Anal. Chem. 1904, 56, 1449-1451
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Cyclic Voltammetry of Organic and Inorganic N-Chloramines in Aqueous Solution Frank E. Scully, Jr.,* Donald M. Oglesby, and Henry J. Buck Alfriend Chemical Laboratories, Old Dominion University, Norfolk, Virginia 23508
Aqueous solutlons of organic and lnorganlc Nthloramlnes as well as hypochlorlte were examlned by cyclic voltammetry at pH 8 and In strong acld (pH C2) wlth platlnum and glassy carbon electrodes. Only hypochlorlte gives a distinct reduction wave at pH 8 at a platlnum electrode. At pH 2 In 0.2 M KCI, aqueous CI, Is reduced at about +0.8 V, N-chlorodlalkylamlnes are reduced at about 4-0.2 V, and N,N-dlchloroalkylamlnes are reduced at about -0.5 V vs. SCE. The lntermedlacy of RNH# wlth a reduction potentlal of +0.05 V was detected In the formatlon of RNCI, from RNHCl In acld. The lnorganlc N-chloramines were characterized In 1 M HC104. NHCI, Is reduced at about -0.5 V and Its converslon to NCI, can be followed by appearance of a wave at 4-0.5 V vs. SCE.
The organic and inorganic N-chloramines are residual oxidants produced along with hypochlorite in the disinfection of natural water and wastewater by chlorine. Most of the electrochemical studies of these compounds have concentrated on their quantitation by amperometric titration (1-6). For analysis of chloramines, however, this usually h V O l V 0 8 addition of iodide to the solution and measurement of the iodine formed rather than direct analysis. Several mechanistic studies of the electrochemical reduction of hypochlorite (7-10) have appeared, ahd a few have examined the corresponding reduction of chloramines (2,11-13). However, most of these studies have been carried out at or near neutral pH. Recently, Evans has reported a voltammetric study of inorganic chloramines from pH 4 to 12 (13). This paper describes the results of voltammetric studies of aqueous chlorine and the Nchlorinated derivatives of ammonia, primary amines, and secondary amines in aqueous solution at low pH. Many of the compounds studied were chosen because their presence in chlorinated water and wastewater has been strongly implicated (aqueous chlorine, inorganic di- and trichloramine, N-chloromethylamine, N-chloropiperidine, N,Ndichloromethylamine, N,N-dichloro-n-propylamine, and N,N-dichloroglycylglycine). Other chloramines were examined in order to determine whether the observed reduction potentials are the same for an entire class of compounds such as the dichloramines or N-chlorammonium ions despite changes in the carbon backbone of the oxidants such as cylic vs. acyclic or long-chain vs. short-chain chloramines. Several chemical reactions must be considered in any investigation of the formation of residual oxidants on water chlorination (14, 15).
+ HzO = HOCl + H+ + C1HOCl = H+ + C10HOCl + RzNH = RzNCl+ H20 HOCl + RNHz = RNHCl + HzO HOCl + RNHCl = RNC1, + HzO Clz
RNHCl + H+ = RNH2C1+
(1) (2)
(3) (4)
(5) (6)
+ RNH2C1+= RNC1, + RNH3+ RzNCl + H+ = R2NHC1+
RNHCl
(7) (8)
Equations 1 and 2 describe the reaction of chlorine with water. At low pH and in the presence of high chloride ion concentration, dissolved Clz is the principal oxidant (16). At higher pH, hypochlorous acid, a weak acid, exists; At pH 8, it is dissociated and C10- is the major oxidant. Amino nitrogen components of natural water systems including ammonia as well as both primary and secondary amines react rapidly and essentially quantitatively with either hypochlorous acid or hypochlorite (eq 3 and 4) to form mono-N-chloramines (14, 15, 17). The monochlorinated amino compounds are sufficiently basic that below pH 3 they are protonated (eq 6 and 8) (18). In addition, in acid solution primary amines can be dichlorinated (eq 5) or undergo disproportionation (eq 7 ) to form N,N-dichloramines (14). Ammonia, which has a third exchangeable hydrogen, forms both dichloramine (NHCl,) and nitrogen trichloride (NC13) in sufficiently strong acid. Consequently, the proper adjustment of pH and excess chlorine provides a spectrum of different chlorine oxidants ranging from hypochlorite and monochlorinated amines a t high pH to hypochlorous acid, perchlorinated amines, and their protonated counterparts in acid solution. It might, therefore, be anticipated that these various oxidants undergo electrochemical reduction over a broad range of potentials.
