942
Anal. Chem. 1980, 52, 942-945
nitric acid concentration to a maximum of 5.4 M. The bismuth ion assay was not dependent on the nitric acid content and the relative standard deviation of both determinations was *0.2%. In a separate determination, nine samples were assayed titrimetrically for bismuth ion and assayed spectrophotometrically following the concentration of the yellow BiLcomplex ( 5 ) . The average relative error for this cross-over study was 1.2%. This study has shown that a tandem titration is both possible and useful in the nitric acid-bismuth ion system. The concentrations of both species may be determined quite accurately in a system that is otherwise tedious, if not difficult to assay. Pyrocatecholsulfonephthalein has always been a reasonable choice as an indicator in a complexometric bismuth titration; the p H control necessary to ensure a sharp end point is built into this procedure. The appearance of four sharp end points in one titration reduces the equipment and preparation of a reasonably sophisticated assay to a minimum. Since aqueous bismuth solutions are generally prepared by the action of nitric acid on reference standard bismuth metal, a simple procedure for joint assay is a useful one. The method can be used over a rather large range of bismuth ion and nitric acid concentrations.
Miklos, I.; Szsgedi, R. Acta Chim. Acad. Sci. Hung. 1961, 2 6 , 356; Chem. Abstr. 1961, 5 5 , 2196a. Beck, M. T.; Gergely, A . Acta Chim. Acad. Sci. Hung. 1966, 50, 155; Chem. Abstr. 1965, 6 3 , 3682e. Dean, J. "Lange's Handbook of Chemistry", 11th ed.,Mc&aw-Hill: New York, 1973; pp 5-62. Schwarzenbach, G.; Fhschka, Hermann, "Complexometric Titrations"; Methuen: London, 1969; pp 309-13. Young, R. S. "Chemical Analysis in Extractive Metallurgy"; Charles Griffin: London, 1971. Kolthoff, I.M.; Elving, P. J. "Treatise on Analytical Chemistry"; Interscience: New York, 1963; pp 147-73. Belcher, R.; Close, R. A. Chemist-Analyst 1957, 4 6 , 86. Belcher, R.;Close, R. A. Chemist-Analyst 1958, 4 7 , 2. Cheng, K. L. Anal. Chem. 1954, 26, 1977. Yamauchi, Osamu; Tanaka, Hisashi; Uno, Toyozo. Tabnta 1968, 75(5), 459-74; Chem. Abstr. 1967, 6 6 , 2069q. Wada, Hiroko; Nakagawa. Genkichi. Nippon Kagaku Zasshi, 1968, 89(5), 499-503; Chem. Abstr. 1968, 6 9 , 326402. Shchurova, L. M. Proizvod. Pirazolona Anal. Reagenty, Metody fir.Khim. Anal. 1976, 145-7; Chem. Abstr., 1975, 82, 67814e. Garcia Montelongo, F.; Perex Olmos. R. An. Quim. 1978, 74(6), 937-944. Sonoda, Kiyokazu; Otomo, Makoto; Kodama, Kazunobu. Bunseki Kagaku 1976, 27(7), 429-34; Chem. Abstr., 1978, 8 9 , 2 0 8 5 0 5 ~ . Kesavan, S.;Garg, B. S.;Singh, R. P. J . Chin. Chem. SOC.(Taipei) 1977, 24(4), 181-6; Chem. Abstr., 1977, 8 7 , 33200q. Gattow, G.; Schott. D. Fresenius' 2. Anal. Chem. 1962, 788, 10. Wilson, Cecil L.; Wilson, David W. "Comprehensive Analytical chemistry", VoI. I;Elsevier: New York, 1960; Chapter 9. Ryba, 0.; Cifka, J.; Maht, M.; Suk. M. Chem. Listy, 1955, 49, 1786-91; Chem. Abstr., 1956, 5 0 , 3147e. Brookes. H. E.; Johnson, C. A. J . Pharm. Pharmacol. 1955, 7 , 836. Barcsa, L.; Koros, E. Chemist-Analyst 1959, 4 8 , 94.
LITERATURE CITED ( 1 ) Katstrup, E. K.; Boyd, J. R. "Facts and Comparison", Facts and Comparison, Inc.: St. Louis, Mo., 1978; p 326.
