Environ. Scl. Technol. 1088, 22, 691-696
General Acid Catalysis of Monochloramine Disproportionatlon Richard L. Valentine" and Chad 1.Jafvertt Department of Civil and Environmental Engineering, University of Iowa, Iowa City, Iowa 52242
NH3C1+
a This paper presents experimental results showing that monochloramine disproportionation, which results in the formation of dichloramine, involves a general acid catalyzed reaction pathway. Rate constants characterizing the effect of hydrogen ion, phosphate, and sulfate were determined by measuring the rate of monochloramine disappearance under pH conditions, which simplified interpretation of results. These rate constants were used to develop a linear free energy relationship that was used to predict the effect of carbonate and silicate. Predictions indicate that carbonate, and possibly silicate, may significantly increase the rate of acid-catalyzed disproportionation at concentrations and pH values typical of many drinking waters. Since this reaction may govern the overall rate of oxidant loss, appropriate consideration must be given to the presence of potential proton donors when predictions relating to chloramine speciation and fate are made on the basis of reaction models or when the results of studies with chloramine solutions are evaluated.
Introduction A fundamental understanding of chloramine chemistry is important in the control of water and wastewater disinfection, biofouling in power plants, and in assessing the fate of chlorinated effluents discharged into the environment. A renewed interest in chloramine chemistry is a consequence of a desire to better understand and control reactions involved in byproduct formation resulting from chloramination disinfection, which is accomplished primarily through the action of monochloramine. Chloramines are inherently unstable even in the absence of organics with which they may react and eventually decompose via a complex series of reactions leading to the oxidation of ammoniacal nitrogen (1-7). A major reaction leading to monochloramine loss is the disproportionation of monochloramine: NH&l+ NHzCl-
kobd
NHClz
+ NH3
(1)
which leads to dichloramine formation. It is generally believed that the subsequent decomposition of dichloramine is primarily responsible for actual oxidant loss, which may be governed by monochloramine disproportionation if disproportionation is rate limiting. Monochloramine disproportionation is generally the consequence of two reaction pathways as first shown by Granstrom and Morris (8, 9). One is the hydrolysis of NHzCl producing HOC1, which then reacts with additional NHzCl to form NHCl,. The other pathway, which is the focus of this paper, is catalyzed by hydrogen ion (8-10). On the basis of the observed catalytic effect of acetic acid/acetate buffers, it has also been suggested (8, 9) to involve a general acid catalyzed mechanism, which is consistent with the reaction of a neutral monochloramine species with a protonated form (NH3Cl+) according to
"$21
+ H+ or HA F! NH,CI+ (+ A-)
(2)
Present address: Environmental Research Laboratory, U.S. EPA, Athens, GA 30613. 0013-936X/88/0922-069 1$01.50/0
+ NHzCl-
NHC1,
+ NH3 + H+
(3)
where HA is any proton donor. The effect of species other than hydrogen ion and acetic acid on monochloramine disproportionation has to date not been investigated even though several commonly occurring substances could potentially influence this reaction in natural waters. Additionally, several studies of the chlorine-ammonia system have employed buffers such as phosphate or carbonate (1,2,4-7) whose effects have not been evaluated. In this study, general acid catalysis of the disproportionation of monochloramine was studied in the presence of phosphate and sulfate by measuring monochloramine disappearance under two sets of conditions that greatly simplified interpretation of the results. The catalytic effects of H3P04,HSO,, and H+ were studied spectrophotometrically in the pH range of 3-4 where the reaction is rapid and dichloramine is stable over the time period of the experiment. The effect of HzPOc was evaluated by a titrimetric approach to measure monochloramine disappearance in experiments conducted at pH 8 where disproportionation is slow and the concentration of dichloramine remains very low relative to monochloramine because of its rapid decomposition. The measured specific rate constants were used to develop a relationship to predict the effect of silicate and carbonate, which is particularly difficult to evaluate experimentally because of problems in controlling pH due to evolution/absorption of carbon dioxide.
