Environ. Sci. Techno/. 1995, 29, 1706-1709
DARREN S. BALDWIN* Murray-Darling Freshwater Research Centre and the CRC for Freshwater Ecology, P.O. Box 921, Albury, New South Wales, Australia 2640
JAMES K. BEATTIE AND LYNETTE M. COLEMAN School of Chemistv, University of Sydney, Sydney, New South Wales, Australia 2006 DAVlD R. JONES Environment Protection Program, CSIRO Division of Coal and Energy Technology, P.O. Box 136, North Ryde, New South Wales, Australia 2113
Introduction Organic phosphorus species can represent a large proportion of the phosphorus present in natural waters, sediments, and in soils ( I , 2). The availability of this phosphorus fraction to support nuisance algal and bacterial growth (eutrophication) is dependent in part on the rate of hydrolysis of the substrate to generate inorganic phosphate ions (inparticular orthophosphate). While the importance of enzymes such as phosphatases (31, phytase (41, and nucleases (5)in the hydrolysis of organic phosphorus compounds has been well documented, the role of inorganic substrates in facilitatingphosphate ester hydrolysis in the natural environment has received much less attention. It has long been recognized that metal ions can facilitate the hydrolysis of phosphate esters in homogeneous solution. For example, in 1938 Bamann and Mersenheimer (6) showed that the hydroxides of lanthanum, cerium, and thorium could promote the hydrolysis of glycerol phosphate. Other metal ions including Ir, Rh, Co, Cu, Zn, and Mn (7-10) have all been shown to facilitate the cleavage of coordinated phosphate esters. It is believed that coordination of the phosphate ester to a metal ion results in an increase in electrophilicityat the phosphorus center, which facilitates hydroxyl ion attack ( I 1 ) . The reaction may be further facilitated by the presence of proximally coordinated hydroxyl groups (12). However, because of the exceptionally low concentrations of dissolved transition metal ions in the aquatic environment (relative to the concentration of colloidal, sedimentary, and/or suspended mineral phases),it is probable that homogeneous reactions of this type would not play a significantrole in the cycling of phosphorus. Conversely, heterogeneous reactions involving suspended or sedimentary mineral phases may be important. It has been shown (13) that both anatase (TiOz) and goethite (a-FeOOH)cause a significantincrease in the rate of hydrolysis of the carboxylic acid ester phenyl picolinate. Silica (SiOz),aluminium oxide ( y - A l 2 0 3 ) , and hematite (aFezOs) showed no surfacecatalyticability for this substrate. The efficiency of goethite and anatase mediated hydrolysis was reduced by co-adsorbed ions ( 1 4 ) and by dissolved organic matter (15).
1706 1 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 29, NO. 6,1995
While it is clear that some mineral phases can facilitate hydrolytic reactions, the importance of such reactions to the biogeochemical cycling of phosphorus has not been explored. In this paper, we survey a range of mineral phases for their ability to hydrolyze a model organic phosphorus compound (p-nitrophenyl phosphate).
Materials and Methods p-Nitrophenyl phosphate, p-nitrophenol, Hepes [N-(2hydroxyethyl)piperazine-N'-2-ethanesulfonic acid], Mes [2-(N-morpholinolethanesulfonicacid],Pipes [piperazineN,N'-bis(2-ethanesulfonicacid)], and Tris Itris(hydroxymethy1)aminomethanel were all purchased from Sigma Chemical Co. and used as received. y-AlzOs (aluminium oxide C) and silicon dioxide (AerosilOx 50)were purchased from Degussa, akhtenskite(activatedMnOz)was purchased from Aldrich, rutile (Tion;type UFTR-Z) was a product of Miyoshi Kasei Inc, anatase (TiOz;type CLDD) was aproduct of Tioxide International Ltd., and barium sulfate was a product of Sachtleben. Amorphous iron hydroxide (Fe(OHl3),hematite (a-Fe203),goethite (a-FeOOH), amorphous/a-MnOZ, pyrolusite @-Mn02), and, a-aluminium oxide were prepared using previously published methods (16-19). Samples ofthe clays kaoliniteand bentonite were gifts from Dr. R. Hunter, School of Chemistry, University of Sydney. The structures of the haematite, goethite,rutile, anatase, amorphousla-MnOz, akhtenskite, and pyrolusite samples were confirmed by powder X-ray diffraction measurements using a Siemens Kristoloflex X-ray diffractometer; the patterns were recorded over an angle range of 35" at a collection rate of l"/min. Specificsurface areas ofthe solidswere determined using a Quantasorb BET surface area instrument. The samples were degassed overnight at 150 "C using NP as the carrier gas. An Nz/Hemixwas used for the adsorption/desorption cycles with Nz as the inert adsorbate. The pH of the point of zero charge (pHpZc) for the amorphous iron hydroxide sample was determined by measuring the effect of ionic strength on acid-base titrations of the substrate (20). The points of zero charge for the goethite, haematite, and pyrolusite samples were determined by measuring the pH dependence of electrophoretic mobility of the particles in dilute solution using a Coulter Delsa 440. The points of zero charge of all other samples measured were estimated by determining the pH at which the 9-potential changes sign with a Matec ESA-8000. The number of P adsorption sites was estimated bysuspending0.1 g ofthe mineral phase in 100 mL of 1.6 x M Na2HP04adjusted to pH 6.00 with diluteHCl andlor NaOH. The suspensions were stirred for 6 h at 25 "C and filtered through 0.45-pmfilters,and the phosphate remaining in solution was determined colorimetrically. The difference between initial and final phosphate concentrations was assumed to represent the amount of phosphate adsorbed. All experiments were done in duplicate and corrected for the blank. Surface area, pH,,,, and the concentration of P adsorption sites of the various mineral phases are summarized in Table 1. All hydrolysis experiments were carried out in acidwashed polyethylene screw cap bottles maintained at 25 f 0.1 "C in a water bath. To limit bacterial contamination, all experiments were carried out under aseptic conditions.