RECEIVEDfor review January 17, 1980. Accepted February 25, 1980.
Water Hardness Determination by the Catalytic Polarographic Wave of Magnesium Ion M. C. Cheney, D. J. Curran,' and K. S. Fletcher 111" Corporate Research, The Foxboro Company, Foxboro, Mass. 02035
A method for the determination of water hardness is based on the polarographic reduction of magnesium ion. Mg2+ is displaced from its EDTA complex by cations contributing to water hardness and the liberated Mg2+ is reduced at the dropping mercury electrode via a catalytic process. The enhanced current which results, yields approximately 100-fold ampliflcation over that obtained from a conventional diff uslon controlled process. This enhanced sensitivity and the ability to measure hardness contribution from several dissolved metals simultaneously Is demonstrated. Precision and accuracy comparable to that obtained using standard EDTA titrations are achieved by this more rapid technique.
T h e determination of hardness in water, both potable and industrial, is one of the most frequently required analytical procedures. The ASTM Standard defines hardness as the total concentration of calcium and magnesium and notes that hardness is also caused by other metals such as copper, iron, manganese, and zinc ( I ) . A gravimetric procedure for calcium P e r m a n e n t address: D e p a r t m e n t of Chemistry. I . n i v c r s i r y 01 Massachusetts, A m h e r s t , >lass. 01001). 0003-2700/80/0352-0942$01.00/0
and magnesium hardness based on precipitation of calcium oxalate and of magnesium ammonium pyrophosphate is relatively free from interference from other metal ions but is tedious. In the EDTA titration procedure, many metal ions block the indicators and need to be masked ( 2 , 3 )or removed ( 4 ) prior to the titration for calcium and magnesium. This is unfortunate since in many cases the analysis of interest is for total metal ion content. In this situation, the Hg-Hg(EDTA)2- couple sensed a t a mercury electrode ( 5 ) or the CU-CU(EDTA)~-couple sensed a t a copper ion-selective electrode (6) may be used for potentiometric titrations and are generally applicable to virtually all metal ions. Several ion selective electrodes for water hardness, based on immobilized liquid ion exchangers, have been described (7). They display Nernstian response to Ca2+and Mg2+over the concentration range lo-' to M but have annoyingly high response to other ions such as H', the alkali metals, and other heavy metal ions. Analyses with these electrodes therefore generally reflect total hardness rather than that due to Caz+ and Mg*+ only. Their application in direct potentiometry requires control of total ionic strength, adjustment of sample pH, and, depending on the total concentration of hardness, account must be made of the contribution of monovalent cations to electrode response. One manufacturer 1980 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 52, NO. 6,M A Y 1980
circumvents these difficulties by providing a nomographic correction of water hardness activity with a measurement of specific conductance (8). Their limited dynamic range and moderate sensitivity to monovalent cations restricts the use of these ion selective electrodes as monitors in potentiometric titrations. The hardness measurement technique being presented here is based on the stoichiometric displacement of magnesium from its relatively weak EDTA complex by most divalent metal ions. Subsequent determination of free Mg2+using the catalytic wave that results when it is reduced a t the dropping mercury electrode (9) provides a substantial amplification effect. Though the relationship between observed catalytic current and concentration of Mg2+is linear over a relatively narrow range of concentration, the technique offers a rapid and sensitive method for determination of water hardness.