Experimental Section Reagents. All solutions were made with reagent-grade chemicals and chlorine demand free water prepared by adding approximately 5 mg/L Cl, (Mallinchrodt, sodium hypochlorite) to distilled-deionized water (Millipore Milli-Q), followed by aging in the dark for 24 h and by dechlorination with a UV lamp. Phosphate solutions were prepared by appropriate addition of an aliquot of filtered and chlorinated/dechlorinated 1.0 M NaHZPO4 and 0.2 M H3P04. Sulfate solutions were made by mixing appropriate amounts of filtered and chlorinated/dechlorinated 1.0 M Na2S04and 0.05 M H2S04. Reacting mixtures of monochloramine and acidic phosphate or sulfate solutions were maintained at an ionic strength of 0.2 M by addition of NaC104 to the phosphate/sulfate stock solutions. Ionic strength was held at 0.1 M in experiments performed at pH 8. Stock monochloramine solutions were prepared fresh daily by dropwise addition of sodium hypochlorite (0.8 M) to a rapidly mixed solution of ammonium chloride adjusted to pH 9.0 by addition of NaOH. Stock concentrations of 1-2 mM were prepared in the low-pH studies, while a concentration of approximately 20 mM was used in the high-pH studies. The titer was determined using the DPD titrimetric method (11)and by UV absorbance at 245 nm measured spectrophotometry with a Model 552 PerkinElmer scanning UV-vis spectrophotometer. No dichloramine was detected. A molar absorptivity of 444 M-l cm-' was calculated for monochloramine at 245 nm. This value compares closely with others (2, 8, 12).
0 1988 American Chemical Society
Envlron. Scl. Technol., Vol. 22, No. 6, 1988
691
A solution of pure dichloramine was obtained by rapidly lowering the pH of the 1-2 mM stock monochloramine to 3.5-4.0 by addition of 1.0 M HC1 as previously described (12). The solution was then aged overnight and its UV spectra obtained. Titer was determined by the DPD titrimetric method and the molar absorptivity calculated to be 208 M-' cm-l at 245 nm, which agrees closely with that obtained by Granstrom (8). Procedures and Apparatus. The disproportionation reaction at low pH (3-4) was initiated by rapidly mixing equal volumes of stock monochloramine with acidic phosphate or sulfate solutions in an SFA-11 rapid kinetics accessory consisting of a thermostatted mixing/observation cell placed in a Perkin-Elmer Model 559 scanning UV-vis spectrophotometer. The absorbance at 245 nm was then monitored for approximately 1min, over which time about 90% of the monochloramine reacted. All low-pH experiments were conducted with an initial monochloramine concentration of approximately 1.4 mM held at 25 "C by circulating thermostatted water through the accessory and cell holder. Five consecutive experimental runs were made at each reaction condition. The pH of the mixture at the beginning and end of the experimental run was determined by a Beckman Model pH171 pH meter in conjunction with a Corning Ag/AgCl general purpose probe, a Calomel reference probe, and a Beckman automatic temperature compensating probe. The initial concentration of monochloramine as measured by its absorbance was determined by substituting water for the acidic phosphate or sulfate stock solution. Total phosphate concentration was varied from 50 to 200 mM. Total sulfate was 50 mM in those experiments that utilized sulfate. The slow loss of monochloramine in 10 and 25 mM phosphate solutions at pH 8.0 was monitored with the DPD titrimetric method (11)over a period of several days. Experiments utilized 1.5 mM monochloramine (N/Cl = 2/1) in phosphate buffer, and sodium perchlorate contained in an aluminum foil covered beaker to omit light. The solutions were stored at 25 "C. No significant concentrations of dichloramine were ever detected in the solution. Data Reduction. Constants characterizing the catalytic effect of various species on monochloramine disproportionation were evaluated from Itobsdvalues determined from integrated rate expressions describing monochloramine disappearance. Two rate expressions were utilized depending on the pH at which the experiments were conducted. Activity coefficients were calculated according to the equation of Davies (13). Acid dissociation constants were obtained from the literature (14). In the range of pH 3-4, monochloramine quickly reacts to form dichloramine, which is relatively stable (12) in this pH range. Reaction 1occurs primarily by the general acid catalyzed pathway (reactions 2 and 3) and is essentially isolated from all other possible reactions. Monochloramine hydrolysis is too slow to appreciably contribute to dichloramine formation, and trichloramine formation is also negligible. A back-reaction between dichloramine and ammonia is much slower than the acid-catalyzed disproportionation reaction (12). In this low-pH range, the species NH3C1+is expected to be less than approximately 1% of the concentration of NH2Cl (10). Under these conditions the rate of monochloramine disappearance can be expressed as d[NH,Cl] /dt = -2kobsd["&1]2 (4) where the observed second-order rate constant is given by kobsd = Cki[HAil (5) 692
