the less abundant elements which Were determined near their detection limits. While accuracy can never be absolutely known, abundances determined for the standard rocks using this method exhibit relatively good agreement with abundances obtained by other investigators. Limitations. The removal of barium as a barium sulfate precipitate is not a completely satisfactory procedure. In particular, if substantial amounts of barium (greater than several per cent) are present in the initial sample, the precipitation of the sulfate results in the coprecipitation of a significant fraction of the total rare-earth elements, yttrium, and scandium. This one step accounts for a large portion of the chemical loss during the analytical procedure. Several other methods were investigated-Le., the solvent extraction method suggested by Schnetzler et al. ( 4 ) and an ion exchange procedure using nitric acid as the eluent. These procedures proved to be even less satisfactory than the precipitation of barium as a sulfate. An alternative procedure for the removal of barium would be desirable. From a consideration of the inter-element interferences, it is evident that more corrections are required for the heavier rare-earth elements. This is of particular importance since most analyses involve mixtures that are enriched in the light rare-earth elements. The greater the amount of light rareearth enrichment, the larger the correction factors which must be applied to the heavier elements. The result is a (15) F. J. Flanagan, Gcwcliim. Cosntocl~int.Acta, 33,81 (1969). (16) G. H. Morrison, J. T. Gerard. A. Travesi, R. L. Curie, S. F. Patterson, and N. M.Potter, ANAL.CHEM., 41, 1633 (1969). (17) H. Higuchi, K. Tomura, N. Onuma, and H. Hamaguchi, Geochem. J., 3, 171 (1969). (18) L. A. Haskin. R. 0. Allen, 1’. A. Helmke, T. P. Paster, h4. R. Anderson, R . L. Korotev, and K. A. Zweifel, “Proceedings cf the Apollo I 1 Lunar Science Conference. Vol. 2,” Pergiimon Press, New Y ork, N. Y ., 1970, p 1213. (19) J. A. Philpotts and C. C. Schnetzler, ibid., p 1471.
reduction in the precision and accuracy of these determinations. The method is also somewhat limited by the amount of material which can be collected on the ion exchange paper. This limitation is required in order to keep the absorption effects to a minimum. Under the particular experimental conditions described above, 500 pg was the maximum allowable amount. This amount changes with changes in the total surface area of the ion exchange paper. For mixtures that are significantly enriched in the light or heavy rare-earth elements, the result is a rather high detection limit for some of the less abundant elements. This limitation may be alleviated if the mass absorption coefficients are known with sufficient accuracy to permit the calculation of corrected intensities by iteration procedures. The ultimate sensitivity of the method will be determined by the statistical capability of distinguishing a weak line from a high background and by the number of corrections which must be applied to any particular analytical line. The latter limitation can be significantly reduced by using an X-ray spectrometer with high resolution capabilities. ACKNOWLEDGMENT
W. H. Pinson and F. A. Frey, Department of Earth and Planetary Sciences, M.I.T., provided laboratory space and equipment and also offered advice on various aspects of the analytical procedure. RECEIVED for review March 14, 1972. Accepted May 23, 1972. This paper represents a portion of a doctoral thesis submitted to the Department of Geology, Boston University. The study was supported, in part, by a grant from the Boston University Graduate School to A. H. Brownlow and a Penrose research grant from the Geological Society of America to the author.
Determination of Trace Amounts of Iron(l1) Using ChemiIuminescence Anidysis W. Rudolf Seitz National Encironinental Research Center, Coroallis Environmental Protection Agency, Southeast Water Laboratory, Athens, Ga. David M. Hercules Department of Chemistry, Unioersity of Georgia, Athens, Ga. Trace amounts of Fetll) are determined by measuring Fe(l1)-catalyzed light emission from luminol axidation by oxygen. Fe(l1) is the only common metal ion to catalyze the lurninol reaction in aqueous solutions in the presence of oxygen alone. The detection limit for Fe(ll) is 0.005 pg/l. The minimum detectable quantity is 5 picograms. Response is linear up to 50 lig/l. High concentrations of organic ligands reduce the intensity of light catalyzed by Fe(ll) but do not affect linearity of response. Excess quantities of Cu(ll), Mn(ll), Co(ll), Cr(lll), and Ni(ll) reduce the light intensity and affect linearity. Chemiluminescence analysis for total iron in natural water samples and standard orchard leaves agreed with values obtained by other methods.