EXPERIMENTAL Polarographic measurements were made with a voltage sweep rate of 2 mV/s from -1.9 t o --2.5 V vs. SCE using a homemade sweep generator and a Princeton Applied Research Model 173 Potentiostat with Model 176 Plug-in. A Metrohrn Cniversal Titration Vessel was used for the cell. A glass pH electrode, nitrogen inlet and outlet tubes, and the polarographic electrodes were introduced through the cap. The reference SCE (Corning, 476002) and platinum auxiliary electrode were isolated from the test solution by salt bridges having fine porosity giass frits and agar plugs made with 0.1 M tetraethylatnmoi-iium bromide (TEAB) electrolyte. The dropping mercury electrode \L)ME) vas constructed with a 6-12 s capillary (E. H. Sargent & Co., S-29417J and the height of the mercury co!umn Esed during the analyrical runs was 57.5 cm. The supporting electrolyte was 0.10 M TEAB (Baker, recrystallized once from i-propanol). All metal ion solutiuns were prepared to be about 5 mM from reagent grade salts !MgC12.6H20, Fe(C104)2-6H20,ZnS04.7H20, CuBr2, and P b ( N 0 3 J 2 )and standardized (using procedures described by Flaschka ( / O ) ) with 0.010 M disodium EDTA (Baker A.R.) which in turn was standardized against CaC03 (Baker A.R.) as suggested by Iiolthoff et al. (11). Another EDTA solution was prepared hy adding a 0.7 M solution of tetraethylammonium hydroxide (TEAH) (Raker Grade) to 1.8 g EDTA in the acid form (Baker A.R.) until the solid was dissolved, then adjusting the pH to 10.0 with TEAH and diluting to 500 mL. This solution was titrated against standardized (0.08 M) MgCl2 (pH 10,Eriochrome Black T indicator (EBT)) to determine t h e relative concentrations of the two reagents and a 0.017 M solution containing stoichiometric amounts of Mg2+and EDTA (with about 10 pM excess of Mgzc) was prepared.
RESULTS A N D DISCUSSION The C a t a l y t i c Mg2+ Wave. The polarographic currentpotent,ial curve for reduction of Mg2+ion in TEAB electrolyte shows a large maximum a t about -2.35 V vs. SCF; the observed current is much higher than would be expected for a diffusion controlled process and the reductior, of Mg2+ is accompanied by liberation of large quantities of hydrogen gas a t the electrode surface. We have evaluated the dependence of current a t the maximum on height of the mercury column and find it nearly exactly proportional to h1;2. These observations are not in conformance with generally stated diagnostic criteria for catal>%icprocesses, i.e., the current should be invariant with h'/' (12)and not show a maximum. Various authors have suggested that the enhanced current and the appearance of hydrogen gas result from reaction of an intermediate species such as Mg+ with water generating Mg'+; presumably the decrease in current after the maximum is due to the increased rate of direct reduction of Mg'+ to magnesium amalgam. Moreover, since the Mg2+ which is reduced to form the catalytic species is uitimate!y removed from the vicinity of the electrode by formation of the amalgam and must be resupplied by diffusion, dependence of wave height on drop
943
1
3c
3-
3
d
A
.
~~
--._ ?.
E^ a3 3" -9T'"\3' v1.-,41
~
~
?I
_
-
~
Figure 1. Catalytic current for Mg2+ reduction at 23.6 OC (A), 30 O C (- ), and 40 O C (0). Inset: (0)slope of linear portion of plots of current vs. [ ~ g * ' ] as a function of temperature
time is to be expected. A discussion, of mechanisms may be found in ref. 9. The Mg(EDT.4);'- complex is not reducible a t the DME and no maximum is observed for this species. As Mg2+is added to the electrolyte or liberated from the complex by addition of other cations forming stronger EDTA complexes, the wave gives only an indistinct maximum until a Concentration of about 20 p M Mg2+is reached (peak current about 15 FA) and then peak height increases linearly with concentration to about 100 pcM Mg" (peak current about 1 C 1 0 p A ) where, because of hydrogen evolution, polarograms become erratic. The background current rises slowly through the maximum, then rapidly a t potentials cathodic to the maximum, making evaluation of background contribution difficult. Though there is a minor shift of the maximum in the cathodic direction as concentration is increased, we used the total current as a measure of the Mg" concentration without mbtracting the hackground. This procedure is particularly convenient since our goal is to determine the free Mg2+ released in the electrolbte by taking the difference between two waves. The wave is relatively insensitive to p H over the range 9.0 to 11.0. In all work reported here, p H was adjust,ed to 10.0 k 0.2 with TEAH and HCl. Because alkali metals and ammonium ions are discharged in the potential range being studied, their use in making supporting electrolytes and the Mg(EDTA)2- solution was avoided. However, introduction of small quantities with reagents or as part of the sample posed no problem since the diffusion current produced is very small compared t o the enhanced current resiilting from the catalytic process. Calibration curves were obtained a t three temperatures, 23.6, 30.0, and 40.0 "C, to demonstrate the range of linear concentration dependence, to determine the sensitivity of the method, and to measure the temperature coefficient. Plots of current vs. concentration are shown in Figure I. for the three temperatures. Straight lines were fitted (linear regression analysis) to the point3 falling in the linear portion of each curve and the slopes of these lines, 0.84, 0.92, and 1.00 pA/pM, respectively, illustrate the extremely good sensitivity of this method over the range of Concentrations studied. These slopes were plotted against temperature (inset, Figure 1). Based on this, a t 30 "C, a 1 "C change in temperature results in approximately a 1% change in slope. The analytical work which follows was performed at ambient ternperature.5 (24 f 1 "C) without thermostating.