Envlron. Sci. Technol., Vol. 22, No. 6, 1988
and ki is the specific rate constant for the ith proton-donating species HAi (including hydrogen ion). The factor of 2 accounts for kobsdbeing defined with respect to dichloramine formation. The observed second-order rate constants (values of kobsd) were determined from absorbance measurements at 245 nm with the integrated form of rate expression 4 (15): (A - X O ) / ( X m
- A) = 2kobsdtCO
(6)
where X = absorbance at time t , Xo = initial absorbance at t = 0, A, = final absorbance at t = 00, t = time, and Co = initial concentration of monochloramine. Five consecutive experimental runs were made and the average value and associated variance of the left-hand term in eq 6 determined at each time from the pooled data. The value of k o b d for each reaction condition was then calculated by a weighted linear least-squares analysis with these averages and their variance estimates as weighting factors. This approach was taken in order to appropriately consider data that becomes less precise with time because the term A, - X approaches zero as the reaction proceeds toward completion. The variances of the kobsd values were used as weighting factors in subsequent weighted least-squares analyses used to obtain species-specific rate constants. The standard deviation associated with each value of kobsd determined at low pH was found to be approximately 5% of its value. Calculations using phosphate data utilized results obtained over a time period in which 90% of the monochloramine disappeared and which the pH did not detectably change. Calculations to determine rate constants in the presence of sulfate utilized results obtained over a time period in which only 20-30% of the monochloramine disappeared. This was done in order to minimize the change in pH, which was estimated to increase by approximately 0.03-0.05 unit. All calculations were based on the initial pH measurements. The value of A, was calculated from knowledge of the initial absorbance in eq 7 , where ED and t M are the molar h, =
(7)
&6D/2eM
absorptivities of dichloramine and monochloramine at 245 nm, respectively. A simplified version of the chloramine decomposition model proposed by Lea0 (1,2) was used to develop a rate expression for monochloramine disappearance at pH 8.0 where redox reactions leading to oxidant loss keep dichloramine at a relatively low Concentration. The mechanism adopted incorporates rate limiting disproportionation (reaction l) and a fast redox reaction (reaction 8), NHClz + NH&l-
fast
N2 + 3C1-
+ 3H+
(8)
which accounts for the majority of oxidant loss. Reaction 8 is hypothesized to actually involve an intermediate produced by dichloramine, which reacts rapidly with monochloramine. Another proposed pathway involving a reaction of the intermediate with dichloramine is not expected to be important at the relatively high monochloramine concentrations used in this study ( I , 2). Monochloramine hydrolysis, which contributes to the overall rate of disproportionation, can be incorporated into the overall kinetics by considering an equilibrium between HOCl and monochloramine, and a rapid reaction of HOCl with NH2C1according to K.
NHzCl+ H2O r HOCl HOCl + "$21
kmd _ . +
+ NH3
NHC12 + H20
(9)
(10)
0 50 mM Phosphate
2.0
25 20
0
10
20
30
40
t
I
A/
50
Time (sec.) Flgure 1. Typical plot of absorbance results used to calculate observed rate constants (kobsdvalues) In the pH range of 3-4.
Under conditions where dichloramine is at pseudo steady state, the corresponding overall rate expression is given by d[NH,Cl] /dt = -3kobsd["&1]2 (11) where the observed second-order rate constant is given by kobsd = Cki[HAil + kmdKe/["3] (12) which incorporates both the general acid catalysis and monochloramine hydrolysis pathways. Reaction conditions were chosen, however, so that general acid catalysis was the dominant disproportionation pathway. Observed rate constants at pH 8.0 were obtained from measurements of the monochloramine concentration-time data with the integrated form of eq 11: 1 / c - l/cO = 3kobsdt (13) where Co is the initial monochloramine concentration and C is the concentration at time t. A constant pH of 8.0 was assumed for all calculations, although pH dropped approximately 0.04-0.07 unit depending on phosphate concentration.
