Langmuir 1985,1, 201-206 Andersonz2applied the Langmuir-type isotherm to adsorption of arsenate ion on amorphous aluminum hydroxide at the isoelectric point, by considering the total concentration of the adsorbate, yet this analysis was designated as strictly empirical? The present work s h o h that there is justification in using such an isotherm in the treatment of adsorption data as long as the adsorption (22) Anderson, M. A.; Ferguson, J. F.; Gavis, J. J. Colloid Interface Sci. 1976, 54, 391.
201
potential and the pH are constant (eq 17 and 18). In Anderson’s case the iep varies over a rather narrow pH range, which explains his ability to fit the results as he did. Further refinement in the analysis employed in this work would require a precise knowledge of the role protons and hydroxide ions play in the overall adsorption process. Acknowledgment. We are indebted to Roberto Torres for some electrophoretic measurements. Registry No. Oxalic acid, 144-62-7; citric acid, 77-92-9;hematite, 1317-60-8.
Interactions of Metal Hydrous Oxides with Chelating Agents. 7. Hematite-Oxalic Acid and -Citric Acid Systems? Yuting Zhang,S Nikola Kallay,* and Egon MatijeviC* Department of Chemistry and Institute of Colloid and Surface Science, Clarkson University, Potsdam, New York 13676 Received July 18, 1984. In Final Form: October 16, 1984 Adsorption of oxalate and citrate ions on colloidal spherical hematite particles and their effect on release of iron from the latter have been studied as a function of solute concentration, pH, and temperature. In principle, acidification enhanced both the adsorption and the dissolution until the pH was sufficiently lowered. Oxalic acid influenced much more strongly these processes than did citric acid. It could be shown that the iron release was dependent on the nature of surface complexes and their adsorption density. Protons affected the kinetics of dissolution only through the control of the adsorption of the two acids. Introduction Oxalic and citric acids are frequently used additives for the removal of rust. Specifically, these acids are common decontamination agents (alone or in combination with other complexing species) for water cooled nuclear power p1anta.l The primary purpose in the latter application is to dissolve the accumulated “crud”, which is made up mainly of hematite, magnetite, and different nickel and cobalt ferrites. This work deals with the interactions of oxalic and citric acids with a colloidal hematite dispersion consisting of spherical particles uniform in size and shape. Adsorption of the two acids on this iron oxide and the dissolution of the solid in their presence was studied as a function of pH, temperature, and the concentration of the reacting species. Since the adsorbent and the state of the solutes as a function of pH are well-defined, this system lends itself well for a theoretical treatment of the experimental results. The adsorption of these two acids on hematite is discussed in detail elsewhere.z An interpretation of the dissolution effects, based on the adsorption phenomena, is offered in this work. Such hematite dispersions were previously used to study the effects of EDTA and other chelating agents with the same purpose in mind.3y4 Experimental Section (A) Preparation of Spherical Hematite. Hydrosols consisting of spherical hematite (a-Fez03)particles of narrow size distribution were prepared and purified as described earlier.3v6 A hot solution (-85 “C) that was 0.018 M in FeC13and 0.001 M ‘Supported by NSF Grant CHE-83 18196. *On leave from Tianjin College of Textile Technology, Tianjin, PR.China. ‘Onleave from the University of Zagreb, Yugoslavia. 0743-7463/85/2401-0201$01.50/0
