Determination of the Complexing Capacity of Natural Water by Cobalt(ll1) Complexation Kenneth W. Hanck* and James W. Dillard Department of Chemistry, North Carolina State University, Raleigh, N.C. 27607
The complexlngcapacity of natural water has been determined by complexation with Co3+. The method eliminates many of the problems of metal complex lability associated with previous procedures. A known excess of Co2+ was added to the sample. The Co(l1) complexes formed were oxldized to klnetically Inert Co(1ll) complexes with H202. After removlng the excess H202with the enzyme catalase, the unreacted Co2+ was determlned by differential pulse polarography in 0.06 M ethylenediamine. The detection limit is 0.4 X mol of ligand/l. Magnesium, calcium, sodium, potassium, chloride, sulfate, carbonate, and bicarbonate do not interfere. A semiautomated procedure was developed and applied to a series of natural water samples.
The role of organic matter which forms soluble complexes with trace metal ions in aquatic environments has been the subject of speculation and experiment for several years (1-9). One of the parameters frequently measured by workers studying metal ion/organic matter interactions is complexing capacity which is usually defined as the moles of metal ion which are complexed per liter of sample. Since complexing capacity is dependent on pH, time of reaction, the type of ligand in the sample, and the metal ion probe used, the values obtained by various workers show considerable variation. Most of the methods used for determining complexing capacity involve measurement of the amount of metal complex formed upon addition of excess metal or measurement of the remaining free metal directly. The specifics of several methods have been reviewed recently (10, 11). A popular method is essentially a micromolar titration with a metal ion titrant using anodic stripping voltammetry (ASV) to monitor the uncomplexed metal ion titrant and thus locate the end point of the titration (1,2, 7, 10). We have recently examined ASV and differential pulse polarography (DPP) as methods for locating the equivalence point of micromolar titrations of complexing agents. The results indicated t h a t the use of a labile metal ion titrant to probe the concentration of ligand was not feasible because of metal ligand dissociation a t the electrode surface prior to electron transfer. The subject of this investigation is the development of a procedure to utilize the use of inert metal complexes in determining the complexing capacity of natural water. Whether a metal complex is labile or inert is primarily a function of the electronic structure of the metal ion in the complex. Metal complexes of such metals as Cu(II), IntIII), and Co(I1) are considered labile while those of Cr(III), Fe(III), and Co(II1) are inert. Co(I1) reacts rapidly with ligands, forms stoichiometric complexes, and has a large formation constant (Kf) for most ligands of interest. In the presence of mild oxidizing agents such as dilute HzO2, Co(I1) complexes are converted to their respective Co(II1) complexes which are considered kinetically inert. For these reasons, Co(I1) was proposed as a metal ion probe for determining the complexing capacity of natural water. The cobalt procedure involves reaction of a known excess of Co(I1) with the unknown ligands found in natural water 404
ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
samples followed by chemical conversion of the Co(I1) complexes to their respective Co(II1) complexes. Only the Co(I1) which is complexed will be oxidized, for in aqueous solution aquo cobalt(II1) rapidly oxidizes water and returns t o t h e stable Co(I1) form (12). After conversion, either the amount of unreacted Co(I1) or the amount of complexed Co(II1) must be determined. This requires an analytical scheme for determining one oxidation state of cobalt in the presence of the other. Aquo cobalt(I1) would be retained by a complexing ion exchange medium ( 1 3 , 1 4 ) ;the Co(1II) complexes would not interact. Conventional analytical methods such as atomic absorption could then be used to determine the Co(II1). The ligand levels expected for natural water would produce ca. 0.5 kmol Co(1II)h. The minimum concentration of Co which can be reliably measured by flame atomic absorption is ca. 2.5 kmol/l. Carbon furnace atomic absorption techniques reduce this value to ca. 0.05 pmol/l. Since flame atomic absorption does not have an adequate determination limit and a “carbon furnace” instrument was unavailable, we have chosen not t o separate Co(I1) and Co(II1) but develop an electrochemical procedure for the determination of Co(I1) in the presence of Co(111). Although Co(I1) can be determined using ASV ( 1 5 ) , the procedure requires considerable sample treatment. Co(11) can be determined polarographically by observing the oxidation of Co(I1) in ethylenediamine (16).We have studied this procedure a t the micromolar level and verified its suitability for our purposes.
