(16) Aden, G. D.; Buseck, P. R. In “Microbeam Analysis-1979”; Newbury, D. E., Ed.; San Francisco Press: San Francisco, 1979; p 254.
(11) Iwanaga, H.; Shibata, N.; Suzuki, K.; Takeuchi, S. Philos. Mag. 1977,35,1213. (12) Yoshiie, T.; Iwanaga, H.; Shibata, N.; Ichihara, M.; Takeuchi, S.Philos. Mag., [Part]A 1979,40, 297. (13) Bradley, J. P., Buseck, P. R., in preparation. (14) Germani, M. S.; Small, M.; Zoller, W. H.; Moyers, J. L. Enuiron. Sei. Technol. 1981,15, 299. (15) Small, M.; Germani, M. S.; Small, A. M.; Zoller, W. H.; Moyers, J. L. Enuiron. Sei. Technol. 1981,15, 293.
Received for reuiew September 11,1980.Accepted April 29,1981. This research was supported by grant ATM-8022849 (toP.R.R.)from the Atmospheric Sciences Division of the National Science Foundation.
Nature of Bonding between Metallic Ions and Algal Cell Walls Ray H. Crist,” Karl Oberholser, Norman Shank, and Ming Nguyen Messiah College, Grantham, Pennsylvania 17027
Introduction
shrimp, Sunda, Engel, and Thuotte (6) found that mortality decreased with increasing salinity and concentration of the chelating agent NTA; Chakoumakos, Russo, and Thurston (7) showed that for various species of copper the toxicity for cutthroat trout was inversely correlated with water hardness and alkalinity. In our laboratory work has been directed to understanding the nature of the initial process, the interaction of metallic ions with an algal cell wall. Earlier it was shown that the adsorption of copper on algae was a reversible system and could be represented by the Langmuir adsorption isotherm (8).Furthermore, metal ions could displace other metal ions or protons, and a reversible pH titration indicated the existence of labile protons whose loss did not lead to substantial structural changes (9). Though the cell wall composition for the alga Vaucheria s. used here is not known, Siegel and Siegel (10)report a protein content of 16-27% for the V a u c h e r i a group. Amino acids in the proteins could provide such functional groups as
Trace elements enter biological systems through the cell walls of plants or the membranes of animals where as constituents of enzymes within the organism they perform many vital functions. A level of availability in agriculture and nutrition is often of concern ( I ) , while toxic effects may appear at various levels of the elements occurring naturally or resulting from waste discharges to the environment. The toxic effect of trace metals on aquatic organisms frequently is dependent on the species of the metallic ion which in turn may be determined by the pH or the varieties of the complexing agents found in natural waters. Interaction with cell walls or with membranes is the initial process in any biotic action. These exterior surfaces have a common composition of proteins and carbohydrates with which the metallic species could react. Bacteria and diatoms are of particular importance because they have a large surface area, are ubiquitous, and are at the low point in the food chain. Situations that emphasize the interaction process are the following: bacteria are the primary agents in the activated sludge treatment of wastewater where they adsorb the trace elements which through sludge disposal as amendments to soil find their way into food crops (2);for the diatom N i t z s c h i a p y r e n o i d o s a , Nielson and Anderson ( 3 ) found that copper influenced the rate of photosynthesis more than for the alga chlorella while the reverse was found for the growth rate, this being attributed to the excreta of the diatom decreasing the concentration of solution copper; Gross, Pugno, and Dugger ( 4 ) found for chlorella that the pigments were affected and photosynthesis was inhibited after short contact with CuS04 at 10-100 pM; for the alga Microcystis a e r u g i n o s a , Allen, Hall, and Brisbin ( 5 )showed zinc toxicity to be due primarily to Zn2+ and Zn(OH)+ in comparison to chelated species; for the larger organism, grass