944
ANALYTICAL CHEMISTRY, VOL. 52, NO. 6, MAY 1980
L-
o
23
do
MI
ac
ADDED CONCE'ITRATlOhi
oc 31
i:
MQ"IN
Figure 2. Catalytic current observed for [Mg"] reduction resutting from addition of various metals to 1.0 mM solution of Mg(EDTA)'- in 0.1 M TEAB at 24 OC (a,Mg2+; 0 , Zn2+; 0 ,Ca2+,A , Pb2+)
Replacement of Mg2+f r o m Its EDTA Complex. The replacement technique is based on a reaction of the type
Mn+ + Mg(EDTA)2-
-
M(EDTA)(n-4)+Mg2+ (1)
where the metal ion M"+ displaces Mg2+ from its relatively weak EDTA complex. In principle, Equation 1 extends the amplification effect which results from the catalytic polarographic reduction of free Mg2+to any metal ion which will displace it from its EDTA complex. In practice we have found this limited t o metal ions whose EDTA complexes are not themselves directly reducible a t the DME over the potential range of interest, since reducible complexes diffusing to the DME will react, liberating free EDTA which decreases the amount of Mg2+ in the vicinity of the electrode. It should be noted that since only Mg2+gives an amplified wave under the conditions of this experiment, the technique can be used to determine hardness due only to Mg2+ independent of contributions from other metals; if no Mg(EDTA)'is added to the supporting electrolytes, other hardness-causing metals will give ordinary diffusion-limited waves, while Mg2+ will give an enhanced wave which will be larger by nearly two orders of magnitude for a given concentration than a diffusion-controlled wave. Other metals, such as alkali metals and NH4+, which are reducible a t these potentials but do not displace Mg2+from its EDTA complex, are also discriminated against by two orders of magnitude. Thus if alkali metals are present in the same concentration as hardness-causing metals, a n error of only 1%will result. Calibration data were obtained for incremental additions (50- to 100-pL additions of 0.005 pM solutions with a 100-pL calibrated syringe fitted with Chaney adaptor) of Mg", Ca2+, Zn2+,Pb2+,Fez+,and Cu2+to the supporting electrolyte which was 0.1 M TEAB, 0.001 M Mg(EDTA)2-, and approximately 10 pM free Mg2+. A fresh solution was used for each metal analysis and addition of samples was stopped when the total measured current approached 100 FA. The data shown in Figure 2 illustrate that the currents which result from additions of Mg2+, Ca2+,Zn2+,and Pb2+ fall on a straight line having the slope 0.837 pA/pM (24 "C) and a correlation coefficient of 0.991. Addition of Cu2+to the supporting electrolyte containing Mg(EDTA)2- results in about half the expected catalytic current, even though the copper complex with EDTA is much stronger than that of magnesium. We propose that this is due t o direct reduction of Cu(EDTA)'- at the electrode surface,
since the EDTA thereby released would combine with free Mg2+in the vicinity of the DME and depress the catalytic current. This interpretation is supported by our observation that a 1.0 mM solution of Cu(EDTA)'-, 0.1 M in TEAB and adjusted to pH 9.4, shows a normal polarographic wave with a half-wave potential a t 4 . 4 8 V vs. SCE. The magnitude of the error resulting from the presence of copper is a function of its concentration. In the volumetric titration procedure, copper, in even trace quantities, blocks the EBT indicator and is normally masked by addition of cyanide (2). The presence of small amounts of Cu2+ in the solutioiis analyzed by our present technique is not nearly as serious, e.g., if Cu2+contributed 1% to the hardness of a sample, the result from the present method would be 0.5% low, as the Cu2+would only produce in the vicinity of the DME, half the stoichiometrically expected concentration of Mg2+. The use of cyanide as a masking agent is still useful since the cuprous cyanide formed is totally inactive polarographically (13).To test this, we made a supporting electrolyte containing 1 mM Mg(EDTA)2-,0.10 M TEAB, and 0.02 M tetraethylammonium cyanide (by treatment of 0.02 M TEAH with HCN to a p H of 9.7) and added Cu2+up to a final concentration of 30 pM. No Mg2+ wave resulted. Additions of Mg2+and Ca2+,however, produced