Anal. Chem. 1982, 54, 1579-1582
1579
Voltammetric Determination of the Decomposition Rates of Combined Chlorine in Aqueous Solution Otis M. Evans Envlronmental Research Laboratory,
U.S. Environmental Protection Agency,
The kinetics of the aqueous reactions In free chlorine and excess ammonia were followed voitammetrlcaily in the range pH 3-12. Molar ratlos d ammonia to hypochlorite varied from 6.3 to 1.3. At pH 12 tile pseudo-flrst-order constant for the formation of chloramine, In a large excess of ammonia was 1.1 min-’. At lower pHr the rate was too fast to be measured. The disappearance of icombined chlorine was followed over the same pH range. Typical first-order constants in buffered solutions were 2.5 X IW2 mln-‘, 1.3 X I O ” min-’, and 1.Q X IO-’ min-‘ at pH 5.0, 7.5, and 11.0, respectively, depending on the dominant chloramine species. An unexpected rate minimum of 1.1 X min-’ was observed at pH 9.8, where monochioramine should be the only reactive species. The peak potential for the reduction of chloramine at the hanging mercury drop electrode (HMDE) was an S-shaped function of pH. Essentlaliy constant peak potentials of -0.52 V vs. saturated Ag/AgCi were observed at pH 8-12. At lower pHs the potential shifted ancdlcaliy indicating compoSne potentials. The predominant chloramine species at each pH were confirmed by UV spectroscopy.
The use of chloramines as disinfectants in drinking water to prevent the production of trihalomethanes (THMs) and other chlorinated organics has resulted in an expression of concern over the possibility that chloramines may not persist long enough to give sufficient biological protection in the distribution system. PU a consequence, important issues in the use of chloramines, must be resolved. Key among these are the primary disinfection species, the rates and equilibria of the pertinent chemical reactions, the effects of hydrogen ion activity (pH), temperature, and amount of chloramine added, and the contact period. The chemical literature contains little information on electrochemicalstudies of the chloramines. Heller and Jenkins (1)investigated the reduction of “free” chlorine (hypochlorite) and “combined” chlor ines a t a dropping mercury electrode with a saturated calomel reference. The electroreductions were irreversible and involved two electrons per molecule. Johannesson (2),in his study of the chloramines, reported frequent interferences frorn other electroreducible species. Marks and Bannister (3) studied the reduction of chloramine a t mercury, rotating platinum, and rotating platinum-silver microelectrodes. The “combined” chlorine is more difficult to reduce as compared to “free” chlorine, which is evidenced by a more negative shift in the half-wave potential. Amperometric titrations ((3-7) of residual chlorine have been carried out with differenit types of electrode substrate materials as the polarizable cathode. Iodide selectively reacts with residual chlorine a t different pH values to liberate iodine, which is titrated amperometrically with arsenite or phenylarsenoxide. The kinetics of the decomposition-disappearance, formation, and other intermediate reactions of chloramines have been extensively studied via wet chemical, colorimetric, and spectrophotometric methods of analysis (8-13). Most of the
Athens, Georgla 306 13
comprehensive kinetic studies have been performed in highly alkaline solutions, because the Raschig (8-11) reaction and its modifications are carried out in basic solutions. Audrieth and others (14,15) have characterized the reactions of chloramine: (a) nitrogen formation as a result of chloramine undergoing decomposition (self oxidation-reduction), (b) formation of hypochlorite (hydrolysis), and (c) reactions in which no decomposition occurs (direct reactions). In aqueous solution, hypochlorous acid or hypochlorite reacts quickly with ammonia or ammonium ions to yield monochloramine (NH2C1),dichloramine (NHCl,), and nitrogen trichloride (NC13). These products are pH dependent and are probably formed successively in solution. The reactions may be represented by the equations
-
+ NHS OC1- + NHzClOC1- + NHCI, OC1-
NH,C1+ OH-
(1)
NHCl2 + OH-
-
NCl:,
+ OH-
(3)
These reactions are extremely important when considering water and water supply disinfection with chlorine, hypochlorous acid, and/or hypochlorites (16-24). This paper presents the results of a polarographic study of chloramines (the analysis of reactions between “free” chlorine and excess ammonia). The rates of disappearance of the chloramines are followed over an extensive pH range; the identification of the species is confirmed with ultraviolet spectroscopy. The quantitative rate information was extraded primarily from the first 40-50 min of reaction time. No rigorous attempt was made to identify the decomposition products of the chloramine reactions. Product identification will be the subject of another paper.