The polysaccharides of the cell wall could also provide the amino and carboxyl groups as well as the sulfate. The amino and carboxyl groups, the imidazole of histidine, and the nitrogen and oxygen of the peptide bond could be available for characteristic coordination bonding with metallic ions like Cu2+; such bond formation could be accompanied by displacement of protons dependent in part on the extent of protonation as determined by the pH. Metallic ions could also be electrostatically bonded to unprotonated carboxyl oxygen and sulfate. In earlier unpublished work two types of metalalgal bonds were evidenced by the appearance of two slopes in the Langmuir isotherm plots. With these considerations in mind, it was thought that differences in adsorption for transition elements, e.g., Cu2+ and Zn2+,with their strong coordination tendencies as compared to alkali and alkaline-earth elements might show contrasting bonding character. Also, recent work with colloidal inorganic oxides for anion and cation adsorptions and pH titrations (11-13) provides a possible framework for understanding metallic ion-algal systems. Thus, the charge development on increasing the pH, with a zero point charge at pH -3 (14), would suggest similarities of behavior. However, it must be noted that the well-defined and highly crystalline A1203 and S i 0 2 in aqueous systems are bound to show sub-
w Metallic ions adsorbed by algal cell walls at pH 4.5 ranged from 600 to 100 pmolg-l for Cu2+and Na+, respectively, with a pH dependence in the case of Sr2+of -50 pmol g-l per pH unit. A reproducible pH titration was found which required -1000 pmol of NaOH per gram for pH 3-8. Protons displaced by metal adsorption gave the following ratios for H+ displaced/M2+ adsorbed: 1.2 (Cu2+),0.66 (Zn2+),0.59 (Mg2+), 0.30 (Sr2+),and 0 (Na+). These ratios will vary some with concentration and pH. Ion exchange showed a strength of adsorption in the order Cu2+ > Sr2+ > Zn2+ > Mg2+ > Na+ suggesting a trend from probable covalent to ionic charge bonding. This latter was demonstrated directly by Na+ decreasing adsorption with positive metallic ion complexes and i n c r e a s i n g it with negative ones. Ionic charge bonding was thought to arise from a surface charge generated by increasing pH, and covalent bonding from constitutent proteins.
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0013-936X/81/0915-1212$01.25/0 @ 1981 American Chemical Society
stantial differences from the heterogeneous matrix which constitutes the algal cell wall ( 1 5 ) . We report here data on proton displacements, pH titrations, and ion exchanges involving the transition elements Cu and Zn, and alkaline earths Ca, Sr, Ba, and Mg, and the alkali Na. The algal material used was cell wall fragments as these were found to be comparable to whole cells and were more reproducible. In work with fresh cells, no attempt was made to ensure organism vitality as it was thought that normal functions could hardly be maintained in the high salt concentrations and extremes of pH used. The discussion of results is in terms of the covalent and ionic charge bondings stemming from the bonding properties of metallic ions with nitrogen, carboxyl, and sulfate groups. M e t h o d s a n d Materials
