228
Anal. Chem. 1981, 53, 228-232
Reduction of Ionic Species by Fulvic Acid R. K. Skogerboe" and S. A. Wilson' Department of Chemistty, Colorado State University, Fort Collins, Colorado 80523
Studies have shown that a fulvlc acld derived from sol1 is capable of reducing Hg(I1) to Hg(O), Fe(II1) to Fe(II), and I2 and Is- to I- under conditions generally characteristic of natural waters. The evaluation lndlcates a reduction potential approximating 0.5 V (vs. normal hydrogen electrode) for fulvlc acid whlch Is in agreement with a prevlous estlmate. Although the effect of pH on the reduction processes changes between redox couples, the results generally lndlcate that hydrogen ion is consumed when fulvic acid is oxldlzed. These results quite distinctly Indicate that fulvlc acid may play a prominent role in influenclng the redox equlllbrla occurrlng In surface and ground waters.
Reviews by Gamble and Schnitzer ( I ) and by Reuter and Perdue (2)have emphasized the probable importance of fulvic acid (FA) as a complexing agent for metal ions in surface and ground water systems. Such possibilities have led to the evolution of quite intensive research efforts to define the metal complexation characteristics of FA (3-25). A few publications have considered the redox properties of FA and humic acid (HA) (26-33) but most have focused on HA (26-32). A recent paper by Wilson and Weber (33)demonstrated that FA was capable of reducing vanadium(V) to vanadium(1V) and provided rather convincing data indicating that the reduction potential (Eh) for FA at p H 2 was 0.5 V vs. the normal hydrogen electrode ("E). This indicated that FA was a better reducing agent than HA; Szilagyi (26,27)has reported a HA reduction potential of 0.7 V (vs. "E). Since the latter has been shown capable of reducing Hg(I1) to Hg(0) and Fe(II1) to Fe(II), the lower Ehvalue for FA strongly suggested that it would also reduce these species (34). The purpose of this investigation has been to examine this possibility. EXPERIMENTAL SECTION
Reagents. All reagents used were of reagent grade or better. The KOH solutions were prepared by dissolving recrystallized KOH in distilled, deionized water. Nitric acid solutionswere made from acid doubly distilled in quartz and distilled,deionized water. The fulvic acid used was isolated from the Bzhorizon of a podzol soil obtained from Conway, NH, using the procedure outlined by Weber and Wilson (35). The freeze-driedmaterial obtained was dissolved in water that was triply distilled over KMn04 and stored in sealed polyethylene containers. Since this material was obtained from the same soil and isolated by the same procedures used by Weber and Wilson (35), it has been assumed that its general properties are those of that described by Weber and Wilson (36, 37) and studied by Weber and associates (38, 39). Thus, the dissociation corrected number average molecular weight of 644 has been used in all calculations. Further, the percentage C, H, N, and ash values of 53.1, 3.24, 0.90, and 0.8, respectively, may be generally representative of the soil fulvic acid used in the present studies. Similarly, total acidity, carboxyl, phenol OH, and carbonyl functional group values of 13.4, 8.2, 5.2, and 3.5 mequiv g-l, respectively,may be nominally representative of the present FA (37). Standard Fe(II1) solutions were prepared by dilution of a stock solution made by dissolving a known amount of ferric nitrate in an appropriate volume of triply distilled water. Iron(I1) standards were prepared by dissolving ferrous sulfate in Present address: US.Geological Survey, 5293 Ward Road, Arvada, CO 80002.