THEORY The reaction sequence is believed to involve two steps: co2+
+
+ rnL + COL,2+
2C0Lm2+ Hz02
(1)
+
+ 2H+
2CoLm3+ 2H20
(2)
The net reaction is the sum of the two steps: 2C02+
+ 2mL + HzOz + 2H+ e 2CoLm3++ 2Hz0
(3)
The overall equilibrium constant for this reaction is: Keq=
[ CoL,3+]2 [CO~+]~[L]~~[H~O~][H+]~
In order to gain an insight into the magnitude of the formation constants required for a quantitative determination, K,, must be estimated. The standard free energy of reaction 3 may be calculated from the standard free energies of reactions 1and 2. These in turn depend upon the equilibrium potential of H202/H20 and Co3+/Co2+ and the formation constants of CoLm2+and C0Lm3+(Kfz and Kf3 respectively).
AGO = -2RT In (Kf2)
-
+
E O C ~ ~ + / C ~ Z +E O H ~ O ~
Substituting 1.77 V for E O H ~1.84 ~ ~V, for E o ~ 0 3 + ~ and ~oz+ noting that A G O = - R T In Keq,we obtain a t 25 OC: log (Keq)= -2.37
+ 2 log (Kf3)
PAR J36
Table I. Variation of Equilibrium Constant with m
.V
Minimum values of K f 3 On
m
K,, (5% error)
5% error
95% error
1
4.6 x 1031 4.1 x 1 0 4 5 6.2 X 1060 1.0x 10103
1.0x 1017 1.0 x 1024 3.8 x 1031
2.6 x 1014 1.4 X lozo 9.5 x 1025 5.5 x 1043
2 3 6
4.7 x
1052
Off
Chort
x-Y on
The magnitude of the overall equilibrium constant and hence the quantitativeness of the determination depends on KO,the formation constant of c0L,3+. If t h e H202 reaction scheme converts 95% of the ligand in t h e sample to CoLm3+,a 5% error will occur. Such an error does not seem unreasonable a t the molar level. One may compute K,, by evaluating all of the concentration terms. Assuming a 95% conversion of ligand t o CoLm3+,an initial ligand concentration of M, an initial Co2+ concentration of M, an initial H202 concentration of M and a buffer of p H 7, minimum K,, values for m = 1,2,3,and 6 were evaluated (Table I). T h e minimum values of Kf3 needed to achieve K,, and hence a minimum of 5% error are also tabulated. Ligands forming complexes with Co(II1) weaker than indicated in Table I will be only partially determined. T h e weaker the ligand, the smaller its percent conversion of C0Lm3+and the larger the error in the analysis. The minimum value of Kf3 that a ligand must have t o be detected may be estimated by computing a K,, based on a minimum conversion of L to CoLm3+;a value of 5% conversion was used to compute the data in Table I. Thus a 1:l complex will be deM level if Kf3 2 2.6 X 1014and satisfactorily tected at the determined if Kf3 >- 1.0 X 1017. Unfortunately relatively few Co(II1) formation constants have been tabulated, making it difficult to predict which ligands can be determined.
EXPERIMENTAL Reagents. The water used in this study was deionized twice, passed through an activated charcoal column, and filtered through a 0.2-pm membrane filter. All standard solutions of synthetic complexing agents were prepared from dried reagent grade chemicals. Ethylenedinitrilotetraacetic acid (EDTA) (Fisher Scientific), 2-hydroxy-1,3-propylenedinitrilotetraacetic acid (Aldrich Chemical), and propylenedinitrilotetraccetic acid (PDTA) (Aldrich Chemical) were standardized by spectrophotometric titration in hexamethylenetetramine buffer (pH 5-6) at 5&0nm using standard cobalt solution as titrant and xylenol orange as indicator. Nitrilotriacetic acid (NTA) (Eastman Kodak) was standardized in ammonia buffer at pH 9.2 using murexide as indicator. Solutions of histidine (Fisher Scientific), tryptophan (Fisher Scientific), glutamic acid (Fisher Scientific); and aspartic acid (Fisher Scientific) were prepared by dissolving the dry sample in dilute base (NaOH) and then buffering at pH 7. Solutions of catalase (300 units/ml) were prepared from the dehydrated powder (Calbiochem, bovine liver 3100 units/mg) and autoclaved deionized water buffered at pH 7. One-milliliter portions were stored separately in autoclaved 4-ml glass vials for convenient allocation. Stock solutions (3 M) of ethylenediamine (Fisher Scientific) were prepared from the distilled 70% azeotrope. Hydrogen peroxide solutions (0.5 M) were prepared from 30% HzOz (Fisher Scientific). The strength was periodically checked by titration with standard Ce(1V) using ferroin indicator. Cobalt solutions were prepared by dissolving the appropriate amount of high purity metal (Alpha Inorganic) in Ultrex nitric acid (J. T. Baker Chemical Co.). The phosphate buffer (pH 7, 0.1 M) was prepared from (Fisher Scientific) reagent grade NaZHP04 and NaH2P04.H20. The supporting electrolyte (KN03) was recrystallized five times from water to eliminate a lead impurity prior to preparing a 1 M stock solution. Synthetic “World Average River Water” was prepared from Fisher Scientific reagent chemicals according to the composition of Reuter