The alga used in this investigation was V a u c h e r i a s. obtained from a local limestone spring and held in trays under window light. The cell wall fragments were prepared as follows: cleansing with deionized water a t pH 3.0-3.5 to remove extraneous material and effect partial displacement of naturally adsorbed metallic ions by the protons, maceration with a Thomas tissue grinder, washing by decantation with deionized water a t pH 3.0-3.5, and storage of the moist product in a refrigerator. All data are based on the weight obtained after drying to constant weight (several hours) at 60 OC. Analysis for metallic components was with a Perkin-Elmer atomic adsorption spectrophotometer. For this the dried samples (0.010-0.030 g) were ashed at 450-500 "C, and the ash was treated with several drops of concentrated HN03 followed This procedure for three samples by 5 mL of 0.01 M "03. treated with 0.01 M Cu2+gave an adsorption value of 514 f 5.6% pmol 8-l. As a check three samples from the same treatment after ashing were extracted twice with 0.01 M EDTA, and the ash was then taken up with H N 0 3 as before. This gave 420 f 1.6% or an average of 467 f 10% for the two procedures. Thus a series could have an internal reliability of -5%. For calcium, errors of 10-20% were experienced, particularly for fresh whole material, due in part to contamination from adhering sediment particles and to variability of the amount adsorbed from the natural waters. The water in which Vaucheria s. grew had Ca2+ and Mg2+ concentrations of 2.6 x 10-3 and 1.7 X M, respectively, while fresh whole alga grown in it contained 70 f 20 pmolg-l for Ca and 21.5 & 0.5 for Mg. Because of uncertainties in calcium data, strontium, being very low in most natural circumstances, was selected for the systematic study of an alkaline-earth element. pH titrations were made with a Corning pH meter. Algal samples of 0.10-0.15-g dry weight were suspended in 20 mL of water and stirred under nitrogen. Blanks were run on water and on the several metallic ion solutions used. In each case the data for a run was corrected for the blank which could be from 40% a t low pHs to 10%a t the high values (16). For proton displacements metallic-ion solutions and algal suspensions both of pH 4.5 were mixed, whereupon the pH decreased, indicating release of protons. The pH was brought back to 4.5 with 0.10 NaOH, and the algal sample was removed and analyzed for metal adsorbed. In this way a measure of protons displaced vs. metal adsorbed was obtained. T o determine the competitive relation of metallic ions for available sites, we performed ion-exchange experiments with three sets of binary mixtures; (1)Mg2+, Sr2+, Zn2+, and Cu2+ ions at 0.01 M each with Na+ at 0-0.50 M; (2) Mg2+, Zn2+, and Cu2+ ions at 0.01 M each with Sr2+ at 0-0.20 M; (3) Sr2+a t 0.01 M with Ba2+at 0-0.20 M, the pH throughout being held at 4.5. After equilibration the algal sample was removed and analyzed for the amount of metal adsorbed. For the study of the effect of Naf on adsorption of Cu2+,Zn2+,and Ca2+in the
presence of complexing agents, the reacting mixture was as follows: 20 ml of algal suspension (-0.020-g dry weight) with 0.01 M metallic ion and 0.05 M complexing agent. Leaching experiments with water and with 0.05 M NaCl solution were performed. The metallic ion was first adsorbed on an algal sample (e.g., 0.020 g in 20 mL of 0.01 M Cu2+);the sample was then removed and dried by pressing with filter paper, one half being then extracted with 10-mL aliquots of water and the other half with 0.050 M NaC1. It is noted that no difference within experimental error was found for NaCl or NaN03 in an ion-exchange experiment. Chemicals were obtained as follows. Triethanolamine, malonic acid, acetylacetone, oxaloacetic acid, and a-ketoglutaric acid were from Aldrich Chemical Co.; ethylenediamine and 2,2-bipyridyl were from Mallinckrodt; the salts used, CuS04-5H20, Mg(N03)~-6H20,Zn(N03)2*6H20, Ca(N03)2*4H20,Ba(N03)2,and Sr(N03)2, were of standard CP grade. R e s u l t s a n d Discussion