triply distilled water. All pH adjustments were made by adding microliter amounts of standard KOH and HN03to obtain a 1.0 X lo4 M stock solution. Dilutions of this were prepared just prior to use in each instance. The sodium borohydride (5 M) used for calibration of the Hg analysis system was prepared in distilled deionized water and stored in a refrigerator until used. Iodine solutions were prepared by dissolving recrystallized Iz in water and standardizing via thiosulfate titration. Iodide solutionswere prepared from KI while the 1, solutions were prepared by mixing 12 and I- solutions in known ratios. Procedures. For evaluation of the reduction of Fe(II1) to Fe(I1) by fulvic acid, the FA concentration was held constant at 5.0 X lo4 M while the Fe(II1) concentrationswere systematically varied to 6.0 X lo4 M at the respective pH over the range of 5.0 X values studied. All test solutions were initially adjusted to a pH greater than 8.0 and then readjusted to the required pH value for each test series. All solutions were prepared in acid-washed Erlenmeyer flasks, covered tightly with parafilm to exclude air, and allowed to stand at ambient temperature for 48 h prior to analysis. The Fe(I1) concentrationsproduced by reaction of Fe(II1) with FA were determined by the o-phenanthrolene (0-P) method described by Sandell (40). Aliquots of 4 mL of 1.0 X lo-$ M of 0-P solution were placed in 10-mLvolumetric flasks and diluted to volume with the Fe-FA test solutions, and the color was dowed to develop until constant absorbance values were observed. At low pH (54.0) this condition was attained in about 4 h but -22-24 h were required for higher pH solutions. Therefore, the effects of pH on the absorbance and time stability of the Fe(I1):O-P complex were examined carefully by using standard solutions without FA being present. These studies showed that pH did not affect the absorbancesignificantlyand that the color developed was stable for at least 48 h. Therefore, the absorbance values for all test solutions were made at 508 nm after the color had been allowed to develop 24 h. These measurements were made in 1-cm quartz cells with a spectrophotometer (Coleman-Hitachi, Model 101) that was calibrated with standard Fe(I1) solutions treated in parallel via the same procedures. Replicate measurements generally agreed to fl-2% or better. All pH measurements were made with a combination electrode calibrated against buffer solutions. To determine the effects of FA on Hg(II), we placed measured volumes of solutions of known FA concentration in a reaction chamber with a fritted bottom; a nitrogen stream (500 mL/min) passing upward through the frit and the FA solution provided extensive solution agitation. This stream passed through a Mg(C104)2/CaS04packed drying tube to remove water vapor and thence into a quartz cold cell (16.5 cm long X 0.4 cm diameter) placed in the optical path of an atomic absorption unit (Techron, Model AA5). Thus, Hg released from the cell was monitored by measuring the absorbance at 253.7 nm by using a hydrogen hollow cathode lamp for background correction. In a typical experimental M FA solution adjusted to a sequence, 1 or 2 mL of a 5 X particular pH was placed in the reaction cell. Then an aliquot (0.2 mL) of a standard Hg(I1) solution, adjusted to the same pH just prior to use, was injected into the cell with a gastight syringe, and the absorbance signal due to metallic Hg vapor produced was recorded on a strip-chart recorder. Successive aliquots were injected into the same FA solution to determine the reduction capacity of said solution. The absorbance signals obtained were quantitated by measuring the peak areas with an integrator. These data were converted to mass of Hg released by reference to standard curves obtained by the redyction of known amounts of Hp(I1) with addition of excess amounts of sodium borohydride soiution (41). For determination of the reduction of I2and 1, by FA, solutions of known concentrations and pH of these respective entities were
0003-2700/81/0353-0228$01.00/00 1981 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 53, NO. 2, FEBRUARY 1981
229
41
a
0.
Figure 1. Example of the effects of sequential Hg(I1) injections on the rate and extent of Hg(0) generation. Each peak due to injection of 5 X M FA without replacement mol of Hg(I1) into 1 mL of 5 X of the latter.
I
I
1
0.5
1.5
2.5
,
3.5
CUMULATIVE MOLES Hgilll INJECTED [xlO*l
placed in the sample cell of a spectrophotometer. An equal volume of water at the same pH was placed in the reference cell. Decreases in the absorbance levels due to I2 and Is- were measured, at the absorption wavelengths characteristic of each, when successive 25-pL volumes of a FA solution of known concentration and the same pH were added to both the reference and sample cells. In this way, any absorption at the respective wavelengths due to the presence of FA was cancelled and the pH at which the reactions occurred was controlled. The depletions in the concentration of I2and 1, due to their reduction by FA were corrected for dilution for each increment of FA solution added.