7
FETI
1 reyt
1
I re!st
I
I
sv-a,r 0
FETZ
< Trigger in Figure 1. Schematic diagram of semi-automatic potential control ac-
cessory
and Perdue ( 1 7 ) and contained the following: magnesium (4 ppm), sodium (6 ppm), calcium (15 ppm), potassium (2 ppm), bicarbonate (58 ppm), sulfate (11 ppm), and chloride (8 pprn). Natural water samples were collected in 300-ml Pyrex bottles with polyethylene lined caps. Since no preservatives were added, all samples were stored at 5 “C and analyzed within 48 h. All stock and standard solutions were stored in polyethylene containers at 5 “C and at concentrations not lower than 10W M. Nitrogen gas (Air Products) used in purging sample solutions was deoxygenated by passage through a Ridox scrubber prior to use. Electrodes and Cells. A Corning fiber tipped saturated calomel reference electrode, a platinum auxiliary electrode (3 centimeters of 26 gauge Pt wire) and a standard dropping mercury electrode (DME) (Sargent-Welch S-29417 capillary, m = 3.896 mg Hgh) equipped with a PAR 172 drop timer (Princeton Applied Research Corp.) were used in the three-electrode cell. The cells used were 100-ml Berzelius Pyrex beakers fitted into commercially available polyethylene cell tops (Leeds and Northrup cell cover No. 067513 and ring No. 127182). Electronic Equipment. The electrochemical measurements were taken using the PAR Model 174 Polarographic Analyzer (Princeton Applied Research Corp.) and a Hewlett-Packard Moseley 7001A X-Y recorder or a Sargent-Welch Model SRG chart recorder. The electrometer and current output of the PAR 174 were multiplexed and displayed on a DTC Model 3312 digital panel meter (Data Technology Corp.). The PAR 174 was modified to allow semi-automatic differential pulse polarographic measurements. The circuit as shown in Figure 1 allowed continuous cycling between baseline and peak potential every 100 drops or automatic termination after 3 cycles of 100 drops each. The Heath EU-801A Analog-Digital Designer was used as a source of power and as a support for printed circuit cards. The following Heath cards were used: EU-&OO-LA,EU-&OO-CC,EU-&OO-JC, (2)EU-&OO-DEand EU-900-JA. The field effect transistor (FET) switches on card EU-900-JA are opened by a logic 0 and closed by a logic 1 except for FET 7 which is opened by a 0 and closed by a 1.The FET symbols used in Figure 1 are those used by the Heath Company and represent both the FET switch and its driver circuit. The monostable output pulse is a logic l of l-s duration. The relay circuit was constructed from an AMF (Potter and Bumfield, RlO-El-Y2, 6VDC) relay energized by a 2N-3393 transistor. Resistor R1 has the value 4.7 kfl. A Health SU-50-JA permanent patch accessory was used to hard wire the connections between cards. When making semi-automatic differential pulse measurements, an initial scan was made to estimate the peak potential and the baseline potential. The baseline potential was adjusted to the desired value with initial potential control of the PAR 174. The potential control accessory was then activated by placing switch S2 in the “On” position which in turn caused FET 4 to close. The peak potential could then be adjusted using the variable resistor R (100 kR) with the start switch S4 on and the trigger input disengaged. After baseline and peak potentials were set, the reset switch S5 was momentarily closed causing the Q output of t_heflip-flops (FF1, FF2, and FF3) to be initialized at logic 1. Since $1 was at logic 0, the output of AND gate 4 is low causing FET 3 to be open and the instrument to be applying the baseline potential. Initially, since NAND gate 3 was low and FETl was open, the timing pulses which occur approximately 75 ms before ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
405
0.0
Table 11. Co2+Calibration Data
t
x-intercept Run
Slope
(PM)
1 2
0.0033 f 0.0004 0.0037 f 0.0002 0.0037 f 0.0004 0.0036 f 0.0004 0.0038 f 0.0001 0.0036 f 0.0001 0.0037 f 0.0002 0.0036 f 0.0001