in HC1, contained in a 1-L Pyrex bottle, was aged in a forcedconvection oven with turbulent air circulation at 105 “C for 24 h. After cooling in cold water, the hydrosol was centrifuged, the supernatant solution discarded, and the precipitate resuspended with a HN03solution (pH 3) in an ultrasonic bath. The washing was repeated several dozen times until no presence of chloride or ferric ions could be detected in the supernatant liquid. The final dispersion (kept at pH 3) had a concentration of -100 mg of hematite per cm3as determined by dry weight analysis. The average particle diameter obtained from electron micrographs was 0.090 pm with the standard deviation of 0.010 pm. The specific surface area of hematite particles was determined by the BET multipointa method with a Monosorb Surface Analyzer (QuantachromeCorp.). The obtained value of 12.7 m2/g was in good agreement with the geometric specific surface area of 11.9 m2/g calculated from the average particle size and density (5.24 g/cm3) of hematite. (B)Techniques. (a) Electrophoresis. Electrophoretic mobilities were measured as a function of pH and the concentration of organic acid at the constant ionic strength (1 x mol dm-3 NaN03) with the Rank Brothers Mark I1 microelectrophoresis apparatus using a cylindricalthin-walled van Gils cell. (b) Adsorption and Dissolution Measurements. All measurementswere carried out in tightly stoppered Pyrex culture test tubes with Teflon-lined screw caps. Sodium hydroxide or nitric acid were used to adjust the pH while each system contained 1X mol dm-3 NaN03. The tubes with given amounts of a-Fe203and of the organic acid were shaken gently in a temperature-controlled bath. The adsorbent was then separated by centrifugation at 12000 rpm at the end of the predetermined reaction time. The supernatant solution,which still contained some a-FezOs particles, as detected by a Tyndall beam, was filtered through a 0.05-pm pore size (1)For a bibliography, see: Electric Power Research Institute, Interim Report EPRI NP-1033,March 1979. (2) Kallay, N.; MatijeviC, E. Langmuir,preceding paper in this issue. (3)Chang, H.-C.; Healy, T. W.; MatijeviE, E. J.Colloid Interface Sci. 1983,92,469. (4)Chang, H.-C.; MatijeviE, E. J.Colloid Interface Sci. 1983,92,479. (5)MatijeviE, E.; Scheiner, P.J. Colloid Interface Sci. 1978,63,509.
0 1985 American Chemical Society
Zhang, Kallay, and Matijeuic'
202 Langmuir, Vol. 1, No. 2, 1985 2
0 L.
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PH Figure 1. Electrophoretic mobilities of hematite (a-Fe203) particles (0.09 r m in diameter) in the absence and in the presence of oxalic acid (upper) and citric acid (lower),as a function of pH. Nuclepore membrane, its pH was measured, and then it was used for the analysis of solutes. (c) Analyses. The concentrations of oxalic and citric acids were determined from the decrease in the absorption at 492 nm caused by the addition of these acids to the iron(II1) 5-nitrosalicylate complex in acidified (pH 2.6-2.8) water/methanol solutions? It is important that the solution of the iron(II1) 5nitrosalicylate reagent was never stored for longer than 5 days. The calibration curves of the absorbances vs. concentrations of organic acid gave straight lines up to 5 X lo4 mol dm-3oxalic mol dm-3citric acid. acid and up to 1.4 X The technique for the determination of the concentration of ferric species was the same as described earliera3v7An excess of EDTA was added into the supernatant solution of a sol in order to replace the dissolved iron(II1) oxalate and iron(II1) citrate species by the Fel"EDTA complex, since the stability constant of this chelate is much larger than that of the former solutes.* A linear relationship was obtained between concentration of the ferric ion (up to 7 X mol dm-3) and (Azss- aASo6), where A is the absorbance at a given wavelength in nm and a is a constant.
Results (A) Electrophoretic Mobilities. The electrophoretic mobilities of hematite particles in the absence and in the presence of oxalic or citric acid as a function of pH at 28 "C are shown in Figure 1. The isoelectric point (iep) of pure hematite is at pH 7.6, in good agreement with the previously reported value? The complexing agents cause a significant shift of the iep to lower pH values which depends on the concentration of the added acids. It is noteworthy that the anions increase the negative mobility of hematite particles above their isoelectric point, indicating solute adsorption under these conditions. (B) Interactions of Hematite with Oxalic Acid. (a) Adsorption of Oxalic Acid on Hematite. Adsorption (6) Lee, K. S.; Lee, D.W.Anal. Chem. 1968,40, 2049. (7) Bhattacharyya, S. N.; Kundu, K. P. Talanta 1971, 18, 446. (8) Hegenauer, J.; Saltman, P.; Nace, G. Biochemistry 1979,18,3865. (9) Kuo, R. J.; MatijeviE, E. J. Colloid Interface Sci. 1980, 78, 408.