Flgure 2. Relative Hg reduction efficiencies for sequential injections at varlous pH levels: 1 mL of 4.2 X lo-' M FA; 5 X lo-' mol of Hg(I1) per injection; (0)pH 2; (A) pH 3.1; (0)pH 4.1; (0)pH 6.1; (0)
PH 7.8.
RESULTS AND DISCUSSION Mercury. Examples of the absorption signals observed for the mercury vapor generated when standard mercury solutions of known pH were injected into a fulvic acid solution at the same pH are given in Figure 1. Two general features may be noted from these absorption curves. First, the appearance of mercury vapor in the absorption cell was almost instantaneous in relation tQ the injection. However, as the reduction reactions involving fulvic acid proceeded with concomitant decreases in the ionic mercury and the oxidized fulvic acid concentrations, the rate of mercury vapor generation decreased also. This effect is illustrated by the peak broadening and increased tailing in Figure 1 when successive aliquots of a Hg(I1) solution were injected into a fulvic acid solution without replenishment of the latter. Although the absorption peak heights in Figure 1decrease with each injection, the peak areas remained nearly constant while there were changes in the reduction kinetics as evidenced by the tailing of the peaks. This continued until a cumulative excess of Hg(I1) had been added. Examples of this which also indicate the general effect of pH are given in Figure 2. These data indicate that successive injections of 5 X lo4 mol of Hg(I1) in 0.2-mL aliquots into 1.0 mL of a 4.2 X M FA solution without replenishment of the latter resulted in the reduction of up to 20% of the Hg (1X mol) per injection. The amount converted tended to decrease with successive injections depending on pH. It must be emphasized, however, that the apparent decreases in conversion efficiency are due at least in part to reductions in the rate of conversion. As the amount of FA available for reaction decreased with successive injections or the pH of the reaction medium increased, the rate of conversion decreased. This was evidenced by the increased tailing of the absorption peaks at higher pH levels and with successive injections (Figure 1). Thus, in the latter stages of each injection cycle, the rate of Hg evolution was often reduced to the point where the Hg concentration delivered to the absorption cell per unit time was not detectable. The upward shift of the base line in Figure 1, for example, was actually due to continued evolution of Hg vapor at a slower rate; injection of additional aliquots of Hg(I1) solution shifted the Hg(I1) concentrations upward with concomitant initial increases in the rate of production. The apparent degree and
a
"r 1
2
,-*,/',
I
5
,
6
\
'., , I 7
PH
Flgure 3. Correlation of the effects of pH on the amount of Hg reduced with the extent of FA lonizatlon.
rate of Hg(I1) reduction tended to decrease with increasing pH (Figure 2). Since hydrolysis of Hg(I1) to form HgOH+ and complexation of metal ions by FA both tend to increase with pH, the reduction process may be hindered by either. Since the first Hg(I1) injection into a FA solution tended to show sharp Hg evolution peaks that returned to the base line in a few seconds, the general effects of pH on the reduction process were evaluated on the basis of single rather than sequential injections. The results summarized in Figure 3 verify that the fraction of the mercury concentration reduced decreased sharply at pH levels above -4 and attained a relatively constant value at pH values above -5.5. Thus, the reduction reaction(s) occur at pH levels characteristic of natural waters although the rates and extents tend to decrease with pH. Figure 3 also draws a parallel between the extent of Hg reduction and the effects of pH on the degree of ionization of the fulvic acid. Titrations of the acid with standard KOH have indicated the presence of two types of ionizable functional groups having mean pK values of 2.48 and 4.98, respectively (42). These values have been used to calculate the fractional distributions of the fulvic acid as the singly (HFA-) and doubly ionized (FA2-) species shown in Figure 3. The inverse parallel between the FA%fraction and the extent of Hg(I1) reduction may be due to increased complexation of the metal by the FA2- or it may be indicative of the necessity of the FA being protonated to enhance its role as a reducing agent. Calculations to assess the possible significance of the formation of HgOH' ( K = 3 x loll) (43)indicate that nominally
230
ANALYTICAL CHEMISTRY, VOL. 53, NO. 2, FEBRUARY 1981 I
3.
I
4 j
0
,
Ti"
L
2.