-1.6 f 7.1 0.5 f 2.0 2.1 f 5.0 -0.2 f 5.0 0.2 f 0.6 0.4 f 0.7 1.2 f 1.4 0.3 f 1.9
3
4 5 6 7 Composite
t 0 2 p A
.1
Ranee of Co2+ concentration: 2.9-58.0 oM. I
Table 111. Determination of Known EDTA Samples Run 1 2
3 4 Average 1 2 3
4 Average
A. 5.140 pmol EDTA per liter Slope pmol/l. found 0.0032 f 0.0001 0.0033 f 0.0001 0.0034 f 0.0001 0.0034 f 0.0001
B. 1.798 pmol EDTA per liter 0.000482 f 0.000007 0.000486 f 0.000006 0.000434 f 0.000013 0.000440 f 0.000013
6.3 f 1.9 4.7 f 2.2 4.5 f 1.2 5.1 f 2.1 5.1 f 1.3 1.4 f 0.5 1.4 f 0.4 2.8 f 1.1 1.3 f 1.1 1.7 f 1.1
mercury drop termination were not transmitted from the PAR 174 scope trigger output to the decade counting unit (DCU). Upon momentary closure of start switch S4, FF1, FF2, and FF3 are cleared and FF4 set; $4 and $1 both change to logic 1causing AND gate 4 to be high, closizg FET 3, and applying the peak potential; the Q output of FF2 and Q of FF4 falls to 0 closing FET 1and allowing timing pulses to pass. Simultaneouslythe sweep relay is activated and the X-Y or chart recorder started. For the X-Y recorder, only a momentary closure of the normally open sweep trigger relay was needed. Therefore, switch S3 was placed in the “X-Y” position causing FET 7 to close and FET 6 to be open. A 5-V pulse from start switch S4 would then energize the relay circuit when the cycling sequence was initiated. For the chart recorder, the sweep trigger relay had to remain closed for the entire cycling sequence. Switch S3 was placed in the “chart” position causing FET 7 to open and FET 6 to close. When cycling was initiated, FET 5 would be closed by a logic 1 from FF4 allowing 5 V to pass through FET 6 and energizing the relay circuit. The relay circuit would remain on until FET 5 was opened by a logic 0 from FF4 at the end of the cycling sequence.When 100 timing pulses were counted by DCU 1 and DCU 2, the output of NAND gate 1 changed from high to low triggering the monostablewhich in turn supplied a logic 1 pulse to FF1. Q1 is now high and $1 low opening FET 3 and applying the baseline potential to the cell. The high to low conversion of NAND gate 1is inverted by NAND gate 2 and used to reset the DCU’s and initiate a new cycle. This recycling continues indefinitely if switch S6 sets the J input of FF3 to logic 0 preventing final conversion.Otherwise, the NAND operation of $3 and Q2 after three complete cycles triggers conversion of FF4 which causes NAND gate 3 to go low opening FET 1 and preventing further transfer of timing pulses. To initiate a second sequence, the start switch S4 is depressed momentarily. Methods. The cobalt procedure involved adding an aliquot of Co(I1) to a known volume of sample (typically 50 ml) which was buffered at pH 7 . The size of the aliquot was chosen so that the Co(I1) concentration exceeded the expected ligand concentration. Twenty-five micromoles of hydrogen peroxide was added to oxidize the Co(I1)complexes formed to the corresponding Co(II1) complexes. Reaction time depends on the ligands and trace metals present in the sample, but should be at least 15 min. The excess HzOz must be removed from the solution before electrochemical analysis.The addition of 60 units of the enzyme catalase to the sample converts the H202 406
ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
I
I
04
0.5
I
0.6
I
0.7
- E , volts
Figure 2. Differential pulse polarogram of 5.8 X M Co(ll) in 0.06 M ethylenediamine. Drop time = 2 s; modulation amplitude = 100 mV
to oxygen and water. This step was carried out at 36 “C for at least 12 min. The dissolved oxygen was removed from the solution by deaeration with nitrogen. One millimole of KN03 was added and a differential pulse polarogram of the solution was obtained to verify that both the oxygen and HzOz had been sufficiently removed. The test solution was then made 0.06 M in ethylenediamine and the free Co(I1) measured by observing the current resulting from the differential pulse polarographic oxidation of C0(en)3~+ to C0(en)3~+. An aliquot of standard Co(I1)was then added to the solution and a differential pulse polarogram obtained. At least four such additions of metal titrant were made per sample. The complexing capacity of the sample was deduced from the x-intercept of a plot of the DPP peak current vs. the pmoles of Co(I1) added. A standard method procedure for the cobalt method is available (18).
