Environ. Sci. Technol. 1993, 27, 1864-1870
Kinetics of Fe(I1) Oxygenation at Low Partial Pressure of Oxygen in the Presence of Natural Organic Matter Liyuan Llang,' J. Andrew McNabb, John M. Paulk, Baohua Guy and John F. McCarthy Environmental Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 3783 1-6038
Rates of Fe(I1) oxidation were measured in aqueous solutions of neutral pH, with partial pressure of oxygen (PO,)at 0.005,0.02,0.20 atm. The effects of natural organic matter (NOM), fulvic acid (FA), and polyglutamate on Fe(I1)oxidation were investigated. Atpo, of 0.2 atm, NOM had little effect on the rate of Fe(I1) oxidation. The rate exhibited first-order behavior with respect to Fe(I1) concentration,and the oxidation half-time (tip) was about 20 min. However, at low PO, (0.005 atm), the oxidation rate showed an apparent zero-order behavior, with t l p of 40 h. Organic compounds (NOM, polyglutamate,and FA) enhanced the oxidation reaction. The half time, tllz, shortened to 8 h in the presence of 2.0 mg L-l of NOM. This finding was in contrast to the limited effect of NOM on Fe(I1) oxidation at high poZ. It is interpreted by considering oxidation through two pathways: oxidation of FeOH+ and oxidation of Fe(I1)-organic complexes,with Fe(I1)-organic complexes being the more reactive species.
Introduction The transformation of Fe(1I) and Fe(II1) has significant effects on many natural processesand engineered systems. Ferric oxidefhydroxide formation caused by changes of the redox potentials in water, such as spring and fall turnover in lake systems (2-4), has an impact on the distribution of trace metals and organic compounds through adsorption and coprecipitation (5-7). In addition, the Fe(II1)-oxide particles in surface water attenuate light, affecting photosynthesis in water. Aeration of groundwater that has high Fe(I1) content is an example of iron removal in engineered systems to improve water quality (8-1 1) . A large body of literature has addressed the rate of Fe(11) oxidation under oxygen-saturated conditions (1,8-10, 12-19). Stumm and Lee (1)established a first-order rate law for Fe(I1) oxidation in pH neutral solutions:
Similar first-order oxidation rates were observed in studies of heterogeneous oxidation of Fe(I1) (14, 15). In the presence of inorganic/organic anions and natural organicmatter (NOM),Fe(I1) oxidation also exhibits firstorder kinetics (16,18,19). Theis and Singer (13) showed that although some organic compounds accelerated Fe(11)oxidation (e.g., citric acid), most organic compounds inhibit the reaction, varying in effect from mild (e.g., glutamic acid) to strong (e.g., tannic acid). Humic materials have little effect on the initial stage of oxidation, but thereafter, the rate is greatly reduced. Fe(I1)-organic complexation was hypothesized to be the mechanism underlying the reduced rate of oxidation.
* Author to whom correspondence should be addressed. 1884
Environ. Sci. Technol., Vol. 27, No. 9, 1993
In contrast to the abundance of studies at oxygensaturated conditions, there is little information on rates of Fe(I1) oxidation at low dissolved oxygen (DO). A recent field investigation showed that Fe(I1) oxidation and Fe(111) (hydr)oxide formation occurred in groundwater of neutral pH and fairly low DO, (DO between 0.2 and 1.5 ppm) (20). However, the half-time for Fe(I1) oxidation in this groundwater was 2 orders of magnitude longer than that in oxygen-saturated solutions. Our interest in Fe(I1) oxidation at low DO levels is motivated by current efforts to clean up groundwater contamination. As in surface waters, iron oxidation states and speciation depend on the aquifer chemistry in the subsurfaceenvironment. In turn, the formation of colloidal Fe(II1) (hydrloxides may influence the distribution of contaminants through sorption and transport processes. In an anoxic/suboxic groundwater aquifer, Fe(I1) is the predominant species. The introduction of oxygen into the groundwater will initiate the conversion of Fe(I1) to Fe(III), but the oxidation rate will depend on the level of DO and other ions present in the aquifer (20). At low DO levels, the oxidation rate will be relatively slow, and the effect of organicsolutes and NOM on the oxidation reaction may be more pronounced. The objectives of this paper are to extend the understanding of the kinetics of Fe(I1) oxidation to conditions relevant to many groundwater systems. Specifically, experiments were designed to investigate (1) the rates of Fe(I1) oxidation at low DO conditions, and (2) the effects of a synthetic organic solute and NOM on the rate of Fe(I1) oxidation. Methods and Materials General Information. All chemicals used in the study were of ACS analyzed reagent grade. Pure gases or gas mixtures were certified standard grade (Liquid Air, Oak Ridge, TN). The NOM used in the study was obtained from a pond at Clemson University's Baruch Forest Science Institute in Georgetown, SC. The pond drains a mixed hardwood forest and contains high levels of NOM (66 mg L-l). The raw water was filtered (0.1 pm Amicon hollow fiber filters), and no further treatment was applied to the NOM. The characteristics of the NOM were the same as those in the field investigations (20,22). Suwannee River fulvic acid (FA) was obtained from the International Humic Substances Society and used without further treatment. Iron Analysis. Fe(I1) concentration was determined colormetrically by the 1,lO-phenanthrolinemethod. Powder pillows obtained from Hach were used, and the absorbance was measured at a wavelength of 510 nm with a Hach DR/2000 spectrophotometer. Total iron was measured with Hach FerroVer reagent, which is a basic phenanthroline method, after Fe(II1) was reduced with hydrosulfite and bisulfite. Iron standards were measured daily, and a linear relationship between the concentration and the absorbance was achieved. The error involved in iron measurements was determined to be i3%. The data 0013-936X/93/0927-1864$04.00/0