!I
P
1
I
1.
i
O \
B
P/'
4INITIAL Felllll CONCENTRATW, mio4
Flgure 4. Degree of Fe(II1)reduction by 5 X lo4 M FA related to inltlal Fe concentration: (A) pH 5.5; (A)pH 5.7. Other symbols are same as Figure 2.
23,75,97, and 100% of the Hg(1I) would be present in this form a t pH levels of 2, 3,4, 5 and above, respectively. The fact that significant fractions of the Hg(I1) could be present in this form at pH levels coincident with maximum reduction (pH 54) and that significant reduction is maintained at pH levels where essentially all Hg could be in this form, tends to negate the significance of Hg hydrolysis as a possible limitation on the reduction reaction. Iron. The fact that Hg(I1) with a standard reduction potential ( E O ) of 0.85 V (vs. NHE) was reduced by fulvic acid, combined with the FA reduction potential estimate of 0.5 V (331, suggested that FA should also be capable of reducing Fe(1II) to Fe(I1) (Eo = 0.771 V vs. "E). In initial measurements of the amount of Fe(II1) converted to Fe(II), it was noted that development of color due to the formation of the iron(I1)-o-phenanthroline (0-P)complex occurred much slower in the presence of fulvic acid; i.e., a few minutes were sufficient for full color development in the absence of FA but 20-24 h were necessary in its presence. Therefore, the possibility that the redox kinetics were slow was checked by sampling a solution of FA and Fe(II1) at a 5/1 molar ratio at expanding time intervals after mixing, immediately adding the color developing reagents, and monitoring the absorbance of the Fe(I1) complex a t regular intervals until color development became constant. The results indicated that the Fe(II1) reduction process occurred fairly rapidly reaching completion in 12-15 min. Thus, the reduction reaction occurred quite rapidly but the development of the iron(II)-ophenanthroline color was apparently delayed by Fe(I1) complexation competition between FA and the 0-P. All subsequent measurements allowed 6 h or more for reaction completion and 24 h for full color development. The effects of varying the initial concentration of Fe(II1) added to a solution of 5 X lo-" M FA are illustrated in Figure 4. Although these two experiments were carried out at pH levels of 5.5 and 5.7, the general reduction reaction trends are apparent. Appreciable fractions of the Fe(II1) were reduced by FA at all concentrations investigated. Although the fractions of the total Fe(II1) reduced decreased with ita initial concentration, the results indicate that 1mol of FA is capable of reducing at least 0.5 mol of Fe(II1) at this pH. To evaluate the effect of pH on the reduction reaction, we measured the degree of reduction for solutions, at each of several pH values, containing Fe(II1) at 1.0 X M and FA at 5.0 x M so that the latter was present in excess. The results summarized in Figure 5 indicated a low degree of reduction at pH 2 that increased sharply to a nearly constant value in the 3-5.7 pH range and dropped thereafter. For
... ....a.
I
1'
2
3
.*..*. ... 4
8
,
6
*.
,
7
pn
Figure 5. Effect of pH on the degree of Fe(II1)reduction compared to calculated chemical dlstribution of Fe species. (Vertlcal bars connect
points from duplicate measurements.)