been reported (1-4). In this paper we report a method for Fe(I1) in natural waters based on Fe(I1) catalysis of aqueous lurninol (5-amino-2,3-dihydro-1,4-phthalazinedione)oxidation in the presence of dissolved oxygen. Other catalysts for luminol chemiluminescence either require peroxide. or do not commonly occur in natural waters where Fe(l1) is present. The previously reported chemiluminescence method for Fe required H202and, therefore, was subject to interference by other metals.
TRACEMETAL ANALYSIS based on catalysis of luminol chemiluminescence is fast, sensitive, and does not require expensive instrumentation. Methods for Cu, Fe, Co, and Cr have
(1964). (3) A. K. Babko and I. E. Kalinichenko, /hid., 31, 1316 (1965). (4) W. R. Seitz, W. W. Suydam, and D. M. Hercules. AVAL. CHEM.,44,957 (1972).
(1) A. K. Babko and L. I. Dubovenko. Zacad. L d . , 30, 1325 (1964).
(2) A. K. Babko and N. M. Lukovskaya, Ukr. Kliim. Zh.. 30, 388
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Figure 1. Diagram of flow pattern in cell Dark arrows show the principal directions of flow, and the dashed line shows the path a slug of sample follows while it mixes in the cell
I n the absence of oxygen, Fe(I1) is quite soluble in natural waters (I& Fe(OH)2 = Ksp Fe(C03) = 10-10.5). Oxygen rapidly oxidizes Fe(I1) to Fe(III), which forms an insoluble hydroxide. Morgan and Stumm (5) have presented a complete account of iron chemistry in natural waters. If the iron concentration in a water supply becomes too high, several problems can occur: Brown stains on plumbing fixtures and cooking utensils, discoloration of food and clothing, formation of iron oxide scale on surfaces contacting the water, unpleasant taste, and the establishment of iron bacteria, which can cause clogging of mains and unpleasant odors (6). To monitor the iron system in lakes and ponds, an analytical method for measuring Fe(I1) specifically is desirable. Although colorimetric methods for Fe(I1) using 1,lO-phenanthroline and its derivatives have been reported and critically evaluated (7-12), these methods have drawbacks. Chemical treatment of a sample may change the Fe(I1)--Fe(III) ratio as in reference 12, where acidification followed by digestion caused organic matter to reduce Fe(II1) to Fe(I1) resulting in erroneously high values for Fe(I1). Shapiro (8) pointed out the error. Since different variations of the colorimetric method do not always give the same value for Fe(I1) concentration, accuracy of this approach is questionable. The chemiluminescence method requires no chemical operations on the sample. It is faster and much more sensitive than standard colorimetric methods. ( 5 ) J. J. Morgan and W. Stumm, Proceedings of the Second Inter-
national Water Pollution Research Conference, Tokyo, 1964. (6) J. E. McKee and H. W. Wolf, “Water Quality Criteria,” 2nd ed.. California State Water Quality Control Board, Pub. No. 3A (1963). (7) G. F. Lee and W. Stumm, J . Amer. Writer Works Ass., 59, 897 (1960). (8) J. Shapiro, Limnol. Oceuuogr., 11, 293 (1966). (9) M. M. Ghosh, J. T. O’Connor, and R. S. Englebrecht, J . Amer. Water Works Ass., 59, 897 (1967). (10) J. W. McMahon, Limnol. Oceamgr., 12,437 (1967). ( I 1) J. W. McMahon, Wnrer Res., 3, 743 (1969). (12) E. T. Gjessing. L i m ~ o lOcrumgr., . 9, 272 (1964). 2144
0
Log
-8
F e ( l l ) Concentration
7
(Molar)
Figure 2. Log Fe(I1) peak height DS. Log Fe(I1) concentration pH 10.5-11.0; 02,80 cm3 minute; 4 X 10-4Mluminol. The ranges 2-10 X 10-lOM Fe(II), 2-10 X 10-9M Fe(II), and 2-10 X lO-*MFe(II) were all run separately. They are put togefher for compactness of presentation. The relative units for Figure 2 are the same for Figures 3-8,lO EXPERIMENTAL
