at a concentration of 1 × 10-5 M oleate. Using the slope of the linear regression, the value of kf[B] was determined and used to predict adsorption density versus time for the purpose of comparing predicted values with experimental values. Figure 4 shows how the experimental data compares with the predicted adsorption density versus time on the basis of eq 5, which was rearranged to a more useful form shown in eq 7. The predicted values in Figure 4 were calculated using the slope and maximum adsorption density values determined from the regression analysis of the data presented in Figure 3. From Figure 4 it is apparent that eq 5 gives a reasonable mathematical description of what is experimentally observed. Results from additional experiments and regression analyses for oleate and linoleate adsorption over a range of temperatures are given in Table 2. Because eq 5 is simple and fits the data so well, it was selected as the kinetic model for oleate adsorption at the fluorite surface for short adsorption times.
kf ) QeEa/RT
(8)
where Q is the Arrhenius rate coefficient, Ea is the adsorption activation energy, R is the gas constant, and T is the absolute temperature. Then, Equation 8 can be rearranged to give
ln kf ) ln Q -
Ea RT
(9)
Using eq 9, Ea and Q can be determined from the slope and intercept of a plot of ln(kf) versus 1/T. Figures 5 and 6 show plots of ln(kf) versus 1/T for oleate and linoleate adsorption, respectively, using the data in Table 2. Results from a linear regression of the data in Figures 5 and 6 yield slopes of -5070 and -3650 K, respectively, and intercepts of exp(25.2) and exp(20.7) L‚min/mol, respectively. From the slopes and intercepts the resulting values of the activation energies and rate constants were calculated as 42.1 kJ/mol and 8.79 × 1010 L‚min/mol, respectively for oleate, and 30.4 kJ/mol and 9.77 × 108 L‚min/mol, respectively for linoleate. The oleate adsorption activation energy of 42.1 kJ/mol is significantly higher than the value of 16.3 kJ/mol determined by Peck for oleate on fluorite,1 but it is much closer to the value of 31.3 kJ/ mol determined by Huang and Miller19 for the adsorption of amyl xanthate at pyrite and marcasite surfaces by an electrochemical surface reaction. These high experimental activation energy values indicate these adsorption reac(19) Huang, H. H.; Miller, J. D. Int. J. Miner. Process. 1978, 5, 241.
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kf
Kinetics of 18-Carbon Carboxylate Adsorption
Figure 7. Effect of deuterate ions on adsorption kinetics as revealed by in-situ FTIR/IRS. Each experiment was conducted with 1 × 10-5 M sodium oleate in deuterium oxide at 25 °C. Note that pD ) pH + 0.4.21
kf
Figure 5. Arrhenius plot of oleate data in Table 2.
Figure 6. Arrhenius plot of linoleate data in Table 2. Note that one data point at 37 °C was omitted because it was a statistical outlier due to the poor R2 value of 0.61 obtained with it present.
tions are difficult to reverse, suggesting that chemisorption occurs rather than physical adsorption which, due to lower activation energies, is more easily reversed. These values, together with the transport analysis, indicate that the adsorption process is controlled by a surface reaction, rather than by a transport process. Effect of Competing Anions on Adsorption Kinetics. Some researchers have found that competing anions affect the extent of adsorption as well as the nature of the adsorbate when anionic surfactants such as oleate are used.19,20 In this study the effect of deuteroxide/hydroxide, carbonate, and fluoride ions on oleate adsorption was examined. The effect of the deuteroxide ions on adsorption kinetics was evaluated by performing three in-situ adsorption testsseach at different pH levels (It should be noted here that all in-situ experiments were performed in D2O where the pH adjustments were made with NaOD and where pD ) pH + 0.4).21 Data from the three in-situ adsorption tests are presented in Figure 7. Figure 7 shows that the rate of adsorption drops as the pH is increased above pH 9.5. Measurements of fluorite in water show that at pH 9.5 (using NaOH for pH adjustment) the fluorite surface has a +16 ( 2 mV zeta potential, whereas at pH 10.9 the zeta potential changes to -10 ( 2 mV. These zeta (20) Fuerstenau, M. C.; Miller, J. D.; Kuhn, M. C. Chemistry of Flotation; SME/AIME: New York, 1985. (21) Glascoe, P. K.; Long, F. A. J. Phys. Chem. 1960, 64, 188.