EXPERIMENTAL SECTION Cyclic voltammograms were obtained with a Princeton Applied Research Model 174.4 polarographic analyzer in conjunction with a EG&G PARC Model 175 universal programmer. The voltammograms were recorded on a Hewlett-Packard 7049 X-Y recorder. A platinum electrode with an area of 0.12 cm2 pretreated with aqua regia and a PAR glassy carbon electrode with an effective area of 0.289 cm2were used as the working electrodes. The glassy carbon electrode was polished with 900 grit alumina before each voltammogram. This was necessary to obtain reproducible results. Current-potential curves were obtained at ambient temperature (26 1 "C) vs. a saturated calomel electrode (SCE), with a platinum auxiliary electrode. A scan rate of 50 mV/s was used in all cases except where otherwise specified. The effective surface area of the glassy carbon electrode was determined from a plot of it'/' vs. t , where i is the current obtained at time, t , from the oxidation of a solution of 1.132 X M ferrocyanide in 0.1347 M KCl at pH 2 at a potential 0.2 V greater than the peak potential according to the method of Adams (19). The graph was extrapA @,was used olated to zero time and the zero value, 4.55 X in the equation
*
it1/' = (54.5 X 103)nAD1i2Cb (9) to calculate the effective surface area, A , where n is the number of electrons involved in the oxidation of ferrocyanide. The diffusion coefficient, D, was extrapolated from the data of von Stackelberg (20) and Cb is the concentration of ferrocyanide in the bulk solution. The effective surface area was calculated to be 0.289 cm2. The area based on the diameter of the electrode was 0.29 cm2. The cyclic voltammograms of solutions of chloramines and aqueous chlorine were routinely recorded without deoxygenating the solutions. In general, bubbling nitrogen through solutions of the compounds studied resulted in loss of the compounds due
0003-2700/84/0356-1449$01.50/00 1984 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 58, NO. 8, JULY 1984
to their volatility. However, since oxygen was suspected of catalyzing the reduction of the dichloramines, one experiment involving Nfl-dichloromethylamine was performed, using every effort to exclude oxygen; consequently, the pH 2 buffer was deoxygenated by bubbling with nitrogen for 30 min. In addition, the solutions of hypochlorite and methylamine hydrochloride were deoxygenated by bubbling with nitrogen for 15 min. N-Chloromethylamine was then prepared by tangential mixing as described below for the inorganic chloramines and collected under a blanket of nitrogen. The dichloramine was prepared by diluting the chloramine with pH 2 buffer. The cyclic voltammograms of the resulting solution were recorded under a blanket of nitrogen. Stock Solutions of Organic Chloramines. All amines were from commercial sources and were of the highest purity available. If contamination was suspected, the amines were distilled from a small amount of lithium aluminum hydride before use. Methylamine hydrochloride was used instead of the free amine. N-Chloropiperidine and N-chlorodiethylamine were prepared by methods previously described (21) and the pure compounds were used to prepare M stock solutions. Stock solutions of other chloramines were prepared by slowly adding with stirring 100 mL of a solution of standardized sodium hypochlorite (2 X M) to 100 mL of an aqueous solution of a 5- to 10-fold molar excess of amine or amine hydrochloride. If the basicity of these stock solutions exceeded the capacity of the buffer used to prepare the working solutions, then the pH of the stock solution was adjusted with concentrated HC1 before diluting with buffer. Working Solutions of Organic Chloramines. Stock chloramine solutions were pipetted into volumetric flasks and diluted to the mark with the proper buffer solution to give a working solution of the desired concentration. A pH 8 buffer was prepared by adjusting the pH of a 0.1 M solution of potassium phosphate with sodium hydroxide. A pH 2 buffer solution was prepared by adjusting the pH of a 0.2 M solution of KCl with concentrated HCl. Solutions of Inorganic Chloramines. Solutions of inorganic chloramine (NH2Cl)were prepared by rapidly mixing a solution (30 mL) of ammonium sulfate (6.8 X M in ammonium ion) with a solution (30 mL) of sodium hypochlorite (5.1 X M). To achieve intimate mixing, two 30-mL syringes charged with solutions of the reactants were discharged rapidly into a Y-connector from which the reaction mixture was fed into a receiving flask. Total mixing time was less than 3 s. Aliquots (10 mL) of this stock solution were diluted to 100 mL with 1.1M HC104 The formations of NHClz (A,,= = 294 nm) and NC13 (Arnm = 336 nm) were monitored by following the UV absorption spectra of the working solutions simultaneouslywith their electrochemistry. The concentration of oxidant before and after mixing and acidification was determined by quenching an appropriately sized aliquot (0.50 mL, 1.0 mL, or 10 mL) in aqueous KI at pH 4 and titrating the resulting solution with a 0.00564 N solution of phenylarsine oxide to an amperometric end point (22).