EXPERIMENTAL SECTION Instrumentation. Unless stated otherwise the polarographic experiments described herein were performed with an EG&G Princeton Applied Research Corp. (PARC) Model 384 polarographic analyzer in the linear sweep mode. A platinum auxiliary electrode and a Ag/AgCl (saturated KC1) reference electrode were used in conjunction with the EG&G PARC Model 303 static mercury drop electrode (SMDE) in the HMDE mode. All wave forms were recorded on a HiPlot digital plotter (Houston Instruments). Ultraviolet spectroscopic measurements were performed on a Perkin-Elmer Model 356 two-wavelength double beam spectrophotometer. Spectra were recorded on a Coleman strip chart recorder (Hitachi 165). Reagents. Water used in this study was passed through a reverse-osmosis unit, followed by a mixed-bed resin, resulting in water with a resistivity of approximately 15 MR. The water is further treated by passing through several Illinois Water Treatment Co. (IWT) cartridge ion exchangers for removal of any residual ionizable constituents, organics, free chlorine and chloramines, phosphate complexes, and turbidity. Analytical reagent grade chemicals were used throughout the experiments. Polarograms were recorded at room temperature. All solutions were degassed at least 20 min with high-purity nitrogen prior to undertaking polarographic experiments. The nitrogen used was passed through three solutions to ensure the absence of oxygen and to ensure that the gas was saturated with
This artlcie not subject to U S . Copyright. Published 1982 by the American Chemical Soclety
1580
ANALYTICAL CHEMISTRY, VOL. 54, NO. 9, AUGUST 1982 10ti0.0,
A c e t a t e b u f f e r KHJ'0,-&HPO,
-o.61
buffer
Na,CO,-NaHCO,
I
I
I
I 0 000 0 0 0
800.0'
a
buffer
600.0]
~iiu
4
,om-. '
moo-; J
1
NHQ
oo
~
00
-03
-6s
T-
-09
-iz
-i5
-ia
E(\'), vs Ag/LgCl Flgure 1. Stationary electrode voltammogram for the reduction of 2.82 X M chloramlne (NH,CI).
water: (1) H20, (2) CrC12 in HCl plus Zn(Hg), and (3) HzO. Procedure. Prior to analysis, standard sohtions of chlorine M were prepared by diluting a at a concentration of 7.9 X 4 4 % laboratory grade sodium hypochlorite solution. The solutions were standardized iodometrically against a 0.10 N solution of standard sodium thiosulfate. Standard solutions were prepared fresh each day and stored in the dark to prevent degradation. The presence of chloramine was confirmed by using the DPD ferrous titrimetric method (25). Polarographic analyses were carried out in cells containing 10 mL of the appropriate buffer solution. A predetermined amount of ammonia was added to the cell followed by accurately dispensed quantities of free chlorine solution. The species were mixed and a polarographic curve recorded immediately. Subsequent curves were obtained at 3-10-min intervals for approximately 4-5 h. The M. The free quantity of ammonia remained fixed at 8.9 X to 7.1 X lo4 M chlorine concentrations varied from 1.4 X (at a specified pH) yielding molar ratios of ammonia to free chlorine of 6.3 to 1.3. The choice of the initial or starting potential was dependent on the pH of the solution. As the pH became more acidic, the initial potential was shifted in an anodic direction. The scans covered a range of approximately 1.5 V at a scan rate of 50 mV/s. Polarograms of solutions were recorded with the cell contents unprotected and-protected from the light with tape; this had no observable effect on the rate of disappearance of the chloramines. When the order of addition of the reagents was reversed, there was no apparent effect on the rate of disappearance, rate of formation, or the relative heights of the chloramine peaks. Since the rate of formation of the species at certain pH values could be monitored, a maximum of nine polarograms were recorded, stored, and subsequently recalled for analysis. Peak heights and positions were determined by the controlling polarographic analyzer. The existence of the particular species of combined chlorine was confirmed by ultraviolet spectroscopy.