RESULTS A N D DISCUSSION For theoretical as well as experimental reasons, pH control of the samples is required. Since the pH range of natural water samples is typically between 6.5 and 7.5, buffering the samples a t pH 7 would disturb the natural equilibria t h e least. Therefore, all samples in this study were buffered at p H 7 with a 0.1 M phosphate buffer. Removal of Excess H202. The HzOz which is required in t h e first step of the procedure must be removed prior to the electrochemical analysis. T h e presence of H202 upon addition of ethylenediamine would result in conversion of C0(en)3~+ t o C0(en)3~+,inducing error into the analysis because C0(en)3~+may not be distinguishable from the Co(II1) complexes with sample ligands. T h e electrochemical reduction of O2 and H202 also results in large undesirable background currents. Hydrogen peroxide can be removed from solution by vigorous boiling but this harsh treatment would obviously disrupt the complexation equilibria of a natural water sample. T h e milder enzymatic conversion of H202 t o 0 2 and water is preferable. Since the optimum conditions for the enzyme catalase are at pH 7 and 36 “C, incubation of the samples was required. T h e addition of 60 units of catalase per sample was found to remove 25 pmol of H202 in 15 min. The effectiveness of removing H202 was checked for each sample prior t o addition of the ethylenediamine by running a differential pulse polarogram of the solution. Since the supporting electrolyte KNO:j inhibits catalase activity, it was added after the incubation step but prior to deaerating the solution for electrochemical analysis. Quantitation of Co(en)S2+. The DPP oxidation of Co(en)32+t o C0(en)3~+was found to be a satisfactory method for t h e determination of Co(I1) in the presence of Co(II1). A
Table IV. Determination of Synthetic Ligand Samples Ligand
Fmol added
pmol found
EDTA PDTA 2-Hydroxy-PDTA NTA Histidine Aspartic acid Glutamic acid Tryptophan
0.257 0.322 0.290 1.04 1.12
0.258 f 0.064 0.320 f 0.019 0.289 f 0.076
1.10 1.22 1.07
ND ND ND ND ND
typical differential pulse polarogram of Co(I1) in 0.06 M ethylenediamine is shown in Figure 2. T h e calibration or “blank” experiment was conducted on samples containing no complexing agents using the procedure outlined in the Experimental section. This not only served as a check of the linearity of the calibration curve over the Co(I1) concentration region of interest, but also served as another check on the efficiency of removal of H202 by the enzyme catalase. The least squares analysis of seven Co(I1) calibration curves is summarized in Table 11.A statistically significant x-intercept was not observed in any case; the slopes of all runs are identical t o within the uncertainty of measurement. Determination of Synthetic Ligand Samples. Synthetic samples of EDTA over the concentration range of ligands anticipated in natural water were determined using the cobalt procedure (Table 111).The percent relative error of the mean for the first concentration is 0.19% and for the second concentration is -3.8%. T h e determination of several other ligands was attempted; the results are listed in Table IV along with the Co(I1) and Co(II1) conditional formation constants of the ligands. T h e “ND” indicates t h a t the uncertainity in the x-intercept was greater than the value of the intercept and is therefore “not detected”. The failure of the method to determine simple amino acids was disappointing. The reasons for this are unclear; Co(I1) is a borderline acid on the hard/soft scale (22) and consequently coordinates strongly with 0, N , and S donors (23). These simple ligands apparently do not form sufficiently strong Co(II1) complexes to permit determination; this hypothesis cannot be tested, however, since very few Co(II1) formation constants have been tabulated ( 2 4 , 2 5 ) . Interferences. Depending on their action, interferents fall into one of three categories: those which inhibit catalase, those which precipitate Co(II), and those which are electruactive in the same potential region as Co(en)32+. Chloroform, carbon monoxide, hydrogen cyanide, hydroxylamine, and ascorbate are known to inhibit catalase activity (19). If the excess H202 is not completely decomposed, Hz02 will oxidize Co(en)32+to Co(en)s3+.This will not cause a significant error since D P P cannot distinguish between the reduction of C0(en)3~+and the oxidation of Co(en):I2+ unless t h e electrode reaction is totally irreversible. Our experiments indicate t h a t i,/Cco for Co(en)33+is 8% smaller than t h a t of Co(en)32+.Thus the additional residual current caused by the reduction of undecomposed HzO2 is more significant than the partial oxidation of Co(en)32+. Of the substances likely t o precipitate Co(II), sulfide is the M S2- does most important. Under the conditions used not decrease the D P P peak current of a M solution of cobalt. Apparently the nucleation of COSis slow or significant amounts of COS dissolve after ethylenediamine is added forming Co(en)32+.Precipitative interferents are troublesome only in that it is difficult to distinguish bound forms of cobalt which are soluble from those t h a t are insoluble. If one wishes to measure only binding capacity, then both precipitating agents and soluble ligands should be determined. The scheme
1.09 x 2.06 x 1.40 X 4.78 X 1.57 X 3.97 x 1.85 x 1.77 X
1.50 x 1038 2.06 x 1038 8.47 x 1037 1.02 x 107
1013 1013
10” 10” 1O1O 107
... ... ... ...