Results Catalysis by H+ and H3P04. Figure 1 shows a typical plot of relationship 6 used to calculate kobsd values for experiments performed at approximately pH 3.0 clearly indicating that the rate of monochloramine disproportionation is increased with increasing phosphate concentration. The expected linearity also supports the initial assumption of a second-order reaction. In the absence of monochloramine hydrolysis, the observed second-order rate constant can be expressed by kobsd = kH+[H+l+ kH8P0,[H3P041+ kHzPO4-[H2P04-1(14) which accounts for the catalybic effect of hydrogen ion and all significant phosphate species between pH 3 and pH 4. Substitution of the acid dissociation constant Ka, = 10-2.1 for H3P04,with consideration of ionic strength effects, results in eq 15, where y1 = activity coefficient for mo-
novalent ions (0.75) and (H+}is the hydrogen ion activity, which can be used to evaluate kH+from the intercept of a plot of kobs. vs HzP04- concentration. The second-order rate constants obtained at approximately pH 3.0 with variable phosphate concentrations are plotted in Figure 2 as a function of H2P04-concentration. From the linear (weighted least-squares) relationship, a value for kH+ of 6.35 X lo3 M-2 s-l f 1.40 X lo3 M-2 s-l
0
25
75
50
100
125
150
175
200
H2P04- Concentration (mM)
Figure 2. Observed rate constants obtained at pH 3.0 plotted as a function of H2P0,- concentration.
25 20 v)
H
15
:
t
//
-
0
0.2
0.4
0.6
0.8
1.2
/ H + /x 10-3(M) Flgure 3. Observed rate constants obtalned as a function of hydrogen ion activity in the presence of phosphate and sulfate in the pH range of 3-4.
(95% confidence level) can be calculated from the intercept. The value of the hydrogen ion specific rate constant can be compared to a value of 7.3 X lo3 M-2 s-l given with respect to a rate expressed in terms of hydrogen ion activity obtained by Granstrom (8) at an ionic strength of 0.45 at 25 "C. On a similar hydrogen activity basis, the value obtained in this work is equal to approximately 8.47 X lo3 M-2 s-l. This value is somewhat less than the value of 2.74 x 104 M-2 8- determined by Gray et al. (IO)at an ionic strength of 0.5 and temperature of 25 "C. Rearrangement of eq 15 yields kobsd
=
(16) which can be used to calculate a value of kHgo from a plot of kobsd vs hydrogen ion activity at a fixed phosphate concentration, as shown in Figure 3 for results obtained Environ. Sci. Technoi., Vol. 22, No. 6, 1988 693
2.0
,
I
1
I
I "
0
10
20
30
43
60
60
70
80
90
100
TIME (hours)
Flgure 4. Second-order plot of monochloramine concentratlon as a fUnCtlOn of time used to calculate kobsd for results obtalned at pH 8.0 at phosphate concentrations of 10 and 25 mM.