RESULTS AND DISCUSSION Figure 1 shows a stationary electrode voltammogram for the reduction of chloramine. Immediately obvious is the broadness of the wave which attests to the irreversible nature of the electrochemical reaction at mercury electrodes. Effect of pH. The change in E, of chloramine was found to be dependent on the pH of the buffered solution (supporting electrolyte). A plot of E , vs. pH (Figure 2) yields an S-shaped curve. This S-shaped relationship has been reported for the polarographic reduction of various types of acids involving organic carbon-halogen bond fission (26-31). To our knowledge this is first report of an S-shaped relationship between pH and E , for the polarographic reduction of chloramines, involving inorganic nitrogen-chlorine bond cleavage. This curve has essentially the same shape as the theoretical chloramine equilibrium curve reported by Chapin (32). Analysis of this theoretical curve suggests that there would be equal amounts of mono- and dichloramine around pH 7.0.
/
I
3
4
NH,CI/NHCl,
I I
5
I
1
I
'
,
6
7
NH.C1
, 8
I
,
I
' 0 1 1
9
I
I
1 2 1 3
PH Flgure 2. Chloramine peak potential vs. pH. ']Acetate
buffeq
KH8P0,-&HP0. buffer 1
Na,CO,-NaHCO.
buffer
- ti 0 4
6
7
8
9
10
11
'2
PH Flgure 3. Chloramine observed firstorder rate constant vs. pH at room temperature.
Results in this laboratory indicate that, in most instances, this theory does not apply. The dichloramine always occurred in significantly smaller amounts. In the region from pH 8 to 12 the peak potential is pH independent. This indicates the presence of only one species in solution. (Monochloramine is known to exist in this range (1,8-10,32,33).) Beginning at approximately pH 7.5 the peak potential shifts in an anodic direction with decreasing pH (Figure 2). As a result of this shift in peak potential the initial potential was also shifted anodically. This necessary shift in potential had no apparent effect on the solution equilbria at intermediate pH values; however, a t lower values the polargrams became distorted and could not be interpreted. Reasons for this will be discussed later. The gradual slope of the E,-pH curve in the range from pH 4 to 7.5 is indicative of "composite potentials", Le., more than one component in solution. The broad waves made resolution impossible. In the region from pH 5.5 to 7.5 from three to four components were observed. Because of the extent of overlapping, the peak heights and peak positions of the reduction waves were greatly influenced. As the predominant species were monitored through successive scans, the minor species were noted to also disappear. With the disappearance of these minor components the peak potential of the predominant component shifted slightly and remained constant until it disappeared from solution. Disappearance of Chloramine. The rate of disappearance of chloramine was monitored in acetate, phosphate, and carbonate buffer solutions, respectively. The rate constant-pH curve (Figure 3) exhibits minima at pH 7.5 and 9.8. The first minimum occurs in the neutral pH region where dichloramine is first seen in solution. As the pH is decreased the proportion
ANALYTICAL CHEMISTRY, VOL. 54, NO. 9, AUGUST 1982 100,
Table I. Observed First.Order Rate Constants for the Disappearance of Varying Concentrations (1 X lo-“to 7.2 x lo4M ) of Combined Chlorine at 25 i 0.6 “C (Room Temperature) PH nu Kobsd ?: 117,~min-’ 5.5 6.0 6.5 7.0 7.5 8.0
8.5 9.0 9.5 9.8 10.5 11.0 11.5 12.0
3 2 3 3 4 2 2 3 3 3 3 3 3 3
2.39 X lo-’ ?: 1.15 X 2.04 x lo-’ f 2.83 X 1.63 X lo-’ ?: 4.71 X 1.36 X lo-’ 5 3.46 X lo-’ 1.31 X lo-’ 8.60 X lo-’ 1.62 X loT2f 1.41 X 2.37 x 1 0 - 2 ?: 2.12 x 10-4 1.85 x io-’ ?: 7.35 x 1 0 - ~ 1.36 x f 4.87 x 1.22 x 1 0 - 2 1.03 x 10-3 1.82 X lo-’ i 5.13 X lo-‘ 1.86 X lo-’ ?: 1.15 X 1.86 x 1 0 - 2 f 1.10 x 10-4 1.97 X lo-’ i 2.08 X
1581
+_
”” -05
-03
-01
-07
-11
-09
-15
-13
-17
E(1-), vs. Ag/AgCl Flgure 4. Variation of chloramine (2.82 X M ”&I) peak current with time: (A) 0 min, (B) 20 min, (C) 30 min, (D) 40 min, (E) 140 min, (F) 250 min.