RESULTS AND DISCUSSION At pH 8 in a phosphate buffer the cyclic voltammograms of solutions of sodium hypochlorite show an irreversible re0.45 V vs. SCE, using a duction wave at approximately stationary platinum disk electrode. The peak current is proportional to the hypochlorite concentration and the reM with a slope M to 1 X sponse is linear from 2 x of (2.46 & 0.008) X lo4 pA/M and an intercept of 1.28 0.07 pA ( r = 0.998). With a glassy carbon electrode, only a large background current is observed with no distinctive reduction peak from +0.8 V to -0.5 V vs. SCE. However, the organic chloramine, N-chloropiperidme, shows no distinctive reduction peak at pH 8 either at a platinum or a glassy carbon electrode. At pH 2, primary amines are N,N-dichlorinated, secondary amines are monochlorinated, and ammonia is di- and trichlorinated. In addition, N-chloro secondary amines are protonated so that the active oxidant is an N-chloroammonium species, R2NHC1+. These three types of compounds are electrochemically distinguishable. Figure 1shows the superimposed cyclic voltammograms of chlorine, N chloropiperidine and N,N-dichloro-n-butylamine in 0.2 M
+
*
I
I
I
08
06
OL 02 00 E vs S C E ivolkl
-02
-0L
Figure 1. Cyclic voltammograms with a glassy carbon electrode of 3.6 X lo3 M aqueous chlorine (- - -), 3.0 X lo3 M N-chloropiperidine (-), and 1.4 X lob3M N,N-dichloro-n-butylamine(-e-) in 0.2 M KCI at pH 2.0. Scan rate = 50 mV/s.
aqueous KCl at pH 2.0 using a glassy carbon electrode. It is clear that there are three distinct reduction potentials for the three different types of oxidants. Solutions of aqueous chlorine give an irreversible reduction wave at approximately +0.75 V vs. SCE. On the other hand, cyclic voltammograms of N-chloropiperidine, N-chloropyrrolidine, and N-chlorodiethylamine at pH 2 all show irreversible reduction waves at approximately +0.2 V vs. SCE. Primary amines which have two exchangeable hydrogens form N,N-dichloroamines in acid solution by the reactions shown in eq 5 and 7. Shortly after mixing a solution of aqueous sodium hypochlorite with a 10-fold excess of npropylamine and diluting with 0.2 M aqueous KC1 (adjusted to pH 2.5 with HCl), cyclic voltammograms of the resulting solution exhibit two irreversible reduction waves, one at +0.05 V vs. SCE similar to that of monochlorammonium ions (R2NHC1+)and another a t about -0.4 V vs. SCE. For a solution containing l X M active chlorine, the latter peak increases to a maximum within half an hour with equivalent disappearance of the +0.05 V wave. See eq 6 and 7. The N-chloro-n-propylamine generated at basic pH becomes protonated when the pH is dropped to 2.5 and the protonated monochloramine is then reducible a t a potential similar to dialkylchlorammonium ions. The N-chloro-n-propylammonium ion, not only a strong oxidant but also a potent chlorinating agent, converts its neutral counterpart to a dichloroamine which is reduced a t a more negative potential. The peak potentials of the N,N-dichloroamines in 0.2 M KC1 at pH 2.0 vary slightly with the structure of the dia t -0.35 V, N,Nchloramine: N,N-dichloro-n-butylamine dichloro-n-propylamine a t -0.40 V, and both N,N-dichloroethylamine and NJV-dichloromethylamine at -0.55 V, all vs. SCE. Two features of the voltammograms are worthy of note. First, the peak potential becomes more positive as the carbon chain length increases. Secondly, the cyclic voltammograms of both N,N-dichloro-n-propylamineand N,N-dichloro-nbutylamine contain a shoulder following the initial reduction wave at about the same potential as the peak maxima for both N,N-dichloromethylamine and N,N-dichloroethylamine. Coincidentally, oxygen is also reduced at about -0.5 V vs. SCE and subtraction of the background from the voltammograms completely eliminates this shoulder. The presence of oxygen in the solution is not essential for obtaining cyclic voltammograms of the dichloramines. When all efforts are made to exclude oxygen in the handling of reagents, the preparation of the dichloramine, and recording of voltammograms, there is no diminution of the electrode response. Solutions of N,N-dichloro derivatives of sec-butylamine and the dipeptide glycylglycine behave similarly to those of the n-butyldichloramine. The peak currents from the cyclic voltammograms of chlorine, N-chloropiperidine, and N,N-dichloromethylamine are proportional to the concentrations of these species. Evaluation of the peak current-concentration data by the
ANALYTICAL CHEMISTRY, VOL. 56, NO. 8, JULY 1984