in 100 and 200 mM phosphate over a pH range of 3.0-4.2. Over this range the concentration of H2P04- is approximately 94 and 188 mM in the 100 and 200 mM total phosphate solutions, respectively, varying by only about 6%. A kHsPO4value of 870 M-2 s-' can be calculated as the average of 880 and 860 M-2 s-l determined for each phosphate concentration from the slopes calculated with weighted linear least-squares with weighting factors equal to the estimated variances. The intercepts of 1.4 f 0.9 M-l s-l and 1.3 f 0.8 M-' s-l (95% confidence limits given) for the results obtained in the 100 and 200 mM solutions, respectively, are not believed sufficiently large or precise enough to allow their use to reliably estimate kHIP04-.Included on Figure 3 is also a plot showing the theoretical relationship between kobsdand hydrogen ion activity for a system containing no other catalytic species other than hydrogen ion. Catalysis by HS04-. The same procedure as that used to calculate the catalytic effect of H3P04and H+ can be used to determine constants characterizing the effect of sulfate. In the pH range 3-4, the primary protonated species is HS04-, which varies in concentration by only a few percent. The observed rate constant kowcan therefore be defined as kobsd
= kH+[H+l+ kHS04-[HS04-l
(17)
Substitution yields
where y1 = activity coefficient of H+ or HS04- = 0.75, y2 = activity coefficient for SO?- = 0.31, and Ka2= M, which can be used to obtain an estimate of kHS04-from a plot of kow vs hydrogen ion activity. A plot of relationship 18 is also shown on Figure 3 for data obtained in the , presence of 50 mM total sulfate. The value of k ~ s o was determined to be 1.72 X lo3 M-2 s-l from the slope of 1.2 X lo4 M-2 s-l and the previously measured value of kH+. The effect of higher concentrations of sulfate was not investigated because the constant ionic strength assumption would be excessively violated. Catalysis by H2P04-. Figure 4 shows data obtained at pH 8.0 plotted according to eq 13 from which kow values and 2.64 X lo4 M-l s-l were calculated from of 5.83 X the 25 and 10 mM total phosphate data, respectively. The dominant catalytic effect of phosphate is evident from the ratio of the slopes, which is 2.2. This ratio is fairly close to the value of 2.5 expected if hydrogen ion catalysis and monochloramine hydrolysis had negligible effects on the 694
Environ. Scl. Technol., Vol. 22, No. 6, 1988
overall rate of disproportionation. The results are consistent with calculations using values of K, and kmdof 6.7 X M-' and 1.5 X lo2 M-l s-l, respectively, obtained at an ionic strength of 0.5 (10, 16), which indicate that monochloramine hydrolysis occurring in the presence of 1.5 mM excess total ammonia is expected to account for only approximately 11% of the kobsdvalue measured in the presence of 10 mM phosphate and only 5 % of the value obtained in 25 mM phosphate. The effect of phosphate is assumed to be primarily due to H2P04-not HP042-because the acid dissociation constants of these species differ by 5 orders of magnitude. On the basis of the measured value of kHBPO,&PO4 is not expected to appreciably contribute to catdysis at this pH. Subtraction of the estimated contributions due to monochloramine hydrolysis (neglecting ionic strength differences) and to hydrogen ion catalysis yields an average kH2P04-value of 0.25 M-2 s-l from estimates of 0.27 and 0.23 M-2 s-l obtained in the presence of 25 and 10 mM total phosphate, respectively.
Discussion Importance of General Acid Catalysis. If catalysis by hydrogen ion results in the formation of dichloramine at a rate faster than attributable to the monochloramine hydrolysis pathway (reactions 9 and lo), then any species whose catalytic effect exceeds that of hydrogen ion must be considered important in the overall rate of dichloramine formation. It is therefore instructive to first evaluate the conditions under which the hydrogen ion catalyzed pathway dominates in the net formation of dichloramine, and then to evaluate the conditions under which the effects of other potentially catalytic species are expected to exceed that of hydrogen ion. As a first approximation applicable to conditions where chloramines slowly decompose near or above neutral pH, eq 1 2 can be used to estimate the importance of each pathway. By considering hydrogen ion catalysis alone and equating the first and second terms in eq 12, relationship 19 can be derived, where NT is the concentration of total
+
uncombined ammonia (NH3 NH4+)and Ka is the acid dissociation constant of ammonia. This equation relates the total uncombined ammonia concentration to hydrogen ion concentration under conditions where the hydrogen ion catalysis pathway and monochloramine hydrolysis pathway contribute equally to the net rate of dichloramine formation. Hydrogen ion catalysis is expected to be more important in the presence of total uncombined ammonia concentrations that exceed the value given in eq 19. Evaluation shows that dichloramine formation is expected to be due primarily to the hydrogen ion catalyzed pathway when total uncombined ammonia nitrogen concentrations exceed approximately 0.32 mM or about 5.4 mg/L as N at pH values well below the pKa of ammonia. At pH 9.3 the requisite uncombined ammonia concentration would be twice this value. These concentrations, while exceeding values of several milligrams per liter sometimes found in chloraminated and natural waters containing chlorinated discharges, are low enough to suggest that species may measurably increase the net rate of dichloramine formation in these waters if their catalytic effects are greater than that of the hydrogen ion. Furthermore, dichloramine formation may be driven primarily by general acid catalysis in experimental systems with ammonia concentrations in excess of these values.