Cell Wall Fragments vs. Intact Cells. The results were generally of the same character for both. Thus, in an earlier pH titration with intact cells, the NaOH per unit for pH 4.5 was 100 pmolg-l (curve of Figure 1) (9) while for the fragments it was -73 ymol g-1 (curve with no M2+ of Figure 3). Again, in an adsorption experiment with copper, the Langmuir isotherm gave for cells vs. fragments maximum adsorption and 415 X values of 457 X pmolg-l with the corresponding equilibrium constants being 6.5 X 10-4 and 13.3 X Some of the differences here could be due to residue cytoplasmic material in the fragments, to a decrease in the mucilage coating of the cells by the preparation process, or to a difference in surface area exposed. Interaction of Protons. The reaction of protons with hydroxyl ions is shown as a pH titration in Figure 1. The reversibility indicated by the points on the curve was achieved by exercising care in the equilibration process which is somewhat slow at the higher pHs and may be associated with some kind of reversible structural change. However, with one sample the reversal was performed 10 times, indicating a considerable degree of stability. For the pH range of 3-9, -1.00 mL of 0.10 M NaOH was required, which for the 0.10-g algal sample was equivalent to 1000 pmolg-l. Since, with increase of pH, sites become negative by proton loss, the total charge generated would be equivalent to the NaOH used. Effect of pH on Metal-Ion Adsorption. The results are shown in Figure 2 for strontium for the pH range 3.5-6.5, where for each pH there appears to be a saturation level; a t p H 4.5 the Sr2+value is -250 ymol g-l, while those found a t
I
I
I
I
I
I
200
400
600
800
IO00
1200
NaOH. p M o l
I
g-1
Figure 1. Titration of the alga Vauchefia s.;0.01 M NaOH was added to a 0.10-g suspension of the alga in 20 mL of water and stirred under nitrogen. Volume 15,
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this pH for Ca2+,Mg2+,Zn2+,and Cu2+ were 150, 175, 230, and 500, respectively (9). The overall equilibrium is Sr2+
+ (A) s (SrA)
(1)
for which an equilibrium constant can be evaluated by using the Langmuir adsorption isotherm plot. These plots for the various ions had breaks at very low concentrations (in the 10-4-10-5 M region for Cu2+),indicating that a different type of bonding might be involved. However, the data were not sufficiently reproducible to establish this quantitatively. The increasing adsorption of strontium with p H may be explained on the basis of an ionic charge bonding. As noted above, the generation of charge corresponds to the NaOH used to change pH, which from Figure 1and the top curve of Figure 3 is roughly linear. One would expect then that, if ionic charge bonding is occurring, the adsorption would be linear with pH; this is approximately the case for the saturation adsorptions of Sr2+as shown in the figure, namely, -50 pmol g-l per pH unit. The adsorptions of Ca2+and Mg2+ at pH 4.5 were found to be about the same; however, the 2-fold greater adsorption for Cu as noted earlier can be taken as an indication that covalent bonding may come into play. It should be observed that the effect of pH may be explained equally well by a similar pH dependence for complex formation with surface groups as found by Hohl and Stumm (17) for P b adsorption on A1203. Proton Displacement. The ratios mol of H + displaced/mol of M2+adsorbed for Cu2+,Zn2+,Mg2+,and Sr2+were 1.2,0.66, 0.59, and 0.30, respectively, with Na+ showing none. They were all taken at pH 4.5 and ion concentrations of 0.01 M while unpublished data show appreciable differences at other conditions. Though a full range of data would be necessary to develop an adequate understanding, the trend from Cu2+to Sr2+to Na+ lends support for a difference in bonding character.