method of least squares showed the slope of the response curve for aqueous chlorine is (7.31 f 0.04) X lo4pA/M (r = 0.993) with an intercept of 6.94 f 1.14 pA. Similarly, N-chloropiperidine gives a slope of (5.85 f 0.05) X lo4pA/M (r = 0.993) with an intercept of 2.15 f 1.27 pA and N,N-dichloromethylamine gives a slope of (1.11f 0.01) X lo5 pA/M (P = 0.998) with an intercept of -8.53 f 2.01 pA. Because of a combination of a large background current and broad reduction wave, the lower limit of detectability in the case of all three is about 2 X 10" M where the peak current is about twice that of the background current. Since oxygen is reduced at a potential similar to that of the dichloramines, its presence further limits the sensitivity of their electrochemical detection. Deoxygenating the solution by bubbling with nitrogen fails to enhance the sensitivity but instead leads to inconsistent variations in concentrations as chlorine or the volatile organics are purged from the solution. If the buffer solution is deoxygenated before diluting the stock solution of dichloramine, the lower limit of detectability of both N,N-dichlorois approximately methylamine and NjV-dichloro-n-butylamine 1X M and good linear response with varying concentrations can be obtained for solutions above this concentration. The inorganic N-chloramines are considerably more difficult to characterize because of their lower basicity, the presence of a third exchangeable hydrogen, and the oxidative side reactions such as hydrazine and nitrogen formation, which they seem to undergo more rapidly than the organic N chloramines. To minimize the effect chloride ion has on the equilibria involved in the formation and decomposition of the various chloramines, the electrochemistry of the inorganic chloramines was carried out in 1 M HC104. Ammonia and hypochlorite were mixed in a molar ratio of 1.3:l in the absence of a buffer. Since oxidative side reactions can compete with chloramine formation, the concentration of oxidant was determined before and after mixing by iodometric methods. In a typical run, only a few percent (0-2%) of the total oxidant was lost by using the method of mixing described in the Experimental Section. When an aliquot of this solution is diluted with 1M HC104, an irreversible reduction wave at about -0.5 V vs. SCE is observed shortly after mixing. Again in a typical run analysis of total oxidant concentration showed that 82% of the theoretical amount of oxidant remained after dilution with acid. The reaction was also followed simultaneously by UV spectrometry and an absorption maximum at 294 nm corresponding to that of NHClz (14) was observed. This fact and the similarity of the reduction potential to that of the organic N,N-dichloramines suggest that the -0.5 V wave is due to the reduction of NHC12. In that case, the -0.5 V wave is similar to the organic dichloramines and the peak current is proportional to concentration with an electrode response identical with that for N,N-dichloromethylamine. Evidence for the intermediacy of a protonated monochloramine (NH3Cl+),similar to that found when solutions of monochlorinated organic amines are pH-jumped to pH 2, is not observed. However, Gray has shown that the secondorder rate constant for the disproportionation of monochloramines (eq 7) is about 16 times faster for NHzCl than for N-chloromethylamine. This would make it sufficiently fast that NHSC1' would not be easily detected within the 4-5 min time restriction between acidification of the solution and recording of the voltammogram. Over the period of an hour after mixing and acidification, both the 294-nm absorption band and the -0.5 V reduction wave decrease with a corresponding formation and increase in a weaker absorption band a t 336 nm and a new reduction wave at about +0.5 V vs. SCE. Since the 336-nm absorption is due to nitrogen trichloride (NC13) (14),the +0.5 V wave is
1451
assigned to this compound. It is interesting to note that over the first half hour after acidification, the increase in the current at +0.5 V is equivalent to the decrease in current at -0.5 V. This is consistent with the clean conversion of the dichloramine to nitrogen trichloride without decomposition by competing reactions. After about 45 min, the change in current at the two potentials is no longer equivalent suggesting a net loss of total oxidant. A platinum disk electrode has also been examined for use in the detection of chloramines in 0.2 M KCl at pH 2.0. Hypochlorous acid (5 X M) is reduced at +0.85 V vs. SCE M) is reduced at +0.52 V and N-chloropiperidine (5 X vs. SCE. However, the working range of the platinum electrode does not extend low enough to detect the dichloramines. In addition, it gives a much greater background current than the glassy carbon electrode and so was disregarded early in these studies.