0 1993 American Chemical Society
presented in subsequent figures are mostly averages of duplicate measurements. Kinetics of Iron Oxide Formation. Experiments were conducted in a 1000-mL reaction kettle (ACE Glass, Vineland, NJ). The top cover accommodates four no. 24 glass stoppers. Adapters (ACE Glass) were used to insert the pH probe, thermometer, and gas inlet. A 24/40 septum stopper was used for sampling and gas venting. The cover was clamped firmly to the reactor base. Experiments were conducted at room temperature (20-22 "C), and the variation in temperature during an experiment was typically less than 1 OC. The solution pH was buffered by a C02/HC03- solution (total bicarbonate was 4.8 mM) and was measured continuously with a combination pH electrode (8103 Ross, Orion). To start an experiment, 1 L of bicarbonate buffer solution (with or without organic compounds) was equilibrated with a gas mixture (3% CO2,0.5-20% oxygen, and a balanced amount of Nz) for 0.5 h at 1000 mL min-l. The DO was measured at the end of equilibration with a colormetric method (Rhodazine-D, Chemetric), and gas purging continued throughout the entire experiment. An aliquot of Fe(I1) stock was then delivered to the solution. Following the introduction of Fe(II), pH was observed to drop sharply up to 0.2 unit, and then to rise back to a steady value. Samples were taken to determine Fe(II), total iron, DO, and particle-size distribution at desired time intervals. The particle-size distribution was determined by a dynamic light-scattering technique (N4MD, Coulter). At the end of experiment, electrophoretic mobilities of the colloidal particles were measured by a doppler electrophoretic light-scattering technique (Coulter DELSA 440). The total dissolved iron (filtered through 3000 mol wt Diaflo ultrafiltration membranes, Amicon) was also determined. The NOM concentration was measured by using a total organic carbon analyzer (Shimadzu, TOC 5000).
0.5 rng L-' NOM
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o r n g ~ "NOM
A
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I
I
I
I
I
0.5
1.o
1.5
2.0
2.5
Time (hr) Flgure 1. Effect of NOM (Georgetown, SC) on rate of oxidation of Fe(I1) at T = 21.5 OC, po, = 0.2 atm, and alkalinity = 4.8 mequiv L-l. The solution pH and the initial Fe(I1) concentration are given in Table I. The lines are drawn as a guide.
100 n v)
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20
30
40
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Time (min)
Results High DO Conditions, Partial Pressure of Oxygen at 0.2 atm. For homogeneous oxidation of ferrous iron, the general rate law has been established as in eq 1. At constant PO, and pH, eq 1 reduces to
Figure 2. Change in total light extinction (measured in thousands of counts per second) with time during oxidation of Fe(I1) at pop = 0.2 atm, alkalinlty = 4.8 mequiv L-I, and near neutral pH.
initial stage of the oxidation shown in Figure 1 is largely the result of the drop in pH. Sung and Morgan (14) also observed this pH variation at the initial stage of oxidation, and they give an extensive discussion of the effects of pH Where kl is the first-order rate constant with a unit of change on the rate constants. Following these author's inverse of time, and kl = k[021[OH-12 = k ~ [ 0 2 1 / [ H + l ~ . treatment, the rate constants, kH, were calculated by Brackets are used to represent concentration, k H is the normalizing the instantaneous rates with respect to the rate constant expressed in terms of H+ concentration, with instantaneous [H+l, [Fe(II)I, and PO, at that time. a unit of M/t. According to eq 2, kl can be determined Alternatively, k H can be obtained by linear regression of from the experimental data (the slope of a linear plot of Itexp,where the steady pH is used in the calculation. Both In [Fe(II)]/[Fe(II)lovs time). However, if either PO, or methods compare well and the rate constants, kH, deterpH is not held constant, kl will not be a constant and the mined for high DO condition (Table I) varied by 30 %, slope determined from In [Fe(II)l/[Fe(II)l~ vs time is As oxidation progresses, the formation of Fe(II1) (hydr)denoted by kexp. oxides will enhance the oxidation rate through the Figure 1shows the change of Fe(I1) concentration with autocatalytic process (14,151. Figure 2 shows the dynamic time in the presence of 0-2 mg L-l of NOM and a constant light-scattering intensity during Fe(I1) oxidation, where po2(0.2 atm). The initial pH was 7.0. As Fe(I1) stock was the light scattering is caused by the formation and the introduced, the pH of the buffer solution dropped sharply growth of iron oxides. The scattering intensity shows an (-0.2 unit) and then gradually (in 20 min) increased to initial steep slope and reaches a plateau at the end of the a steady value. Since the oxidation inversely depends on oxidation. Based on Figure 2, Fe(II1) (hydrloxide forhydrogen ion concentration to a second power, a drop in mation is instantaneous following Fe(I1) oxidation, and pH reduces the oxidation rate. The slower rate at the the autocatalytic effect becomes progressively more sig-
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Environ. Scl. Technol., Vol. 27, No. 9, 1993
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Table I. Summary of Fe(I1) Oxidation Experiments (Alkalinity = 4.8 mequiv L-l,
[Fe(II)lo ( r M )
PO, (atm)
PH
65 60 63 47 42 35 93 87 24
0.2 0.2 0.2 0.02 0.02 0.02 0.005 0.005 0.005
6.86 f 0.12 6.96 f 0.10 6.88 f 0.10 6.94 f 0.06 6.95 f 0.02 6.90 f 0.10 7.03 f 0.02 6.97 f 0.02 6.99 f 0.01
NOM (mg L-l) 2
t1p (h)
T = 20.5 f 0.8 "C)
kHa
0.36 0.22 0.35 4.3 5.2 8.0 8.3 11.8 39.6
(M min-l)
FeT/FeTob( % )
EMc