a Number of different concentrations of combined l u = one standard chlorine used in the determinations. deviation.
of dichloramine is thouglht to increase in solution. It is a highly unstable species; consequently, the overall (heterogeneous) rate of disappearance alf chloramines is enhanced as the solution is made more acidic due to the change in the form of the predominant solution species. The equilibrium between “free” chlorine and monochloramine is reached rapidly at pH >7.5 and the solutions are more stable than at lower pH. The equilibrium between monochloramine and dichloramine and between dichloramine and trichloramine is reached slowly. The occurrence of the second minimum at pH 9.8 was totally unexpected. It seems reasonable to expect a constant rate of disappearance for the chloramine from pH 8.0 to pH 12.0, because this regioin consisted of pure monochloramine. It is possible that the piH a t the electrode surface is not the same as that in the body of the solution. The electrode reaction consumes hydrogen ions and therefore the pH at the electrode surface is probably higher than that of the bulk solution. The carbona te-bicarbonate buffer equilibrium is perhaps too slow to maintain a constant pH at the electrode surface. M) was A predetermined amount of ammonia (8.9 X added to deaerated, buffered, supporting electrolyte solutions. Varying amounts of hypochlorite were added ranging from 1.4 X M to 7.1 X I.W4 M yielding excesses of ammonia from 6.3- to 1.3-fold. ECarlier researchers reported yields of chloramine varying boltween 50 to 60% of theory when equimolar amounts of hypochlorite and ammonia were allowed to react (34). Our analyses showed that the chloramine content for these particular solutions were initially higher when ammonia was in excess, giving yields that approached 90% of theory. These solutions appeared to decompose much faster than when the chloramine was prepared from stoichiometric quantities of ammonia and hypochlorite. The rate of disappewance for five different concentrations of chloramine was monitored from pH 3 to 12 in increments of 0.5 pH units. From pH 6 to 12 the rate constants of disappearance for all concentrations of chloramine were independent of initial concentration at a given pH value (Table I). In pH 9.8 Na2C03-WaHC03buffer typical first-order rate constants for the disappearance of chloramine were 6.6 X lo4 min-’, 6.7 X min-l, and 6.6 X min-l for chloramine concentrations of 5.6 X M, 4.2 X M, and 2.8 X M, respectively. Figure 41 shows selected scans that depict the disappearance of Chloramine over approximately a 4-h period. T h e reagents were mixed and a scan was obtained immediately and at 5-min intervals for 1h, followed by 10-min intervals for the next 3 h.
,i
3-.16@7
)iio0
-.14@
~
100
0
, 3cc
,
,
0-3,
,
100
0
Time (min)
,
,
,
-
3cc
203
Time (nun)
Flgure 5. Disappearance of “free”and “combined”chlorine in pH 5.0 acetate buffer, 7.1 X lo4 M hypochlorite and 8.9 X lo4 M ammonia: (A) peak potential vs. tlme, (B) first-order plot for the disappearance of “free”and “combined”chlorine; free chlorine (open circles), comblned chlorine (open squares), “composite”species (closed circles).
-.080
i
0
, 23 4 C
,
60
?C
10$12014C
Time (rnin)
103, 0
,
,
,
2C
40
60
,
I
,
,
80 100120‘4@
Time (min)
Flgure 6. Disappearance of “free”and “combined”chlorine in pH 4.5 acetate buffer, 7.0 X lo4 M hypochlorite and 8.9 X lo4 M ammonia: (A) peak potential vs. time, (B) first-order plot for disappearance of “free” and “combined”chlorine: free chlorine (open circles),combined chlorine (open squares), simultaneous existence of free and combined chlorine (closed circle and closed square).