pH Titrations. These are shown in Figure 3 for data obtained in the presence of Ca2+,Zn2+,and Cu2+.Ca2+is seen to give a moderate shift in the normal curve which for Cu2+ was much greater even at 0.002 M. From the titration of the metallic-ion solutions alone, incipient precipitation for 0.01 Cu2+was found to occur at pH 5.0, while for 0.01 Zn2+it was at pH 6.5. The titration in the presence of Cu2+may be described as follows. A Cu2+ solution and an algal suspension both a t pH 3.5 are mixed, whereupon the pH drops; it is then brought back to pH 3.5 and the titration is performed as indicated in the figure. Subsequent addition of NaOH increases the pH, which at the same time enables increased Cu2+ adsorption accompanied by release of protons for reaction with NaOH, the overall effect being a shift of the curve to the right. For low concentrations this increasing adsorption effects a considerable reduction in the metallic ion; it was found that for 0.002 M Cu2+ its concentration dropped by half on going from pH 3.5 to 4.5. The effect of Cu2+is pronounced as compared to Zn and Ca2+; here Na+ as with proton displacement has little effect indicating a weak covalent boriding. Ion Exchange. The data for the Zn2+-Na2+and Zn2+-Sr2+ systems are shown in Figure 4. Here, for example, NaN03 was added to 0.01 M Zn2+ solutions containing algal suspensions to give Na+ concentrations of 0,0.020,0.050,0.10,0.20,and 0.50 M. With increasing Na+ the Zn2+ adsorbed (Zn(Na)) decreases while adsorbed Na+ (Na) increases. As can be seen, the effect is much greater when Sr2+ is used. For Na+ displacing adsorbed Zn2+,we have the following system with an equilibrium constant of (7.7 f 0.4) X
Na+
+ (ZnA) 5 (NaA) + Zn2+
(2)
K=--(NaA) uznz+
(3) (ZnA) U N ~ + The corresponding constant for Na+ displacing adsorbed Sr2+ is (5.4 f 0.7) X 10-2 with the larger value showing that adsorbed Sr2+ can be displaced easier by Na+. Ease of displacement, as reflected by larger constants, is shown in Table I for the following systems: (1) Na+ displacing Mg2+, Zn2+, Sr2+,and Cu2+,(2) Sr2+displacing Mg2+,Zn2+,and Cu2+,(3) Ba2+displacing Sr2+.These ion-exchange constants give the interesting result that the strength of adsorption of metal ions to algae follows the series Cu > Sr > Zn > Mg (by displacement of Na+ experiments) and Cu > Zn > Mg (by displacement of Sr2+ experiments). Inspection of formation constants for these ions shows a trend of C U ~>+ Sr2+ > Zn2+ > Mg2+ > Ba2+ > Na2+ (very
I
I L
I
100
I
300
I
400
YdOH,fMol g‘l 4
8
Sr”,
I2
M L-’X
16
20
IO3
Figure 2. Adsorption of Sr2+ vs. Sr2+ ion concentration for various pHs. The “saturation” levels increase -50 pmol per pH unit.
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Environmental Science & Technology
Figure 3. Effect of metallic ions on the pH titration. A 0.10-g suspension of the alga with the metallic ion at the concentration indicated was titrated with 0.10 M NaOH. The curves are for data corrected for blank titrations of the metallic-ion solutions.
I
Cu Triethanolamine
Cu Acetyl A
6
Acetonate
I
IW I
1
010
010
I 010 Sr(N01)2,
I
I
040
050
Na N01. M L’ Figure 4. Ion exchange on algal surface, Curves Zn(Na)and Na: mixture of 0.02 g of alga, 0.01 M Zn2+, and 0-0.5 M Na+; the Zn2+ adsorbed decreases and Na’ adsorbed increases. Curves Zn(Sr) and Zn: 0.01 M Zn2+ and 0-0.2 M Sr2+;here stronger adsorption of Sr2+effects a greater decrease in Zn2+ adsorption.
0.10
0 20
0.30
0 50
0.40
M I-’ Figure 5. Adsorption of copper in the presence of complexing agents which form positive ion complexes. The Cu2+ adsorption decreases with increasing concentration of the positive Na+ ion; the increase in Na+ adsorption shown for ethylenediamine (curve labeled Na) is comparable to this for the other complex ion systems. NaNOj,
Table 1. Ion-Exchange Constants for Na+, Sr2+, and Ba2+ Displacing Adsorbed Metal Ions a adsorbed metal Ion
system I b
Mg2+
(8.1 f 0.2) x 10-3 (7.7 f 0.4) X (3.7 f 0.3) x 10-3 (2.6f 0.4) X
Zn2+
Sr2+ cu2+
system 2
system 3
*
1.12 f 0.20 0.54 f 0.07
1.41 f 0.08 0.13 f 0.03
a The amount of metal adsorbed was relatively small so the initial make-up ion concentrations were used in the calculations of activities. K = [(NaA)/ (MA)](a~z+/ab+). K = [(SrA)/(MAXaMz+/a~z+). K = [(BaA)/(srA)](a~z+/
aaa2+).