Apparatus. The chemiluminescence apparatus has been described elsewhere ( 4 ) . A few modifications, such as omitting the peroxide syringe, were necessary to meet the requirements of a different system. The reaction is carried out in a flow system driven by a Harvard Apparatus Model 975 infusion pump. The reactants are continuously mixed in a cell positioned directly in front of a photomultiplier tube, which measures light emission. Luminol is dissolved in a 0.1M KOH-H3B03buffer that controls the pH of the light-emitting reaction. In the cell, luminol is mixed with deionized water. Slugs of sample are inserted into the water line using a Chromatronix SV-8031 sample injection valve. As the sample passes through the cell at steady state, a constant level of light emission is observed. When the sample is consumed, the light level returns to zero. Unless otherwise mentioned, a sample volume of 2 ml and a flow rate of 4.41 ml/minute/syringe were used in this study. The cell is a 2 cm high x 12-mm 0.d. borosilicate glass tube with three 3-mm 0.d. glass nipples on the bottom (Figure 1). The luminol-buffer mixture flows through one nipple, the sample-background through the second, and the stirring gas through the third. Teflon (Du Pont) tubing is fitted inside the glass nipples so that mixing occurs toward the center of the cell. There is a 3-mm 0.d. exit tube at the top of the cell. Cell volume is approximately 1 cm3. In reference 4 , nitrogen bubbling was used to mix the reactants in the cell. In the work with Fe(II), oxygen is the stirring gas as well as one of the reactants. When the stirring gas also serves as a reactant, control of the gas flow rate is particularly critical. Two needle valves in addition to the regulator valve are used to control gas flow. Fluctuations in mixing still produce noise which can amount to 5-10z of peak intensity and must be averaged out in measuring the peak height. Light intensities were measured by a 1P21 photomultiplier operated at 800 volts using a Princeton Applied Research
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
(PAR) Model 280 high voltage power supply. The photomultiplier tube was housed in a PAR Model 180 housing, and the anode current was measured with a PAR Model 210 D C Photometer and recorded on a Hewlett-Packard-Moseley 7100 B strip-chart recorder. Chemicals. Luminol from Eastman Organic Chemicals was converted to the sodium salt and was purified by recrystallization from basic aqueous solution. The purified sodium luminol was dissolved in 0.1M H3B03KOH buffer to control the pH in the reaction cell. The H 3 B 0 3 concentration was maintained constant while the amount of KOH was varied to achieve the desired pH. All reagents were prepared using water from a Continental Water Conditioning Company deionization system. A 0.100M standard Fe(I1) solution was prepared by weighing reagent grade FeS04 and dissolving in acidic solution. Other standards were prepared by dilution. Fe(I1) standards of concentrations below 10-3M were prepared fresh daily. All standards were prepared in 1 M HC1 to stabilize Fe(I1) against air oxidation. Procedures. A 500-ml sample volume was used in this study. Such a large sample is convenient because the amount of sample consumed for one measurement (2 nil) is small relative to the total sample volume. Analysis can be done by the method of standard additions (adding known quantities of iron to the unknown sample) without having to correct for volume changes. This corrects for matrix effects which are likely to be encountered in natural samples. Smaller sample volumes can be used. Standard additions were made with 50- or 10O-pl Grunbaum pipets. All Fe(I1) samples were purged with high purity nitrogen scrubbed by a vanadous solution before being introduced into the flow system.