For pH values below 6 the formation and disappearance of chloramine were more complex. The mixtures passed through rapid alternate formations of the chloramine, disappearing and reappearing as free chlorine. The chloramine formed and stabilized before its subsequent disappearance from solution. A t pH 5.0 (acetate buffer) (Figure 5) the chloramine appeared initially and then disappeared. The “free” chlorine peak was then observed for approximately 50 min. Especially noticeable is a gradual transition to a “composite” species that ultimately gives way to the chloramine. The disappearance of the chloramine was determined to be first order. Lowering the pH to 4.5 further illustrates that the chloramine, which may form very fast initially, disappears, followed by the presence of “free” chlorine in the predominant amount. At this pH alone, it is possible to detect simultaneously both “free” and “combined”chlorine (Figure 6). This simultaneous
1582
ANALYTICAL CHEMISTRY, VOL. 54, NO. 9, AUGUST 1982 - 28ra
at 245 nm. It was the only species in the pH range 8-12. Mono- and dichloramine were present in the range from pH 5.0 to 7.5. The wavelength of maximum absorption of chloramine was approximately 297 nm. Trichloramine appeared below pH 4.0 with its maximum absorption wavelength at 223 nm. Trichloramine was observed at pH 5 also when the ammonia-to-chlorine ratio was almost equal. Its residence time in solution ranged from 3 to 5 min.
m
C
A
- 05c-
Time (min)
mid
:
ACKNOWLEDGMENT
r-7 3
10
23
30
40
50 60
Time (min)
0
10
20
3:
40
50
EO
Time (min)
Flgure 7. Disappearance of “free” and “comblned” chlorine in pH 4.0 acetate buffer, 7.0 X IO-‘ M hypochlorlte and 8.9 X IO4 M ammonia: free chlorine (open clrcles), “combined chlorine” (open squares): (A) peak potential vs. tlme for “free” chlorine, (8)peak potential vs. time for “combined’ chlorine, (C) flrst-order plot for the disappearance of “free” and “comblned” chlorine.
The author expresses his gratitude to David M. Cline, Frank G. Lether, Gary Kurth, and Steven Hodge for their assistance in making this manuscript possible and to Ronnie Moon for his patience and exactness in the construction of the figures. Thanks also to Ken Morris for the use of the ultraviolet spectrometer. I
LITERATURE CITED Heller, K.; Jenklns, E. N. Nature (London) 1948, 158, 706. Johannesson, J. K. Chem. Ind. (London) 1958, 97, 97-98. Marks, H. C.; Bannister, Q. L. Anal. Chem. 1947, 19, 200-204. Marks, H. C.; Glass, J. R. J.-Am. Water Works Assoc. 1942. 34.
1227-1240.
existence is short-lived before complete conversion to the chloramine occurs. The results at pH 4.0 (Figure 7) illustrate dramatically the effect of the transition from “free” chlorine to chloramine. There is an abrupt disappearance of “free” chlorine followed by the slow formation of the chloramine, which eventually begins to disappear from solution. At pH 3.5 and 3.0 “free” chlorine disappears rapidly with no other reactive species observed. The rate of disappearance of free chlorine is first order. As mentioned previously, voltammetric scans were difficult to obtain in this lower pH range. The initial potential was necessarily shifted anodically with decreasing pH. The result was oxidation of the electrode with the subsequent formation of calomel. Formation of Chloramine. The reaction between ammonia and hypochlorite (eq 1)is very rapid, such that the rate is measurable only a t very low concentrations. Under our