-g
M
low). Thus, we have presumptive evidence that covalent bonding may be playing a role in ion exchange. Significantly, however, the fact that Na+ is also effective in displacing Cu2+ would seem to provide a strong argument that ionic charge bonding is also an important factor. Interestingly, Na+ appears in the equilibrium as first order rather than second as found in conventional ion-exchange systems. This result is probably reliable since there was no trend in the values of the constants for an individual system and the same displacement order was given for both the Na+ and Sr2+series. An explanation may lie in a differing surface structure for algae as compared to colloidal inorganic oxides and exchange resins. The latter have a pattern of closely spaced negative sites, while algae cell walls, being composites of proteins and polysaccharides, provide the possibility that reactive sites are interspersed with inert atoms. In such case an isolated negative charge could attach any positive ion, thus accounting for first-order behavior for Na+ as well as for bivalent ions. Ionic Charge vs. Covalent Bonding in Metal Complexes. Early experiments with copper showed an extent of adsorption which was about the same whether it was present
200
?\ 71 Q SI 5
P
5
100
010
020
Na
Ca
Citrate
030
040
050
N Q , M L-1
Figure 6. Adsorption of metal ions in the presence of complexing agents which form negative or neutral complexes. The adsorption is enhanced with increasing Na+ ion concentration; the Na+ adsorption shown for the EDTA complex (curve labeled Na) is comparable to this for the other complex ion systems. Ion species are shown in Table II.
as the free aquo ion or primarily as the ethanolamine complex. This could be due to (a) the copper-alga bond being stronger than that with triethanolamine, (b) the formation of a direct Volume 15, Number
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October 1981 1215
Table II. Speces Distribution for Metallic Ion Complexes at pH 7 a ligand (L)
acid form cons1 ( K )
[L], mol L-1
x
triethanolamine ((CHzCHOHMJ)
K1
8.0
3.8
acetylacetone (CH3COCH2COCH3)'
Ki
8.82
0.05
10-3
( c u ~= + 2.2 4.40
0.05
10.0 7.3 3.69 2.48 5.68 4.35 2.87 5.34 2.85
3.43
x
x
10-4 M)
10-5
0.05 (ZnZf = 3.8 X 0.0335
form cons1 (6)
complex
M)
x
9.48 10-4 (Cu2+ = 1.3 X 10-5 M)
wL.
CUL2+ CUL22+
3.9
0.66 0.32
CUL2+ CUL22+
8.16 14.76
-1
CUL'+ CUL22+ CUL2+ CUL22+ ZnLO ZnL22CaH2LCaHL2-
8.1 13.5 10.55 19.88 3.1 4.9 3.5 8.4
7.97 x 10-5 -1.0 1.37 10-5 -1.0 0.20 0.79 -1.0 2.6 10-4
CULO cuL22-
4.42 7.20
6.0
5.03 X
.o
x
x
0.21 0.79
a Metallic ions were 0.01 M and complexingagents 0.05 with pH 7 except for malonic acid which was at pH 4.5. Adequate data for a-ketoglutaric and oxaloacetic acids were not available; however, the acid constants are low and the complexes would be neutral or negative. Cu2+ and Zn2+ concentrations are given for triethanolamine, tartaric acid, and malonic acid; Mt was very low for the other systems. The formation constants given as the logarithms for the acids (K,) and the complexes (PI) were taken from Sillen and Martel (22)and Ringbom (23).[L], the concentration of the ligand, was calculated by using the method in Chapter 7 of ref 24. Since the amount of a complex is relatively small, the concentrationof L is determined by the pH. a is the fraction of the total metal in a particular complex. It is given by ~ M L , = [MLn]/[MT] = Pn(L)"/(l 4- 2,=,"p,(L)'). For details, see Chapter 9 of ref 24 and ref 20. Hydrogen ions do not prevent complexing.
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i--: Ratio of Complex Species
a-
I
Adsorption of Zn
Sr H20
1 I
5
6
PH
Figure 7. The zinc tartrate system: (A) ratio of ion species; (B) adsorption of Zn2+ at pH 4-7 without and with 0.50 M NaN03. The adsorption is decreased by NaN03 at pH 4 and is increased above pH 4.5. The negative to neutral ion ratio rises from ca. 2 to 3.2 over the range.