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EQUIL. CONC. OF OXALIC ACID CX103M) Figure 2. Adsorption isotherms of oxalic acid on spherical hematite particles for different pH values after 20 min of equilimol dme3NaN03. bration at 25 "C in the presence of 1 X measurements of oxalic acid on hematite at a constant pH showed that an equilibrium was established in less than 5 min at 25 "C. The uptake of the adsorbate was strongly pH dependent with each isotherm having an inflection as illustrated in Figure 2. The adsorbed amount was obtained from the difference in the concentration of the initially added oxalic acid and that found in the supernatant solution after equilibration. It is important to note that a certain, albeit small, fraction of the acid in solution was present as a ferric complex owing to the partial dissolution of hematite in the presence of the oxalate ion. As will be shown later the dissolution of hematite increases with the lowering of the pH; thus, the amount of complexed acid also was increased in more acidic solutions. In determining the adsorption isotherms the amounts of the free and of the complexed oxalate were taken into consideration. The bound oxalic acid was present as 1:l complex with the ferric ionloand could be established from the analysis of the concentration of the dissolved hematite. The temperature effect on the adsorption is illustrated in Figure 3, which shows that the uptake of the oxalic acid by hematite increased somewhat with rising temperature. The shape of the isotherms remained unchanged, indicating the same nature of adsorbed species. (b) Dissolution of Hematite in the Presence of Oxalic Acid. In blank experiments no ferric ion could be detected in supernatant solutions3 in which hematite particles were suspended, after 17 h of aging the sol at 60 "C over the pH range 1.2-11.3. The addition of oxalic acid into the hydrosol caused dissolution of hematite particles, the extent of which depended strongly on the pH, temperature, and acid concentration. Figure 4 gives the amounts of ferric ion released from a-Fez03in the presence of 2 X mol dm-3 oxalic acid as a function of time at three different pH values, while Figure 5 shows the dissolution data in the presence of varying concentrations of (10) Patel, M. N.; Shah, J. F.; Patel, R. P. J.Indian Chem. SOC.1975, 52, 882.
Hematite-Oxalic Acid and -Citric Acid Systems
Langmuir, Vol. I , No. 2, 1985 203
HEMhTlTE ORGANIC ACID NoNO, 17 h , IO cm-'
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Figure 5. Specific amount of iron(II1) released from hematite at pH 5.2 as a function of the initial concentration of oxalic acid for three different temperatures. oxalic acid at 25,40, and 60 OC,after 20 min of aging. The influence of the pH over the range 3-11 on the amount of hematite dissolved after 17 h of reaction time at an initial mol dm-3 at 25 and concentration of oxalic acid of 2 X 60 OC is illustrated in Figure 6. There is a sharp increase in the solubility of the iron oxide at pH y Fez+,pmol g of hematite 130 81 67 29 5.6 -0 0.09 0.07 0.07 0.04 0.009 -0 [Fe2+]/[FJ+]
These results point to a dissolution mechanism controlled by the release of adsorbed complexes, a process that requires breaking the iron-oxygen bonds at the crystal surface of the adsorbent. An increase in dissolution with rising temperature (Figure 6) is obviously a consequence of such a mechanism. Similar enhancement in dissolution of hematite with rising temperature was observed in the presence of EDTA and other chelating agents! An analogous analysis for release of iron in the presence of citric acid gave a value of g 7 and h 0. While the effect of the proton activity is the same as in the system hematitejoxalic acid, the order in terms of surface complexes is much too high. There is no obvious explanation for such a behavior of citric acid. I t has been shown that a reduction of ferric to ferrous ions enhances the dissolution of iron(II1) ~ x i d e s . ~ J ~ J ~ Testa were made to see if oxalic and citric acids cause reduction of ferric to ferrous ions under the conditions of
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(14)Blesa, M. A.; Borghi, E.B.; Maroto, A. J. G.; Regazzoni, A. E.J . Colloid Interface Sci. 1984, 98,295.
206
Langmuir 1985,1, 206-212
the conducted experiments. Table I shows the presence of solute Fe(I1) ions in sols containing oxalic acid at pH