determination of whether the drop at higher pH was due to precipitation of Fe(OH)3,4 mg of Fe(OH)3 precipitate was collected on a Teflon filter without drying. Equal halves of this filter were placed in respective Teflon beakers containing 20 mL of 5.4 X loW4 M FA at pH 6. Five-milliliter aliquots were withdrawn at the end of 15 min and 15 h and were analyzed for Fe(I1). The resulta indicated 0% and 0.3% mean reduction values for the 15-min and 15-h samples, respectively. Thus, the reduction of solid Fe(OH)3does not occur rapidly, but the amount converted over the extended time period confirms that reduction of a common precipitated form of Fe(II1) can occur in nature. The rate of reduction is slower, however, than that observed for soluble forms of Fe(II1). Complexation of Fe(II1) by FA or by OH- (to form Fe(OH),+), might also be responsible for the decrease in reduction observed above pH 5.7. However, the fraction of Fe(II1) complexed by FA should be essentially equivalent a t pH 5.7 and 6.0 since the fraction of the FA in the doubly ionized form would change only about 10% between these pH values (see Figure 3). Similarly, Schnitzer and Skinner (44)have shown that the solution concentration of OH increased when Fe(II1) reacted with organic matter extracted from soil and concluded that this was due to the involvement of Fe(OH),+ rather than Fe(II1) such that OH was released via the complexation reactions. Indeed, the stability constants for FeOH2+and Fe(OH),+ approximate lo1' and 5 X loz1(43),respectively, such that the fractions of the Fe(II1) present in various forms at each pH may be estimated by calculation. The calculated Fe(II1) distribution diagram included in Figure 5 may be used to draw some circumstantial inferences. First, comparison with the fraction of Fe(II1) reduced indicates the lowest degree of reduction at pH 2 where the fraction present as Fe3+ is predominent. This, combined with the increase in reduction as FeOH2+becomes the most prominent species and the decrease as Fe(OH),+ begins to predominant, implies that the reduction reaction may preferentially involve the FeOH2+ species. These possibilities must, however, be considered speculative. Of primary importance is the observation that FA is capable of reducing Fe(II1) to Fe(I1) under conditions nominally representative of natural water systems. Iodine and Triiodide. As a further means of characterizing the reduction capabilities of FA, its ability to reduce 1 2 to I(Eo = 0.62 V VS. NHE) and 13-to I- ( E O = 0.536 V vs. NHE) was investigated. The results for I, reduction summarized in Table I are consistent with those for Fe since less FA was required to
ANALYTICAL CHEMISTRY, VOL. 53,
0.5 1.0 1.5 2.0 2.5 a For 2 x
percentage of I, reduced at pH 3.0
4.0
231
E", V
5.5
16.8 25.6 35.2 52.0 55.2 77.6 76.0 98.3 98.4 100 M solution of I,.
FEBRUARY 1981
Table 111. Standard Reduction Potentials of Selected Half-Reactionsa
Table I. Fractions of I, Reduced by Various Concentrations of Fulvic Acid at Different pH Levelsa molar ratio FA112
NO. 2,
half-reaction
(vs. NHE)
+ 2Ht + e- VOz+ + H,O Hg2++ 2 e - Hg Ag+ + e- -t Ag Hg," + 2e- 2Hg Fe3++ e - Fez+ H,SeO, + 4 H + + 4e- Se + 3H,O humic acid 0, + 2H+ + 2e- -+ H,O, I,(aq) + 2e- -+ 21H,AsO, + 2H+ + 2e- -+ H,AsO, + H,Q I; t 2e' 31Cu+ + e' -+ Cu FA(ox) + 2" + e- FA(red) VOz++ 2Ht +
1.00
33
0.854 0.799 0.789 0.771 0.74
present
0.7 0.682
26, 27
0.62 0.56
present
0.536 0.521 0.5
present
0.361
3 3 (no rxn)
VO;
34.4 67.2 98.4 100 100
ref
-+
--f
-f
-f
present
-+
Table 11. Fractions of I; Reduced by Various Concentrations of Fulvic Acid at Different pH Levelsa molar ratio FA/I;
a
0.2 0.4 0.6 0.8 1.0 1.2 1.4 For 7 x
percentage of I; reduced at pH 2.7
3.8
0 0 0 0 0 3.1 5.5
0 0
5.7 1.6 3.1
3.9 1.3 3.1 5.5 8.7 3.9 6.3 15.7 8.7 22.8 M I; solution.