105
lo2
Table V. Determination of Bound PDTA Metal added, PDTA added, pmol Fmol 0.515 Cazf 0.537 Cd2+ 0.536 Hg2+
0.322 0.322 0.322
PDTA found pmol
K’f
0.321 f 0.049 0.351 f 0.117 0.367 f 0.098
4.40 X lo7 6.97,X 6.66 X l O l 8
of Batley and Florence (9) is the only one of which we are aware which systematically determines soluble and insoluble binding agents separately. Although Cu(II), Fe(III), and Pb(I1) exhibit electrochemical behavior in ethylenediamine in the potential region of Co(II), blank currents of natural water samples were negligible. If an interfering peak appears, the peak current is subtracted from each of the peaks obtained when the sample plus Co(I1) and H202 is run. Magnesium, calcium, sodium, potassium chloride, sulfate, carbonate, and bicarbonate a t levels normally found in ground water do not interfere. This was evaluated by determining PDTA in a “World Average River Water” medium (17).A very slight precipitate was observed upon addition of t h e ethylenediamine to t h e sample, presumably an insoluble phosphate salt of calcium and magnesium. Of the 0.322 hmol of PDTA added, 0.335 f 0.077 Mmol were found. Metal Exchange Reactions. Theoretically, since Co(II1) complexes generally have significantly higher formation constants than their respective Co(I1) complexes, free ligands as well as ligands complexed t o trace metals in natural water should be determinable. Depending on ligand concentration, trace metal ion concentration, Co(I1) concentration, and the lability of the complexes formed, the CoL, 2+ conversion t o CoL, 3+ should drive the equilibrium toward eventual conversion of all the ligands t o the Co(II1) complex provided the minimum Kf3 value is met. The rate of the reaction, however, cannot be predicted by this solely thermodynamic approach. The results of selected bound determinations of PDTA are listed in Table V. Metals which have smaller K( values than the Kf2 (2.06 X for PDTA) for cobalt are easily replaced and the ligand determined; this was observed with Ca. Metal complexes with formation constants larger than Kf2 but smaller than Kf3 should be determined; the PDTA bound by Cd and Hg was determined by the Co(II1) technique. T h e mechanism of ligand exchange reactions depends heavily on the ligands and metal involved. Further work in this area is needed before the types of bound ligands which will be displaced by the Co(II1) method can be firmly established. Detection Limit. The detection limit of the cobalt method under the conditions we employed may be estimated from the variance of the x-intercept of the “standard addition” plot by application of the t test. If four additions in triplicate are made, the null hypothesis that the x-intercept equals zero (Le., no ligands detected) may be rejected in favor of the alternate hypothesis t h a t the x-intercept is greater than zero (i.e., liANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
407
Table VI. Effect of Ligand Concentration on Minimum
m
10-5 M
Keq-C~2m+1 4.6 x 1013 4.1 x 1015 6.2 X l0l8 1.0 x 1 0 2 5
1 2 3 6
Kf3
M
3.3 x 1015 3.1 X loz1 1.2 x 1028 1.5 x 1046
1.0 x 1017 1.0x 1024
3.8 x 1031 4.7 x 1052
10-7 M
10-8 M
1.0 x 10'8 3.1 X loz6 3.8 x 1037 1.5 x 1059
1.0 x 1020 9.8 X loz8 3.8 x 1038 4.8 X lofis
Table VII. Complexing Capacity of Natural Watera 1
2
3
4
5
6
7
8
1.2 f 0.7 1.7 f 0.8 ND ND 0.7
2.1 f 0 . 5 1.1 f 0.3 1.6 f 0.8 1.4 f 0 . 7 1.6
1.5 f 0.7 1.0 f 0.3 1.8 f 0.4 1.7 f 0.4 1.5
2.1 f 0.5 ND 1.9 f 0.2 1.3 f 0 . 4 1.3
2.4 f 0.6 ND ND 0.8 f 0.4 0.8
1.7 f 0.7 ND ND 1.2 f 0.6 0.7
2.0 f 0.2 1.4 f 0.3 0.5 f 0.4
1.5 f 0.3 0.6 f 0.5 0.7 f 0.4 0.9 f 0.6 0.9
A. Sampling network Week I I1 I11 IV Averages
B. Miscellaneous samples Raleigh City tap water Raleigh sewage effluent Briery Swamp (Stokes, N.C.) Tar River (Greenville, N.C.) Aquarium water a
1.3 f 0.3 1.3
ND 4.1 f 0.9 0.6 f 0.1 ND 7 . 5 f 0.1
ymol ligand/l. ND = Not detectable.