ligand-alga bond, or (c) the attraction of the positive complex ion to the negative algal surface (16,18-21). An observation favoring c was that adsorption in the presence of NaN03 gave either positive or negative effects depending on the sign of the charge on the metallic ion complex. A study was made of this latter effect for a variety of positive and negative complex ions. The total metallic-ion concentration was 0.01 M. For the positively charged complexes of Cu2+,added Na+ decreases the extent of adsorption of Cu2+ as shown in Figure 5 for data a t pH 7. The species distribution for these positive complexes with different degrees of complexation is given in Table 11.On the other hand, the presence of Na+ increases the extent of adsorption of Cu2+ when complexed with ligands having negative carboxyl groups as shown in Figure 6. In this 1216
Environmental Science & Technology
I I
2
3
4
10 M L ALIQUOT S OF H20 OR NPCl Figure 8. Leaching of metal from alga by water and 0.05 M NaN03. An alga sample of 0.10 g equilibrated with 0.01 M metal ion was dried with filter paper and then leached successively with 10-mL aliquots.
latter case some neutral complexes are also present at pH 7 (see Table 11). The Zn-tartrate system was used for a study over a wider pH range and involving a neutral and negative complex as a test of the above conclusions. Solutions were 0,010 M in Zn2+, 0.05 M in tartrate, and 0.5 M in NaN03; the speciation of Zn2+ complexes is shown in Table I1 for pH 7, and ratios of negative to neutral complexes for other pH values are shown in Figure 7A. I t was found that Na+ increases the adsorption of Zn2+ in the 5-7 pH range but causes a decrease at pH 4 (Figure 7B). This behavior is in agreement with the previous discussion on
Literature Cited
(2) Bradford, G. R., et al. J . Enuiron. Qual. 1975,4,120. (3) Nielson, S. E.; Wium-Anderson, S. Physiol. Plant 1971, 24, 480. (4) Gross, R. E.; Pugno, P.; Dugger, W. M. Plant Physiol. 1970,46, 183. (5) Allen, H. E.; Hall, R. H.; Brisbin, T. D. Enuiron. Sci. Technol. 1980,14,441. (6) Sunda, W. G.; Engel, D. W.; Thuotte, R. M. Enuiron. Sci. Technol. 1978,12,409. (7) Chakoumakos, C.; Russo, R. C.; Thurston, R. V. Enuiron. Sei. Technol. 1979,13,213. (8) Frey, R. A,; Crist, R. H.; Oberholser, K. M. Proc. Pa. Acad. Sei. 1977,52,179. (9) Sailer, D.; Shellenberger, D.; Crist, R. H.; Oberholser, K. M. Proc. Pa. Acad. Sei. 1980,54, 85. (10) Siegel, B. Z.; Siegel, S. M. CRC Crit. Reu. Microbiol. 1973,lO. (11) Huang, C.-P.; Stumm, W. J . Colloid lnterfuce Sei. 1973, 43, 409. (12) Davis, J. A.; James, R. 0.;Leckie, James 0. J . Colloid Interface Sei. 1978,63,480. (13) Davis, J. A.; Leckie, J. 0. J . Colloid Interface Sei. 1980, 74, 32. (14) Stumm, W.; Morgan, J. J. “Aquatic Chemistry”; Interscience: New York, 1970; p 455. (15) Dodson, J. K., Jr.; Aronson, J. M. Bot Mar. 1978,21,241. (16) Elliot, H. A,; Huang, C.-P. Enuiron. Scz Technol. 1980, 14, 87. I (17) Hohl, H.; Stumm, W. J Colloid Interface Sei. 1976,55,281. (18) Davis, J. A.; Leckie, J. 0. J . Colloid Interface scz. 1978, 67, 90. (19) Davis, J. A,; Leckie, J. 0. Enuiron Sei. Technol. 1978, 12, 1309. (20) Elliot, H. A.; Huang, C.P. J . Colloid Interface Sei. 1979, 70, 29. (21) Davis, J. A.; Leckie, J. 0. Enuiron Sei. Technol. 1979, 1 3 , 1290. (22) S i l l h , L. B.; Martell, A. E. Spec. Publ.-Chem. Soc. 1971, No. 25, Supplement No. 1. (23) Ringbom, A. “Complexation in Analytical Chemistry”; Interscience: New York, 1979; Appendix. (24) Butler, J. N. “Ionic Equilibrium, a Mathematical Approach”; Addison-Wesley: Reading, MA, 1964; Chapter 7.