6.3
7.3
4.7 10.2 16.5 23.6 31.5 40.9 50.4
7.1 15.0 29.1 41.7 55.9 72.4 100
--f
e - + V3' + H,O Cuz++ 2e- 2 Cu UO,Z+ + 4H++
reduce equivalent amounts of 12 at higher pH levels up to at least pH 5.5. Similar observations are recorded for the reduction of ,1 in Table 11. Again, larger amounts of FA were required to cause significant reduction when the solution pH was at levels concomitant with the acid sites being minimally ionized. Thus, these reduction measurements combined with those reported by Wilson and Weber (33) are generally indicative of the redox properties of fulvic acid. Since the standard reduction potentials of I2 and Is- are 0.62 and 0.536 V, respectively, the present results are reasonable confirmations of the fulvic acid reduction potential of 0.5 V estimated by Wilson and Weber (33). Summary. Wilson and Weber (33)reported that a lo-fold molar excess of FA at pH 2 was capable of reducing 50% of a vanadium(V) solution to V02+,but reduction to V(II1) did not occur over a 4-h period. In an equivalent solution at pH 6, only 4.5% of the V(V) was reduced to V(1V). They also reported a potential of 0.57 V (vs. NHE) for a solution containing FA and H2V04-in a 10/1 molar ratio at pH 2.3 and a potential of 0.50 V for FA based on an &,--pH measurement and indicated that the following half-reaction model for the oxidation of FA was most consistent with their results:
FA(Red)= FA,,,
+ mH+ + e-
(1)
The present investigation has concentrated on the reduction of Hg2+(Eo= 0.854 V), Fe3+ (Eo= 0.771 V), I2 (E" = 0.62 V), and 1, (Eo= 0.536 V). The observation that FA is capable of reducing 13- verifies that its reduction potential approximates that for 13-within a few tenths of a volt. Thus, the estimate of 0.50-0.51 V given by Wilson and Weber (33)has been semiquantitatively verified herein. The reduction half-reactions of I2and Is- do not depend directly on pH. An indirect effect may have been operative, however, due to hydrolysis of iodine to form hypoiodite which tends to become more prominent in alkaline solutions. Any hypoiodite formed may undergo a pH-dependent disproportionation to form iodate and iodide. Thus, the observation that the degree of I2 reduction increased with pH (Table I) may have partially reflected the occurrence of such disproportionation reactions. However, scans of the absorption wavelength for iodate ion at the conclusion of each set of experiments failed to indicate the presence of said ion at levels
33 and
present
-+
0.337 0.334
2e- ZU4++ 2H,O Values of standard reduction potentials ( E " ) of all but those for fulvic and humic acid taken from Latimer ( 3 4 ) . above lo+ M thereby negating this possibility. It was consequently concluded that increased reduction with increasing pH was most likely due to an enhancement in the reduction capabilities of FA as the acidic functional groups present became more ionized. At the same time, the reduction of Hg(I1) appears to be enhanced by the presence of H+(Figures 4 and 5). The accuracy of this observation may, however, be brought into question by the occurrence of the reduction at slower rates at higher pH levels such that the Hg vapor measurements were not completely accurate as suggested above. The results in Figure 3, for example, suggest quite clearly that the rate of mercury vapor generation in the later stages of the reaction may have been too slow to permit detection of the Hg. In addition, the rate or extent of reduction may have been affected by complexation of Hg(I1) by OH or FA forming Hg(OH)+ and HgFA, both reactions being enhanced at higher pH. Finally, the diminished degree of reduction of Fe(II1) at low pH is parallel to the effect observed for 1 2 and 13-while the diminution of Fe reduction at high pH parallels that for Hg. Weber and Wilson (33) postulated reaction 1 as the pHdependent half-reaction for the oxidation of FA. In their vanadium studies, it was indicated that the value of m = 2 was