I
I
I
I
0.0
-
+
c
2
C
2 3 u
3
U
04
0.6
05
0.7
-E, volts
Flgure 3. Differential pulse polarogram of sample 11-3after complexation of Co(lll)
gands detected) provided that the computed statistic t exceeds 1.812 (95% confidence level, one tail test, 10 degrees of freedom). Since the variance of the x-intercept is about 1.0 for solutions of EDTA below 2 Kmolar, the minimum concentration of ligands which can be detected with 95% confidence is 0.4 X mol/l. T h e value of Kf3 needed for determination may be decreased by increasing the initial concentration of Co(I1) or H202. However, since the sensitivity of t h e method depends on observing a loss in the concentration of Co(II), a large initial Co(I1) level will result in a small percentage decrease in Co(I1) upon formation of CoL, 3+. If the decrease is of the same order of magnitude as the uncertainty in measuring the Co(I1) concentration, no loss will be detected and hence no complexing capacity found. In this work, a tenfold excess of Co(I1) over ligand was found t o be the most desirable. 408
ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
Time Figure 4. Semi-automatic readout of differential pulse polarogram of
samDle 11-3
Two problems result if the Hz02 level is increased; complicating side reactions with oxidizable organic matter may appear and the time required for H202 removal by catalase will increase. An addition of 60 units of catalase was found to satisfactorily decompose 50 ml of l W 5 M Hz02 in 15 min at 36 OC.Increasing the amount of H202 used by a factor of 10 will cause K,, to decrease by a factor of ten but the time needed to decompose the HzO2 will increase proportionally. The decomposition time could be reduced by adding more units of catalase but a t the risk of introducing errors caused by Co(I1) binding to the protein structure of catalase; 60 units of catalase represent approximately 10-10 mol of protein (20).
The sensitivity of the method toward ligands forming weak complexes with Co(II1) may be improved by changing the pH. A decrease in p H will cause a decrease in the K,, needed for
quantitative conversion of L to CoL, 3+ and a corresponding decrease in the minimum value of Kf3 a ligand must have to be determined. T h e effective formatioh constant for ligands containing carboxyl or other acid groups is also a function of pH; as p H is decreased, the effective formation constant decreases. Thus the effect of decreasing the p H cannot be readily predicted unless the identity of the sample ligands is known. T h e practical p H range is 5 to 10 because t h e efficiency of catalase in decomposing HzO2 decreases rapidly below p H 5; the catalase molecule decomposes above p H 10. T h e effect of ligand concentration on the feasibility of determination is similar to t h a t of increasing the initial conwill + 1be centration of Co2+ or Hz02. T h e product K e q - C ~ 2 m constant provided the initial concentration of Co2+ and H202 are made equal to lOCL (see Table VI). The results of Table VI imply that ligands having a high Kf3 can be determined at lower concentrations than ligands of low consequently a 3.6 Kf3. At pH 7 , Kf3 for EDTA is 4.8 X X M solution should be determinable! In order to perform such an analysis, an analytical method sensitive a t the M level is required; since DPP is not reliable a t conM, determination of ligands a t concentrations below centrations below mol/l. is impossible regardless of p H or initial concentration of Co2+ and H202. Evaluation of Natural Water Samples. Eight sites in the metropolitan area of Raleigh, N.C., were selected for collection of natural water samples. T h e sites were chosen to reflect industrial, commercial, and municipal usage. A more complete description of the sites is available elsewhere (18).The results of analysis of 50.0-ml aliquots of each sample are shown in Table VII. All the data in Table VI1 was obtained using the semi-automatic potential control accessory shown in Figure 1. A conventional differential pulse polarogram of sample 11-3 is shown in Figure 3, a n d the semi-automatic output of this sample is shown in Figure 4. Since the semi-automatic method is static a t zero scan rate, the peak current distortions for scanned polarograms using the PAR 174 should be eliminated (21). Also, the peak current a n d baseline current are much easier to obtain from Figure 4 than from Figure 3. Four known additions of Co(I1) were used in obtaining each data point of Table VII. In computing the concentration of ligands in a sample in which the chemical identity of the ligand is unknown, the value of m must be assumed. Since most of the ligands for which m would be greater than 1form Co(II1) complexes that are too weak to be detected, the value of m has been assumed to be 1. T h e results of the sampling sites can be divided into two groups: sites 1, 5, 6, 8 and sites 2, 3, 4, 7 . T h e average complexing capacity of each site in the first group was about 0.8 pmolll., while t h a t of the second group was about 1.4 pmol/l.