experimental conditions (4-6delay response of the instrument and the relatively slow scan rates), it was difficult to measure the rate of formation of chloramine. Pseudo-first-order rate constants measured for the formation of chloramine with a large excess of ammonia were 1.1 mi& at pH 4.4 and 0.72 min-’ at pH 12.0. Measurement of rate constants at pH values inside this range was not possible. Solutions containing ammonia to which “free” chlorine was added showed an instantaneous appearance of monochloramine and dichloramine. The order of addition of the reagents had no apparent effect on the rate of formation. At pH 4.4, depending on the ammonia to chlorine ratio it was possible to observe the formation of chloramine. As the ammonia-chlorine ratio grew smaller it was possible to monitor “free” chlorine initially before the formation of chloramine. Spectroscopic Confirmation. Ultraviolet spectroscopy was used to confirm the existence of the chloramines at the various pH values. The wavelength region examined was from 380 to 200 nm. Monochloraminegave an absorption maximum
Haller, J. F.; Llstek, S. S. Anal. Chem. 1948, 2 0 , 639-842. Marks, H. C.; Wllllams, D. B.; Glasgow. G. U. ,/.-Am. Water Works Assoc. 1941, 43, 201-207. Wllllams, D. B. Water Sewage Works 1951, 98, 429-433. Colton, E.; Jones, M. M. J . Chem. €doc. 1955, 32, 485-487. Drago, R. S. J. Chem. Educ. 1957, 34, 541-545. Anbar, M.; Yagll, 0.J . Am. Chem. Soc.1982, 8 4 , 1790-1796. Anbar, M.; Yagll, G. J. Am. Chem. SOC.1982, 8 4 , 1797-1803. Berliner, J. F. T. J.-Am. Water Works Assoc. 1931,23, 1320-1333. Kleinberg, J.; Tecotzky, M.; Audrleth, L. F. Anal. Chem. 1954, 2 8 ,
1388-1389. Audrleth. L. F.; Rowe, R. A. J . Am. Chem. SOC. 1955, 77,
4726-4829. Audrleth, L. F.; Schelbler, U.; Zlmmer, H. J. Am. Chem. SOC.1958, 78, 1852-1854. Norman, T. S.; Harms, L. L.; Looyeng, R. J.-Am. Water Works AsSOC. 1980, 72, 178-180. Duke, D. T.; Slria, J. W.; Burton, B. D.; Amundsen, D. W., Jr. J.-Am. Water Works Assoc. 1980, 72, 470-478. Kim, B. R.; Snoeyink, V. L. J.-Am. Water Works Assoc. 1980, 72,
468-490. Wllllams, D. 6. J.-Am. Water Works Assoc. 1963,55, 1195-1205. Palin, A. T. J.-Am. Water Works Assoc. 1983, 55, 1205-1208. Laubusch, E. J. J.-Am. Water Works Assoc. 1983, 55, 1208-1209. “Dlscusslon” J.-Am. Water Works Assoc. 1981,73, 63-64. Hubbs, S . A.; Amundsen, D.; Olhius, P. J.-Am. Water Works Assoc.
1981, 73,97-iOl. Shull, K. E. J.-Am. Water Works Assoc. 1981, 73. 101-104. “Standard Methods for the Examination of Water and Wastewater”. 14th ed.;American Public Health Association: Washlngton, DC, 1975; pp 309-345. Elving, P. J.; Komyathy, J. C.; Van Atta, R. E.; Tang, C. S.; Rosenthal. I. Anal. Chem. I9S1, 2 3 , 1218-1223. Elving, P. J.; Rosenthal, 1.; Kramer, M. K. J . Am. Chem. SOC.1951, 73, 1717-1722. Elving, P. J.; Hlbn, C. L. J. Am. Chem. SOC.1951, 74, 3366-3371. Elvlng, P. J.; Tang, C. S. J . Am. Chem. SOC.1952, 74, 6109-8112. Rosenthal, I.; Tang, C. S.; Eking. P. J. J. Am. Chem. Soc. 1952, 74,
8112-6113. Elvlng, P. J.; Tang, C. S. J. Am. Chem. SOC.1950,72, 3244-3246. Chapln, R. M. J . Am. Chem. SOC.1929, 51, 2112-2117. Chapln, R. M. J. Am. Chem. Soc. 1931, 53,912-920. Rowe, R. A,; Audrleth, L. F. J . Am. Chem. SOC.1954, 78, 583-584.
RECEIVED for review January 20, 1982. Accepted April 23, 1982. Mention of trade names of commercial products does not constitute endorsement or recommendation for use by the U.S.Environmental Protection Agency.