9 .c
2000-
-d
1600-
E01
1200-
.-
f 800-
2 iL
*
0
4
"
8
x
12
O 16
20
L 24
M Fe ( I I)
Figure 3. Effect of oxygen flow rate on chemiluminescence us. concentration curves at high Fe(I1) flow rates A-A-A c--3-n -4
30 cm3 02/minute 60 cm3 02/minute
100 cm3 Oz/minute
- - - - - linear response
pH 10.5-11.0, 4 X 10-4M luminol
RESULTS AND DISCUSSION
Effect of Cell Configuration. In basic solution, oxygen rapidly oxidizes Fe(I1) to ferric hydroxide. Although the mechanism of aqueous luminol oxidation is not known, it is likely that an intermediate produced in the oxidation of Fe(I1) by dissolved oxygen is the species reacting with luminol. It was observed that the geometry of the reaction cell affected chemiluminescence us. Fe(I1) concentration curves. This effect was interpreted by assuming that the Fe(II)-02reaction was the first step of the luminol oxidation. A preliminary study of flow patterns in the cell was undertaken to optimize cell geometry. The positions of the Teflon tubes through which the reactants and the oxygen enter the cell are adjustable. The study of flow patterns was done by reducing potassium permanganate with hydrogen peroxide in the cell using various geometries and observing the path of purple permanganate from its entrance into the cell until it changed color. The predominant flow pattern is illustrated in Figure 1. The rising oxygen bubbles drive liquid toward the top of the cell, producing a downward return flow on the side of the cell away from the bubbling. The Fe(I1) sample solution is best introduced into the downward flow. (Refer to Figure 1.) As the sample enters the cell, it has sufficient momentum from the injection system to start upward before being forced downward by the flow pattern. If the sample goes into the upward flow as soon as it enters the cell, three effects are observed. First, at low Fe(I1) concentrations (2-10 x l0-9M) calibrations of light intensity us. concentration curve upward rather than being linear. Some of the Fe(I1) is lost without catalyzing light emission. The percentage of Fe(I1) lost is greater at lower concentrations. Since to catalyze the luminescent reaction, Fe(I1) must interact
x 1 V 7 M Fe(lI)
Figure 4. Effect of syringe drive rate on chemiluminescence us. concentration curves at high Fe(I1) concentrations A-A-A C-a-D 0-0-0
Syringe drive rate 2.25 ml/min/syringe 4.41 ml/min/syringe 8.64 ml/min/syringe
pH 10.5-11.0; 02,80 cm3/minute; 4 X lO-'M luminol
with oxygen in the dissolved rather than gaseous state, the nonlinearity may be due to Fe(I1) oxidation by 0 2 bubbles before mixing becomes complete. The cell configuration shown in Figure 1 corrects the situation. When injected into the downward flow, the sample mixes completely before contacting oxygen bubbles. The second effect occurs at higher Fe(I1) concentrations. Some of the liquid swept upward by the bubbles leaves the cell through the exit tube. If the sample enters the upward flow, some of it will leave the cell very shortly after it enters, reducing the time available for the Fe(I1) to react with 0 2 . In the region of linear response, all the Fe(1I) is oxidized before leaving the cell. When unreacted Fe(I1) leaves the cell, light emission us. concentration is not linear. Injecting the sample into the downward flow reduces loss of unreacted Fe(I1) from the cell and makes it possible to get linear chemiluminescence us. Fe(I1) concentrations response up to higher concentrations. Third, when Fe(I1) enters the upward flow in the cell, the level of light emission becomes more sensitive to small fluctuations in flow rate, and the noise level becomes much greater.
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-
I
5
200
B
I
I
e
,
,
,
-4
Log
-3 Luminol Concentration
-2
Figure 5. Effect of luminol concentration on Fe(I1) peak height pH 10.5-11.0; 02,80 cm3 minute; lO-7M Fe(I1)
V
I
I
I
I
I
I
0
2
4
6
8
IO
12
14
I
I
16
18
Flow rate (mis/min per syringe)
Figure 7. Effect of syringe drive rate on Fe(I1) peak height pH 10.5-11.0; 02,80 cm3/minute; 2 X lO-*MFe(II); 4 X 10+M luminol
-( . ‘ o = -.