Sites 1 , 5 , 6 , 8 are located on streams which receive relatively little industrial or municipal use. Sample station 1 is the control and is located on a tributary of Crabtree Creek in Umstead State Park. Sites 2,3,4, and 7 are located on streams which receive moderate industrial and municipal use. Based on the limited number of samples examined, complexing capacity is higher in streams which receive waste effluents than in protected streams. T h e Co(II1) method should, however, be employed routinely over an extended period of time on numerous documented samples before drawing firm conclusions on the correlation between complexing capacity and water quality.
LITERATURE CITED (1) H. E. Allen, W. R . Matson, and K. H. Mancy, J. WaterPollut. ControlFed., 42, 573 (1970). (2) M. E. Bender, W. R . Matson and R. A. Jordan, Environ. Sci. Technol., 4, 520 (1970). (3) M. Schnitzer, "Metal-Organic Matter Interactions in Soils and Waters", in "Organic Compounds in Aquatic Environments", S. D. Faust and J. V. Hunter, Ed., Marcel Dekker, New York, N.Y., 1971. (4) A. Siegel, "Metal-Organic Interactions in the Marine Environment", in Ref. 3. (5) C. G. Gregor, Eflvlron. Sci. Technol., 6, 278 (1972). (6) Y. K. Chau, J. Chromatogr. Scl., 11, 579 (1973). (7) Y. K. Chau and K. Lum-Shue-Chan, Water Res., 8, 383 (1974). (8) R. D. Guy, C. L. Chakrabarti, and L. L. Schramm, Can. J. Chem., 53, 661 (1975). (9) G. E. Batley and T. M. Florence, Anal. Lett., 9, 379 (1976). 10) T. A. O'Shea and K. H. Mancy, Anal. Chem., 48, 1603 (1976). 11) K. W. Hanck and J. W. Dillard, Anal. Chim. Acta, in press (1976). 12) F. A. Cotton and G. Wilkinson, "Advanced Inorganic Chemistry", Interscience Publishers, New York, N.Y., 1972, p 875. 13) D. E. Leyden and G. H. Luttrell, Anal. Chem., 47, 1612 (1975). 14) "Bio-Rad Laboratories Technical Bulletin No. 114", Bio-Rad Laboratories, Richmond, Calif., 1964. 15) B. K. Housepian and I. Shain, J. Electroanal. Chem., 12, 397-410 (19661.
(16) "Handbook of Analytical Chemistry", Louis Meites, Ed., McGraw-Hill, New York. N Y.. 1963. D 5.67. (17) J H Reuter and E M Perdue, U S Nat Tech lnform Ser , PB Rep, 1972, No 210714 (18) James W. Dillard, Ph.D. Thesis, North Carolina State University, 1976. (19) H. U. Bergmeyer, "Methods of Enzymatic Analysis", translated by D. H. Williamson, Academic Press, New York, N.Y., 1963, pp 885-972. (20) T. Samejima, M. Karnata, and K. Shibata, J. Biochem. (Tokoyo), 51, 181-187 (1962). (21) J. H. Christie, J. Osteryoung, and R . A. Osteryoung, Anal. Chem., 45, 210 (1973). (22) R. G. Pearson, "Hard and Soft Acids and Bases", Dowden, Hutchinson and Ross, Inc., Stroudsburg, Pa., 1973. (23) H. Sigel and D. B. McCormick, Acc. Chem. Res., 3, 201 (1970). (24) A. E. Martell and R . M. Smith, "Critical Stability Constants", Vol. I, Plenum Press, New York, N.Y., 1974. (25) L. G. Sillen and A. E. Martell, "Stability Constants of Metal-Ion Complexes", Spec. Pub/. 17, The Chemical Society, London, 1964.
RECEIVEDfor review July 22,1976. Accepted November 18, 1976. A grant from the Water Resources Research Institute of the University of North Carolina (A-052-NC) partially supported this work.
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