catalytic currents in good agreement with the calibration curve shown in Figure 2 . The behavior of iron is more complex. Making the electrolyte, a t pH 10, as little as 5 pM in Fez+ or Fe3+gives rise to a wave with a large current maximum of about 20 pA a t potentials 0.25 V anodic to the magnesium maximum. This wave was also accompanied by evolution of gas at the electrode and occurred regardless of the presence of Mg2+ or Br-, but disappeared in the presence of excess EDTA. As the concentration of iron is increased, the current a t the maximum increases very rapidly and becomes erratic, presumably because of gas evolution. We have concluded that some iron species is being reduced at these potentials and that formation of metallic iron, which does not amalgamate, on the mercury surface reduces the overpotential for reduction of water, leading to the large increase in the current. In spite of the great stability of the iron-EDTA complexes, they are known to form slowly and additions of Fez+ or Fe3+ to the basic Mg(EDTA)'- solutions may lead to formation of iron hydroxide complexes with loss of stoichiometry for release of Mg2+. While Fe(EDTA)2- is not reducible at the DME, its incomplete formation leaves a species in solution that is reduced to metallic iron. The large current that results from the apparent discharge of solvent interferes with the analytical wave when the concentration of iron in the supporting electrolyte is about 25 pM. We have observed that making the electrolyte acidic (pH 3) after addition of iron, then returning it to pH 10 prior to analysis, reduces the size of this interfering wave, and enlarges the analytical wave, again suggesting that incomplete formation of the iron-EDTA complex is involved. At lower pH, the competition of hydroxide for iron is less strong. We found on reducing the p H of the electrolyte to 8.1, that addition of iron would contribute quantitatively to release of Mg2+and that the level of iron could be increased to 60 pM without onset of the erratic polarographic behavior. While solutions a t this p H could be used to construct a calibration curve where iron is the species of interest, we chose not to work a t this pH because Mg2+is incompletely complexed by EDTA and thus an error would result for Mg2+ hardness. Cyanide has been suggested as a masking agent for Fe3+ (14). However, we found that making the test solution 1.0 mM in KCN was ineffective in removing the interfering wave. ,4chromatographic procedure using a silica gel activated with an amino silyl functional group has been described for removal
ANALYTICAL CHEMISTRY, VOL. 52, NO. 6, MAY 1980 T a b l e I. R e s u l t s of H a r d n e s s M e a s u r e m e n t s W a t e r S a m p l e s , ppm CaCO,
of
EDTA sample standard h a r d water a r t e s i a n well w a t e r s u r f a c e well water Foxboro t a p w a t e r seawater total
Ma only
titration
45.2 f 106.4 i 35.6 f 38.0 f
0.1 0.1
0.1
0.3
5670 = 10 4760 i 30
polarographic
45.6 i 107.5 i 35.6 i 37.9 *
1.3 0.5 0.3 0.7
5600 4740
i
110
i
35
of both iron and copper from samples prior to water hardness titration ( 4 ) and could be used where the contribution of these ions to the hardness determination is not required. T h e effect of the common anions SO4'-, Po43-, C032-, and F- on the catalytic Mg2+ wave was tested using supporting electrolytes 30 pM in Mg2+to which were added solutions of these anions to make their concentration in the electrolyte 300 pM. Their effect on the catalytic wave was minimal, causing only 1 t o 2% depression of the wave in each case. Since Mg2+ forms stable complexes with oxygen containing anionic ligands such as P207'-and chelating ligands such as EDTA, their presence in the samples could be an interference in this as well as in the volumetric titration procedure. The level of interference depends on the concentration of the ligand and the formation constant for the complex. We expect the determination of hardness to be quantitative by our procedure when the final concentration of the complex MgX is no more than 0.1% of the total hardness added. With this condition t h e