most consistent with their results. Moreover, their &-pH studies under aerobic and anaerobic conditions indicated that the E h values were very similar for both situations and that the FA reduction potentials increased with decreasing pH. Thus, H+ was consumed when FA acted as an oxidizing agent. All of these observations parallel those reported by Szilagyi (26,271for peat humic acid suspensions. The present measurements on 12, 13-,and Fe(II1) reduction generally confirm the above pH dependence of the reduction potential for FA; the pH dependence discrepancy observed for Hg(I1) reduction may be due to the kinetic effect discussed above or to other unrecognized factors. Of primary signifiance is the demonstration that FA is capable of reducing a variety of chemical entities under conditions generally representative of surface and ground water systems. The potential role that FA may play in affecting the oxidation states of constituents in natural water systems may be generally deduced by examination of the selected data
232
Anal. Chem. 1981, 53, 232-236
summarized in Table 111. Indeed, the reduction potential of -0.5 V for FA indicates that it is a stronger reducing agent than humic acid and that it can clearly cause reduction of a variety of ionic and nonionic species which may be present in natural aqueous systems. It appears, for example, that O2 can be reduced by FA with subsequent reduction of the Hz02 formed (Eo= 1.77 V) thereby accounting, at least in part, for the minimal effects of aerobic conditions on the E,,measured for FA (33). Certainly, the ability of FA to reduce Fe and Hg, as well as H2Se03and H3As04and even humic acid, raises a number of possibilities of significance to numerous environmental and geochemical questions. These possibilities require further evaluation, however. Experiments designed to provide this information are in progress. LITERATURE CITED Gamble, D. S.;Schnitzer, M. I n "Trace Metals and Metal-Organic Interactions in Natural Waters"; Singer, P. C., Ed.; Ann Arbor Science Publishers: Ann Arbor, MI, 1973. Reuter, J. H.; Perdue, E. M. Geochim. Cosmochim. Acta 1979, 41, 325-334. Stevenson, F. J. Soil Sci. 1977, 723, 10-17. Sposito, G.; Holtzclaw, K. M. SoilSci. SOC.Am. J. Ig79, 43, 47-51. Sposlto, G.; Holtzclaw, K. M.; LeVesque-Madore, C. S . Soil Sci. SOC. Am. J. 1978, 43, 600-607. Stevenson, F. J.; Krastanov, S.A.; Ardakoni, M. S . Geoderma 1973, 9 , 129-141. Guy, R. D.; Chakrabarti, C. L. Can. J. Chem. 1976, 54, 2600-2611. Buffle, J.; Greter, F. L.; Haerdi, W. Anal. Chem. 1977, 49, 216-222. Bresnihan, W. T.; Grant, C. L.; Weber, J. H. Anal. Chem. 1978, 50, 1675-1679. Chearn, V. Can. J. SoilSci. 1973, 53, 377-382. Brady, B.; Pagenkopf, G. K. Can. J. Chem. 1978, 56, 2331-2336. Whitworth, C.; Pagenkopf, G. K. J. Jnorg. Nuci. Chem. 1979, 41, 317-321. Chau, Y. K.; a c h t e r , R.; Lumrn-Shue-Chan, K. J. Fish. Res. Board Can. 1974, 37, 1515-1519. Mancy, K. H.; Prog. Water Technoi. 1973, 3 , 63. Hanck, K. W.; Dillard, J. W. Anal. Chim. Acta 1977, 89, 329-340. Chau, Y. K.; Lumm-Shue-Chan, K. Water Res. 1974, 8 , 383-388. Shuman, M. S.;Woodward, G. P. Anal. Chem. 1973, 45, 2032-2035.
(18) O'Shea, T. A.; Mancy, K. H. Anal. Chem. 1976, 48, 1603-1607. (19) Davison, W.; Whitfield, M. J. flectroanai. Chem. 1977, 75. 763-770. (20) Buffle, J.; Greter, F. L.; Nembrini, G.; Paul, J.; Haerdie, W. Z . Anal. Chem. 1976, 282, 339-345. (21) Nurnberg, H. W.; Valenta, P.; Mort, L.; Raspor. B.; Slpos, L. Anal. Chem. 1976, 282, 357-361. (22) Greter, F. L.; Buffle, J.; Haerdi, W. J. Electroanal. Chem. 1979, 707, 211-230. (23) Buffle, J.; Greter, F. L. J. Electroanal. Chem. 1979, 707, 231-251. (24) Brezonik, P. L.; Brauner, P.; Stumm, W. Water Res. 1978, 70, 605-612. (25) Siegerman, H.; ODom, G. Am. Lab. (Fairfieid, Conn.) 1972, 4 , 59-68. (26) Szllagyi, M. SoilSci. 1971, 1 7 7 . 