9
10
II
12
PH
Figure 6. Effect of pH on Fe(I1) peak height 02, 80 cm3/minute; lO-’M Fe(I1); 10-3M luminol
Response to Fe(I1) Concentration. Figure 2 plots light intensity catalyzed by Fe(I1) as a function of concentration to lO-’M Fe(I1). Response is linear throughfrom 2 X out this range. Because no light emission is observed when the luminolbuffer is mixed with deionized water, the noise due t o photomultiplier dark current theoretically limits sensitivity. A detection limit (signal equal t o twice the noise level) can be extrapolated from light intensities at higher concentrations to be less than 10-llM. However, measurements at this level are subject to trace metal adsorption-desorption effects, “memory” in the sample bottle, and difficulty in making reproducible standard additions. Measurements in the 10-10-10-9M concentration range required several hours conditioning of the flow system before satisfactory results were obtained. In this range a faster flow rate (8.64 ml/ minute/syringe) and a longer sample loop (5-ml volume) were used to improve the quality of the data. A Fe(I1) conentration of 10-9M may be considered a “practical detection limit.’’ At concentrations greater than lO-$M, conditioning of the flow system was occasionally required, especially after the apparatus had been unused for several days. Conditioning was achieved by running successive slugs of lO-7M Fe(I1) through the cell until the emitted light reached a constant value. When more drastic cleaning was necessary (e.g., after running permanganate through the flow system), 1 M HCI was run through the flow system. The upper limit of linear response to Fe(I1) depends on how efficiently the sample is mixed with oxygen before it leaves the cell. Because the Fe(I1) sample is purged with 2146
nitrogen before being introduced into the flow system and the luminol-buffer is saturated with air ( 2 0 x oxygen), most of the oxygen has to come from the oxygen bubbling through the cell. The rate at which oxygen dissolves in the cell solution varies with the rate of oxygen flow through the cell. Figure 3 shows how the upper limit of linear response t o Fe(I1) concentration varies with oxygen flow rate. The upper limit of linear response can be extended by reducing the syringe drive rate so that the Fe(I1) sample spends more time in the cell, allowing more time for the oxygen to dissolve and react with it. Figure 4 shows Fe(I1) calibrations at three different flow rates. At flow rates slower than 4.41 ml/minute/syringe, there is a high level of noise on the signal from the emitted light, approximately 15 of peak height at 2.25 ml/min/syringe and 25% a t 1.15 ml/ min/syringe. Dilution rather than decreasing the flow rates is a better way t o get a “concentrated” sample into the linear range. Effect of Luminol Concentration. Figure 5 shows the intensity of light catalyzed by lO-7M Fe(I1) as a function of luminol concentration. Although the reaction is most efficient at 5 x 10-3M luminol, 4 x lO-*M luminol was chosen as the concentration for the present study in order to conserve luminol. The decrease in light at the highest luminol concentrations may be due to Fe(I1) complexing with lurninol or with the aminophthalate product of luminol oxidation, as discussed below. Effect of pH. Figure 6 shows the intensity of light catalyzed by lO-’M Fe(I1) as a function of pH. Light emission does not vary significantly from pH 10.5 to 11.0, unlike systems involving hydrogen peroxide which show a continuous change with pH. Since p H control is not critical, the Fe(f1)oxygen-luminol system is potentially useful for systems where analysis can be done using a calibration curve rather than the method of standard additions. Effect of Flow Rate. Figure 7 plots light emission catalyzed by 2 x 10-8M Fe(I1) as a function of sample flow rate. As discussed in reference 4 , light emission is proportional to flow rate only if the reaction is going to completion within the cell. From the light emission us. flow rate curve, one may infer that the reaction goes to completion is less than a second. Effect of Organic Complexation. Formation of Fe(I1)organic complexes reduces the availability of Fe(I1) for catalysis and therefore reduces the level of light observed. Figure 8 plots light intensity as a function of ligand concentration for several ligands. The oxygen-containing ligands,