formation constant expression becomes where Kf is the formation constant and [XI is the concentration of the complexing ligand. Thus, for the simple case of a 1:l complex, the product of the formation constant and the molar concentration of the complexing ligand must be equal to or less than for the procedure to be quantitative. Clearly, application of the method to new samples will necessitate investigation of possible interferences from other species therein. Analytical Results for Potable Waters. Samples of tap water, well water, and an artificially prepared hard water containing 25.0 ppm Ca2+and 20.2 ppm Mg2+ were analyzed by the present method and by the standard EDTA titration method. For the polarographic method, the starting solutions were adjusted to the linear concentration range with Mg2+and the sample sizes chosen to keep the final concentrations less than 80 pM (total current, 80 PA). The polarogram was recorded for the starting solution and again after sample addition. T h e change in the maximum current for these two runs, divided by the slope of the calibration curve, gave the
945
increase in [Mg2+]attributable to the sample; a volumetric correction was applied. The results, shown in terms of ppm CaC03 hardness, are listed in Table I. All the reported values are the average and standard deviations of at least four determinations. Results are accurate to within 1 % of the titration values and precision is of the order of il%. Since diffusion-controlled waves are much smaller for a given reactant concentration than the catalytic wave used for this method, this technique may be used to discriminate against metal ions, which, although reducible a t these potentials, do not contribute to the catalytic wave, e.g., alkali metals and NH4+, making it particularly suitable for the analysis of Mg2+ and Ca2+in such samples as seawater, for which other convenient analytical techniques are not available. We determined both total hardness and [Mg2'] in a sample of seawater taken in the vicinity of Providence, R.I., by the present technique and by EDTA titration; the latter method is cumbersome for the determination of Mg2+ because of the necessity for quantitative precipitation of Ca2+and Sr2+( 1 5 ) . Both methods gave similar results for each analysis; although published values (16, 17) are higher by about 1070, this discrepancy is attributable to the coastal location from which the sample was taken (18).
LITERATURE CITED ASTM Standard Test Method D1126-67, American Society for Testing and Materials, Philadelphia, Pa. "Standard Methods for the Examination of Water and Waste Water", American Public Health Association, Inc., New York, 11th ed.. 1960, p 133. J. S. Fritz, J. P. Sickafoose. and M. A. Schmitt, Anal. Chern., 41, 1954 (1969). J. S. Fritz and J. N. King. Anal. Chern., 48, 570 (1976). C. N. Reilley, R. W. Schmid, and D. W. L.amson, Anal. Chem., 30, 953 (1958). J. W. Ross, Jr., and M. S. Frant, Anal. Chem.. 41, 1900 (1969). "Ion-Selective Electrodes in Analytical Chemistry", H. Freiser, Ed., Plenum Press, New York, 1978. Orion Research Inc., Cambridge, Mass., Instruction Manual of Divalent Ion Electrode Model 92-32, 1971. G. G. Perrauk. "Encyclopedia of Electrochemistry of the Elements". Vol. VIII, A. J. Bard, Ed., Marcel Dekker, New York. 1978, pp 277-280. H. A. Flaschka, "EDTA Titrations", 2nd ed., Interscience, New York. 1965, p 184. I.M. Kolthoff. E. B. Sandell. E. J. Meehan, and S. Bruckenstein, "Quantitative Chemical Analysis", Macmillan Co., New York. 1969, pp 805-807. L. Meites, "Polarographic Techniques", 2nd ed.,Interscience, New Yo&. 1965, p 184. I . M. Kolthoff and J. J. Lingane, "Polarography". Vol. 2, Interscience, New York, 1952, pp 494-495. R. H. Stehi, D. W. Margerum, and J. J. Latterell, Anal. Chem., 39, 1346 (1967). I. M. Kolthoff et al., ref. 11, pp 628-629. M. E. Thompson and J. W. Ross, Science, 154, 1643 (1966). M. E. Thompson and J. W. Ross, Science, 153, 867 (1966). Michael E. Q. Pilson, Professor of Oceanography, University of Rhode Island, private communication.
RECEIVEDfor review November 5 , 1979. Accepted January 29, 1980.