0013-936X/95/0929-1706$09.0010
@ 1995 American Chemical Society
TABLE 1
Characteristics and Hydrolytic Ability of a Range of Mineral Phases normalized by no. of P adsorption sitesr (xlO-* L - ' s d yim
mineral phase
specific surface area (dg-l)
concn of P adsorption sites (x (mol g-l)
goethite (a-FeOOH) hematite (a-FezO3) amorphous Fe(OH)3 a m orp h o u d a - Mn0 2 akhtenskite (Mn02) pyrolusite (P-Mn02) a-AlzO3 Y-Ah03 anatase (Ti02) rutile (Ti021 Si02 BaS04 kaolinite bentonite
32.4 32.7 273 135 78.3 1.79 5.98 98.3 158 45.2 N Da ND ND ND
4.3 6.5 5.6 2.2 1.3 ND 1.5 '8 6.6 1.8 ND ND ND ND
control
pH,, 8.6 6.3 7.4 b b 6.5 9.1 9.4 6.4 7.4 =2.0c 3.9d e3.50 Ti > Fe > Al. Furthermore, the amorphous Mn and Fe phases hydrolyze p-nitrophenyl phosphate approximatelythree times more rapidly than the more highly ordered phases (akhtenskite and goethite or haematite, respectively). This latter observation is of note given that amorphous Fe (28)and Mn (29) phases can be the dominant forms of these elements in the water column of many aquatic systems. To gain some insight into the facilitated hydrolysis of organic phosphorus compounds by mineral phases, the effects of pH, substrate concentration, and mineral loading on the rate of hydrolysis by amorphous iron hydroxidewere examined. The effect of pH on the initial rate of hydrolysis of p-nitrophenyl phosphate (50pM)mediated by amorphous iron hydroxide (1000mglL) is shown in Figure 2. From the figure, it is evident that the initial rate of hydrolysis is essentially constant in the pH range 5.5-7.0. As the pH,,, (the pH at which the surface does not carry a net charge) of this mineral is about 7.4 (Table 11, the mineral surface in this pH regime would maintain a net positive charge. Under more basic conditions, the initial rate of hydrolysis of the substrate decreases substantially-there is an approximate order of magnitude decrease in the rate between pH 7.0 and pH 8.5. This decreasein initial rate of hydrolysis is consistent with a decrease in affinity of the (anionic) substrate for the mineral surface as the surface charge becomes negative (that is the pH exceeds the point of zero charge of the mineral). The change in initial rate of hydrolysis as a function of amorphous iron hydroxide loading at both constant pH (7.7)and constant p-nitrophenyl phosphate concentration (50 pM) is presented in Figure 3. vIntvaries linearly with the amount of amorphous iron hydroxide present. The order of the reaction with respect to mineral loading was determined from the slope of a plot of log vint as a function of the logarithm of mineral loading. The plot is linear (r2 = 0.98, n = 12) with a slope of 0.83.
geochemical cycling of phosphorus, it can be seen that under pH conditions and at FelP ratios often encountered in aquatic systems amorphous iron hydroxide can facilitate the hydrolysis of a model organic phosphorus compound. Studies designed to elucidate the detailed mechanism(s) of phosphate ester hydrolysis by mineral phases, and their importance in natural systems are on-going.
I
z A
Acknowledgments
A A
0.5
0.0
1.5
1.0
2.0
2.5
3.0
Mineral Loading (g L') FIGURE 3. Effect of mineral loading on the initial rate of hydrolysis of pnitrophenyl phosphate (50 pM) facilitated by amorphous iron hydroxide and at constant pH (7.7). Errors bars representthe standard error (n 2 2).
20
1 T 1
A
A
7
3 A
"
I' 0
50
100
150
200
250
Concentration PNPP (pM) FIGURE 4. Effect of substrate concentration on the initial rate of hydrolysis of pnitrophenyl phosphatefacilitated by amorphous iron hydroxide (at a loading of 1 g/L)and at constant pH (7.7). Errors bars represent the standard error ( n 2 2).