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
t2
+I
-6
-5
-4
-3
-2
’
Time
-I
Log Ligand Concentration
Figure 9. Explanation of decaying Fe(I1) peaks in the presence of organic ligands. Before t , deionized water is going through cell. From ti to tz, a slug of sample goes through the cell
Figure 8. Effect of ligand concentration on Fe(I1)-catalyzed chemiluminescence A-A-A E--C--CI
0-0-0
C--@--O
Ethylenediamine Glycine Tartrate 2,4Pentanedione
pH 10.5-11.0; 02,80 cm3 /minute; lO-’M Fe(I1); 4 X 10-PMluminol
tartrate and 2,4-pentanedione, are more efficient in reducing light catalyzed by Fe(1I) than the nitrogen-containing ligand ethylene diamine. At high ligand concentrations, the observed light “peak” is not level as in the absence of ligand. Instead, it rises rapidly to a maximum and decreases gradually to a constant value. Figure 9 illustrates the factors giving rise to this effect. The Fe(I1) steady state concentration, i.e., the concentration when the rate of the species entering the cell is equal to the rate of reaction, is reached rapidly. In contrast, the ligand steady state concentration, reached when the rate of ligand entering the cell balances that leaving the cell, is attained more slowly. The Fe(I1)-ligand equilibrium is fast enough that the observed light from the cell is determined by the instantaneous concentrations of Fe(I1) and ligand (Figure 9C). This effect can indicate the presence of organic complexing agents in an unknown sample. If equal ligand concentrations are present both in the sample and in the background water, the ligand concentration in the cell remains constant while a slug of sample passes through and the light peaks do not decay. The intensity of light catalyzed by Fe(I1) is still linearly proportional to concentration in the presence of complexing ligands. The “practical” detection limit of lO-9M Fe(I1) is not affected by moderate ligand concentrations because the luminol-oxygen system is intrinsically sensitive to Fe(I1) concentrations two orders of magnitude more dilute than 10-9M. By reducing the sensitivity of response to Fe(II), organic ligands increase the upper limit of linear response to Fe(II), provided the concentration of dissolved oxygen in the cell is constant. Addition of organic ligands can be used to adapt the luminescence method to higher Fe(I1) concentrations.
Figure 10. Reduction of Fe(I1) peak height in the presence of varying concentrations of interfering metal ions o-D-O
A-A-A &-B--3
0-0-0 A-A-A
_-_
Cr(II1) Co(I1) Ni(I1) Mn cu Peak height in the absence of interfering metal
ions O , , 80 cm3; pH 10.5-11.0; 4 X 10-8M Fe(I1); 4 X 10-JMluminol
Also, it may be possible to “preserve” field samples against air oxidation by addition of an appropriate ligand. Inorganic Interferences. MnO.,-, OC1-, Is,and Fe(CN)6-3 have been reported to catalyze luminol chemiluminescence in the presence of dissolved oxygen (13-15’). We have found that V(I1) also catalyzes luminol chemiluminescence in the (13) H. H. Seliger in “A Symposium on Light Life,” W. D. McElroy and B. Glass, Ed.. Johns Hopkins Press, Baltimore, Md., 1961, p 204. (14) A. K. Babko and N. M. Lukovskaya. Dopor. Akad. Ncurk. U ~ VRSR. . 1962,619-621. (15) A. K. Babko, L. V. Markova and N. M. Lukovskaya. Zh. A / ~ o /Khin/., . 23, 401-6 (1968).
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Table I.
Comparison of Three Techniques for Fe Determination Atomic Sample CL analysis absorption Colorimetry Orchard leaves 10 240,220 pg/g ... NBS SRM 1571 Orchard leaves 24 237, 251 pg/g ... NBS SRM 1571 Pond water 0.72, 0.83 pg/ml 0.74 pg/ml ... River water 0.39, 0.38 pg/ml 0.34 pg/ml ., , Tap water 0.18, 0.15 pg/ml