Figure 8. Effect of NaF on oleate adsorption kinetics as revealed by in-situ FTIR/IRS analysis. Each experiment was conducted at pH 9.4 (pD 9.8) with 1 × 10-5 M sodium oleate in deuterium oxide at 25 °C.
potential measurements clearly indicate that hydroxide and deuteroxide ions have a strong tendency to adsorb on the fluorite surface and inhibit the adsorption of oleate. The effect of fluoride ions on oleate adsorption kinetics was examined by adding different amounts of NaF to oleate solutions and observing the rate of oleate adsorption. Figure 8 shows that increasing the concentration of NaF decreases the rate of oleate adsorption provided that the NaF concentration is relatively high (>0.01 M) as indicated in Figure 8. In the presence of 0.01 M NaF, the rate of oleate adsorption is 3.3 × 10-3 µmol/m2‚s, whereas with 0.03 M NaF present the rate drops to 3.6 × 10-4 µmol/ m2‚s. The increasing inhibition to oleate adsorption due to enhanced fluoride ion adsorption at higher fluoride concentrations is corroborated by zeta potential measurements on fluorite (without oleate) which show a change from -4 ( 2 mV to -8 ( 2 mV for the corresponding change from 0.01 to 0.03 M NaF, respectively. Increasing the concentration of Na2CO3 in the fluorite/ oleate system also has a pronounced effect, inhibiting oleate adsorption as illustrated in Figure 9. In comparing Figures 8 and 9, it can be seen that 5 × 10-3 M Na2CO3 has a much greater inhibitory impact on the rate of oleate adsorption than 1 × 10-2 M NaFseven though the concentration of sodium is the same. This comparison suggests that the CO3-2 ion has a higher tendency to adsorb
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Figure 9. Effect of Na2CO3 on oleate adsorption kinetics as revealed by in-situ FTIR/IRS analysis. In each test the sodium oleate concentration was 1 × 10-5 M, the pH was 9.4 (pD 9.8), and the temperature was maintained at 25 °C.
Figure 10. Comparison of adsorption density versus time for long oleate adsorption times in the presence of competing anions. All adsorption experiments were conducted at 25 °C with 1 × 10-5 M sodium oleate and the resulting adsorption density values determined from in-situ FTIR/IRS spectra.
at the fluorite surface than fluoride ions. This evidence is supported by the zeta potential of fluorite of -21 ( 3 mV measured without oleate in 0.005 M carbonate solution and -18 ( 2 mV in 0.01 M carbonate solution at pH 9.5. The more negative zeta potential obtained with the lower carbonate concentration may explain why the lower carbonate concentration yielded the greatest inhibition to adsorption. However, it should also be noted that carbonate may also react with dissolved calcium near the fluorite surface and thereby alter the surface properties. A thorough comparison between the different inhibitory effects of competetive ions is difficult to perform, due to the fact that the adsorption studies were performed in deuterium oxide solution, whereas the zeta potential measurements were performed in water. However, despite the difficulty in the comparison, the results indicate clearly that the inhibitory effect is due primarily to the individual ion tendency to adsorb on the fluorite surface. These comparative results are important from an industrial standpoint because Na2CO3 is usually added to fluorite separation processes in relatively high concentrations to increase the pH of the suspension. Reduction of competitive anions such as carbonate in industrial fluorite separations may increase the rate of oleate adsorption, thereby reducing the conditioning time. Additional insight into the effect of competitive anions on oleate adsorption can be obtained by analyzing the adsorption density at longer adsorption times. Figure 10 shows that even though the rate of adsorption is significantly reduced by the addition of competitive anions, the level of adsorption tends to approach the monolayer level of around 6.2 µmol/m2 if given sufficient time. This observation is important because if the addition of competitive anions resulted in an increase in the backward reaction rate without affecting the forward reaction rate (see eq 4), the adsorption density would have reached a plateau at a significantly lower value than that was determined without the presence of the competitive anion.
The fact that the adsorption density appears to reach approximately the same equilibrium value (5.8-6.8 µmol/ m2) with or without significant amounts of competitive anions suggests that the presence of the inhibitive or competitive anions decreases the forward reaction rate constant. Also, the results in this section, which show that oleate adsorbs to nearly the same equilibrium value despite competitive ions, further substantiate the observation that oleate is strongly chemisorbed at the fluorite surface. This conclusion is corroborated by previous spectroscopic evaluation of oleate monolayers using ultrasonic cleaning in organic solvent which showed it is difficult to remove a self-assembled oleate monolayer, yet physisorbed layers of adsorbed calcium dioleate are easily removed.3 Conclusions In this study it has been shown that the rate of oleate and linoleate adsorption is not limited by mass transport or convective diffusion, but rather is controlled by a reaction at the fluorite surface as evidenced by the activation energies of oleate and linoleate adsorption of 42.1 and 30.4 kJ/mol, respectively. The adsorption is believed to be first order with respect to the availability of surface sites when the solution surfactant concentration is maintained at a constant level. It has also been shown that inhibitive or competitive anions such as OD-/OH-, CO32-, and F- strongly influence the rate of oleate adsorption. In addition, this study shows that for the conditions evaluated oleate chemisorbs at the fluorite surface and is not desorbed to a significant extent by the presence of reasonable quantities of competitive anions. Acknowledgment. The authors wish to acknowledge funding for this project from the DOE Basic Sciences Grant No. DE-FG-03-93ER14315. LA960329R