The change in initial rate of hydrolysis as a function of p-nitrophenyl phosphate concentration at both constant pH (7.7) and constant amorphous iron hydroxide concentration (1000 mglL) is presented in Figure 4. From the figure, it can be seen that Yint varies linearly with substrate at low concentrations. The initial rate of hydrolysis tends to level off at higher concentrations of substrate. A plot of log vint as a function of the logarithm of p-nitrophenyl concentration is linear ($ = 0.98, n = 10) with a slope of 0.62. Afractional reaction order in heterogenousreactions is indicative of a transition from first- to zero-order kinetics, suggesting that the mineral surface may be approaching saturation (30). While it is not yet certain that facilitated hydrolysis of phosphate esters is an important pathway in the bio-
T. Kelso and C. Fong of the Department of Chemistry, University of Sydney are thanked for synthesizing the a-MnOz andg-MnOz samples,respectively. Ms. G. Gleeson and Mr. M. Charge are thanked for their technical assistance with the amorphous iron hydroxide experiments. Funding for this project came from the ARC Small Grants Scheme, CSIRO Coastal Zones Program, and from funds made available by the Murray-Darling Freshwater Research Centre (a partner in the Co-Operative Research Centre for Freshwater Ecology).
Literature Cited (1) Broberg, 0.;Person, G. Hydrobiologia 1988, 170, 61. (2) Sommers, L.; Harris, R.; Williams, J.; Armstong, D.; Syers, J. Limnol. Oceanogr. 1970, 15, 301. (3) Jansson, M.; Olson, H.; Pettersson, K. Hydrobiologia 1988,170, 157. (4) Herbes, S.; Allen, H.; Mancy, K. Science 1975, 187, 432. (5) Hino, S. Hydrobiologia 1989, 174, 49. (6) Bamann, E.; Mersenheimer, M. Ber. Dtsch. Chem. Ges. 1938, 71, 1711. (7) Hendry, P.; Sargeson, A. Inorg. Chem. 1990, 29, 97. (8) Hendry, P.; Sargeson, A. A u t . 1.Chem. 1986, 39, 1177. (9) Jones, D.; Lindoy, L.; Sargeson, A. 1.Am. Chem. SOC. 1984, 106, 7807. (10) Tetas, M.; Lowenstein, J. Biochemistry 1963, 2, 350. (11) Dixon, N.; Jackson, W.; Marty, W.; Sargeson, A. Inorg. Chem. 1982, 21,688. (12) Jones, D.; Lindoy,L.; Sargeson, A. 1.Am. Chem. SOC. 1983, 105, 7327. (13) Torrents, A,; Stone, A. Enuiron. Sci. Technol. 1991, 25, 143. (14) Torrents, A.; Stone, A. Enuiron. Sci. Technol. 1993, 27, 1060. (15) Torrents, A.; Stone, A. Enuiron. Sci. Technol. 1993, 27, 2381. (16) Schwertmann, U.; Comell, R. Iron Oxides in the Laborutory; VCH: New York 1991. (17) Sugimoto, T.; Sakata, K.; Muramatsu, A. 1. Colloid Interface Sci. 1993, 159, 372. (18) More, T.; Ellis, M.; Selwood, P. 1.Am. Chem SOC. 1950, 72, 856. (19) Wefers, K.; Misra, C.; Oxidesand Hydroxides ofAluminiumAlcoa Technical Paper No. 19;Alcoa Laboratories: Pittsburgh, PA, 1987. (20) Detenbeck, N.; Brezonil, P. Enuiron. Sci. Technol. 1991,25,395. (21) Chang, R. Physical Chemistry with Applications to Biological Systems; McMillan Publishing Co.: New York 1977; p 373. (22) Holbrook, K.; Ouellet, L. Can. 1. Chem. 1958, 36, 686. (23) Shaung, C.; Huang, P. M.; Stewart, J. W. B. Can. 1.Soil. Sci. 1990, 70, 461. (24) Shaung, C.; Stewart, J. W. B.; Huang, P. M. Geoderma 1992, 53, 1. (25) Gaudin, A.; Fuerstenau, D. Trans Am. Inst. Mater. Eng. 1957, 208, 1365. (26) Yousef, A.; Bibawy, T.; Malali, M. Chem. Ind. 1977, 6, 229. (27) Drover, J. The Geochemistry of Natural Waters Prentice Mall: EnglewoodCliffs,NJ, 1982. (28) Tipping, E.; Woof, C.; Ohnstad, M. Hydrobiologia 1982, 92,383. (29) Lind, C. J.; Hem, J. D.; Robertson, J. E. Reaction Products of
Manganese-Bearing Waters. In Chemical Quality ofwater and the Hydrologic Cycle;Averett, R. C., McNight, D. M., Eds.; Lewis: Chelsea, MI, 1987; pp 273-301. (30) Bond, G. Heterogeneous Catalysis Claredon Press: Oxford, 1987.
Received f o r review October 25, 1994.Revised manuscript received February 28, 1995. Accepted March 7,1995.
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