233-238. (27) Szilagyi, M. Soli Sci. 1973, 775, 434-441. (28) Szalay, A.; Szilagyi, M. Geochim. Cosmochim. Acta 1967, 37, 1-14. (29) Alberts, J. J.; Schindler, J. E.;Miller, R. W.; Nutter, D. E. Jr. Science 1974, 184, 895-902. (30) Bloomfield, C.; Kelso, W. I. J. Soil Sci. 1973, 24, 368-375. (31) Goodman, 8. A.; Cheshire, M. V. Geochlm. Cosmochim. Acta 1975, 39, 1711-1716. (32) Sakatos, B.; Tibai, T.; Meisel, J. Geoderma 1977, 19, 319-323. (33) Wllson, S. A.; Weber, J. H. Chem. Geol. 1979, 26, 345-351. (34) Latimer, W. M. "The Oxidation States of the Elements and their Potentlals in Aqueous Solutions", 2nd ed.; Prentice-Hall: Englewood Cliffs, NJ, 1952. (35) Weber, J. H.; Wilson, S. A. Water Res. 1975, 9 , 1079-1084. (36) Wilson, S. A.; Weber, J. H. Chem. Geoi. 1977, 19, 285-293. (37) Wllson, S. A.; Weber, J. H. Anal. Lett. 1977, 10, 75-84. (38) Bresnlhan, W. T.; Grant, C. L.; Weber, J. H. Anal. Chem. 1976, 50, 1675-1679. (39) Saar, R. A.; Weber, J. H. Can. J. Chem. 1979, 57, 1263-1268. (40) Sandell, E. B. "Colorimetric Determinatlon of Trace Metals"; Interscience: New York, 1959. (41) Natusch, D. F. S.;Tucker; M. D. M.; Miller, J. A.; Schmldt, F. Q. Anal. Chem., in press. (42) Wilson, S. A.; Huth, T. C.; Arndt, R. E.; Skogerbce, R. K. Anal. Chem. 1980, 52, 1515-1518. (43) Sillen, L. G.; Martell, A. E. "Stability Constants of Metal-ion Complexes"; The Chemlcai Society: London, 1964. (44) Schnitzer. M.; Skinner, S. I. M. So// Scl. 1966, 702, 361-372.
RECEIVED for review February 19,1980. Accepted November 6,1980. Research supported by the EnvironmentalProtection Agency under Grant No. R805183-03.
Tensammetric Determination of Polyoxyethylenated Alcohols with Low Oxyethylene Content at Parts-per-Million Concentrations Milton J. Rosen," Xi-yuan Hua, Peter Bratln, and Anna W. Cohen Department of Chemistry, Brooklyn College, City University of New York, Brooklyn, New York 71210
Differential double-layer capacltance vs. potential measurements at the dropplng mercury electrode are used for the determinatlon of polyoxyethylenatedn-dodecyl alcohols wlth low oxyethylene content at part-per-mllllon concentratlons. The method Is sensitive to materlals with 2 or more mol of ethylene oxide. The method has an accuracy of f2.5% for concentratlons down to 6 ppm and 3 5 % for concentratlons down to 3 ppm. The determlnatlon Is done directly on the surfactant solution without extraction or flltration steps and without the addition of any reagent except supportlng electrolyte. The potential of the hlghest desorption peak In the capacltance-potential curve can be used to estlmate the oxyethylene content of the polyoxyethylenated alcohol.
The literature on the analysis of polyoxyethylenated alcohols reveals that there are only a few methods available for the accurate determination of these materials at low con-
centrations (parts per million). A method involving formation of a cobaltothiocyanate complex, extraction of the complex into an organic solvent, and measurement of the absorbance of the solvent layer was proposed by Brown ( I ) and Kurata (2). Since then, the method has been investigated by several workers for various purposes under varying conditions (3-9). It is now regarded as one of the most convenient methods for the microdeterminationof nonionic surfactants which contain polyoxyethylene chains with three or more oxyethylene units. According to Nozawa et al. (3),however, a linear relationship between the concentration of surfactant and the absorbance is not found for pure polyoxyethylenated alcohols and for mixtures prepared by them from two pure compounds of this type. On the other hand, the modification of Buerger's Dragendorff reagent method (IO), in which the concentration of nonionic is measured by the difference in absorbance of a potassium tetraiodobismuth(II1)ate-barium chloride solution after precipitation of the nonionic, gave unreproducible results. This was due to the fact that the absorbance of the reagent
0003-2700/81/0353-0232$01.00/00 1981 Amerlcan Chemical Soclety