ACKNOWLEDGMENT We are grateful to J. Donald Johnson of the University of North Carolina at Chapel Hill for suggesting the use of UV spectrometry to follow the interconversion of NHClz and NC13. Registry No. Clz, 7782-50-5; NHC12, 3400-09-7; NC13, 10025-85-1;NH2C1,10599-90-3;KC1,7447-40-7;HOCl, 7790-92-3; C, 7440-44-0;Pt, 7440-06-4; N-chloropiperidine,2156-71-0;N,Ndichloro-n-butylamine,14925-83-8; N,N-dichloromethylamine, 7651-91-4;N-chloropyrrolidine, 19733-68-7;N-chlorodiethylamine, 5775-33-7;N-chloro-n-propylamine,52548-07-9;N,N-dichloro-npropylamine, 69947-01-9;N,N-dichloroethylamine,24948-83-2; sodium hypochlorite, 7681-52-9;ammonium sulfate, 7783-20-2.
LITERATURE CITED Johnson, J. D. I n "Treatise on Water Analysis"; Minear, R., Ed.; Academic Press: New York, 1983; Vol. 1. Marks, H. C.; Bannister, G. L. Anal. Chem. 1947, 79, 200-204. Marks, H. C.; Glass, J. R. J . Am. Water Works Assoc. 1942, 34, 1227-1240. Haller, J. F.; Listek, S. S.Anal. Chem. 1948, 20, 639-642. Marks, H. C.; Williams, D. B.; Glasgow, 0. 0. J . Am. Water Works ASSOC. 1941, 43, 201-207. Williams, D. B. Water Sewage Works, 1951, 9 8 , 429-433. Chao, M. S. J . Nectrochem. SOC. 1988, 175, 1172-1174. Hine, F.; Yaouda, M. J . Electrochem. SOC. 1971, 778, 170-173. Harrison, J. A.; Khan, Z. A. J . Electroanal. Chem. InterfacialElectrochem. 1971, 30, 87-92. Wolf, H.; Landsberg, R. J . Electroanal. Chem. Interfacial flectrochem. 1970, 28, 295-300. Heller, K.; Jenkins, E. N. Nature (London) 1948, 758, 706. Johannesson, J. K. Chem. Ind. (London) 1958, 97, 97-98. Evans, Otis M. Anal. Chem. 1982, 5 4 , 1579-1582. Gray, E. T.. Jr.; Margerum, D. W.; Huffman, R. P. I n "Organometals and Organometalloids: Occurrence and Fate in the Environment"; Brinckman, F. E., Bellama, J. M., Eds.; American Chemical Society: Washington, DC, 1978; ACS Symposium Series No. 82, pp 264-277. Morris, J. C. In "Prlnciples and Applicatlons of Water Chemistry"; Faust, S. D., Hunter, J. V., Eds.; Wiley: New York, 1967; pp 23-53. Johnson, J. D.; Edwards, J. W.; Kesslar, F. J . Am. Water Works Assoc. 1978, 70, 341-348. Higuchi, T.; Hussain, A.; Pitman, I. A. J . Chem. SOC.6 1069, 626. Weil, I . ; Morris, J. C. J . Am. Chem. SOC. 1949, 77, 3123-3126. Adams, R. N. "Electrochemistry at Solid Electrodes"; Marcel Dekker: New York, 1969; p 216. von Stackelberg, M.; Pilgram, M.; Toome, V. Z . Nectrochem. 1953, 57, 342-351. Scully, F. E., Jr.; Bowdring, K. J . Org. Chem. 1981, 4 6 , 5077-5081. "Standard Methods for the Examination of Water and Wastewater", 14th ed.; American Public Health Association: Washington, DC, 1976; pp 322-325.
RECEIVED for review June 14, 1982. Resubmitted March 8, 1984. Accepted March 19, 1984. This work was funded by the United States Environmental Protection Agency under assistance agreement CR-807254 with the agency's Health Effects Research Laboratory, Cincinnati, OH (F. B. Daniel, Project Officer). The information contained herein has not been subjected to the Agency's required peer and administrative review and, therefore, does not reflect the view of the agency and no official endorsement should be inferred.