(1) Mortvedt, J. J.; Giordano, P. M.; Lindsay, W. L. “Micronutrients in Agriculture”; Soil Science Society of America: Madison, WI, 1972.
Received for review September 30,1980. Revised manuscript received May 18,1981. Accepted June 15,1981.
the effect of Na+ on the adsorption of complexes of unlike charge. The species distribution values, a, for pH 7 are listed in Table 11. Results of the leaching experiments performed to explore further the bonding character are shown in Figure 8. Here, aqueous NaCl is more effective than water, as expected. Copper is retained in both cases while strontium is almost completely removed by NaC1, with zinc showing intermediate behavior. Conclusion
It has been demonstrated that the alga V u u c h e r i a s. has a proton equivalence of -1000 pmol g-l, that metallic ion adsorptions range from -500 pmol g-1 for Cu2+to 100 for Na+, that metals displace each other in the order Cu > Sr > Zn > Mg > Na and they also displace protons, and that Na+ decreases adsorption of positive metallic ion complexes and enhances negative complexes. Most of these observations can be understood in the light of the protein and polysaccharide composition of the algal cell wall where one would expect covalent bonding amino and carbonyl groups and ionic charge bonding carboxyls and sulfates. Cu2+and Na+ then represent extremes in the bondings to these two types of functions. As found by Liecke, Davis, Stumm, and co-workers (22,17,18, 21) for inorganic colloids, bonding of a complexing agent directly to the surface with its attached metal must also be considered here. This could provide an alternate explanation for the observation that the aquo copper ion and the triethanolamine copper complex are about equally adsorbed. Acknowledgment
We acknowledge the help of Mr. Charles Zercher with the experiments on proton displacements and p H titrations.
Identification and Determination of Individual Tetraalkyllead Species in Air Walter R. A. De Jonghe, Dipankar Chakraborti,?and Fred C. Adams+ Department of Chemistry, University of Antwerp (U.I.A,), B-2610 Wilrijk, Belgium
Introduction
Several investigations have been undertaken in recent years to determine the contribution of gaseous tetraalkyllead compounds to the lead burden of our environment (1-5). One major focus of attention is the occurrence and fate of these pollutants in ambient air. As automotive emissions are the main source for both organic and inorganic atmospheric lead, lead levels rise with increasing traffic density. In urban areas, tetraalkyllead (TAL) compounds are generally believed to be present a t a concentration in the range 10-200 ng of P b m-3, which represents typically 1-10% of the total lead concentration (6-13). These data do not suffice for a full assessment of the health hazard associated with airborne alkylleads in view of the dissimilar toxic properties of different TAL species (14,15).
T o date, no comprehensive surveys have been performed regarding the discrimination of atmospheric TAL concenPresent address: Department of Chemistry, Texas A&M University, College Station, TX 77843. +
0013-936X/81/0915-1217$01 25/0 @ 1981 American Chemical Society
trations into specific compounds. The available information is restricted to a few rather occasional measurements, intended to demonstrate the practical applicability of newly developed analytical techniques. Laveskog ( 1 6 ) determined tetramethyllead (TML) and tetraethyllead (TEL) separately in automobile exhaust gases and found a dependence of the concentration ratio TML/TEL on the driving mode of the car. Data on the other tetraalkyllead species were not reported. Similarly, some determinations of TML and TEL in city air were recently reported (17). In the air in the vicinity of a highway, Corrin (18) detected only TML, whereas, in air samples taken close to a car-repair shop, Rohbock et al. (13) found TML and TEL. The low analytical sensitivity of their method probably did not allow them to detect trace quantities of the lead alkyls with mixed ethyl-methyl groups, Le., TMEL (trimethylethyllead), DMDEL (dimethyldiethyllead), and MTEL (methyltriethyllead).The presence of these alkyllead species in city air, however, has been demonstrated by other workers (19,20).Furthermore, it was observed that all TAL compounds mentioned can be found in the atmosphere even Volume 15, Number 10, October 1981 1217