6 5, H 0
;$\w
I
measured predicted
”
\
-5-
H3SiO;
-60
1
2
3
4
5
6
7
8
9
1 0 1 1 1 2 1 3
pKAt Log(P/Q)
Flgure 5. Linear free energy relatlonship (Bronsted plot) relating species-specific catalysts rate constants to acid dlssoclation constants. Solid squares, measured values: hollow squares, predicted. 95 % Confidence region shown. Value for acetic acid (triangle) from Granstom and Morris (8, 9 ) .
Linear Free Energy Relationship. The four rate constants determined in this study can be related to the species individual acid dissociation constant Ka by a Bronsted plot (17)(linear free energy relationship) of the form (20) log ( k i / P ) = C,[pKa + log (P/Q)I+ C2 where C1and C2 are constants, P is the number of exchangeable protons on the species, and Q is the maximum number of protons that the conjugate base could combine with, Relationship 20 is shown in Figure 5 for results summarized in Table I. As expected, the specific acid catalysis constant increases as the acid dissociation constant increases. While a reasonably linear relationship is obtained with C1 = -0.68 and C2 = 4.09, significant uncertainty is associated with use of this relationship as a predictor as indicated by the 95% confidence limit band and associated uncertainty for carbonate and silicate specific rate constants presented in Table I. In support of this relationship, the specific rate constant of 1.96 M-2 s-l obtained by Granstrom and Morris (8,9)for catalysis by acetic acid (pK, 4.7) results in a value of log ( k J P ) , which is shown to lie within the 95% confidence region. In spite of the potential uncertainty in estimating rate Constants, it is worthwhile nevertheless to use this relationship to investigate the expected effects of carbonate, silicate, and phosphate on general acid catalyzed monochloramine disproportionation. Importance of Carbonate and Silicate. The relative importance of carbonate and silicate in general acid catalyzed monochloramine disproportionation can be evaluated by examining the ratio of their contribution in the determination of komto that of hydrogen ion, which serves as a convenient reference species. Figure 6 shows the relationship between carbonate concentration and pH expected to result in several values of the ratio R where = kCARB/kHY, and kCARB = kH,C08[H2C031 + kHC08[HCO,], and km = kH+[H+]. Figure 7 is an analogous plot for the silicate system with the ratio R = kSi/kHYwhere ksi = k~,sio,[H4Si04]+ k~ sio4-[H,SiO4-]. Above the center line where R = 1, the eifect of’carbonate or silicate is expected to exceed that of hydrogen ion in governing the overall rate of general acid catalyzed disproportionation. At R = 0.1, the effect of these inorganics is 10% of that attributable to hydrogen ion catalysis while at R = 10 it is 10-fold greater. The relative importance of carbonate is seen to increase as pH increases but is somewhat insensitive to pH between 7.5 and 8.5. This is due to a compensating reduction in
both H+ and H2C03at a fixed carbonate concentration as pH increases. In effect, general acid catalyzed disproportionation at a fixed carbonate concentration could probably be adequately considered over this pH range by incorporating only hydrogen ion catalysis but with a hydrogen ion specific rate “constant” that is a function of the carbonate concentration and hence not a true constant. The shaded area in Figure 6 defines the region in carbonate concentration and pH where carbonate effects are likely to be significant in natural waters. Here, significance is defined by a carbonate effect exceeding 10% (R 2 0.1) of that attributable to hydrogen ion catalysis over a maximum pH range of 6.5-9.0. This zone has an upper bound in carbonate concentration of 5 mM, which is usually not exceeded in natural waters. As can be seen, the minimum effect level of R = 0.1 is expected to be realized over this pH range with carbonate concentrations between 1and 0.3 mM, respectively, a range exceeded in many drinking waters. The effect of carbonate is shown to be comparable to or exceed that of hydrogen ion in the pH range of 8-9 when carbonate is present at several millimoles per liter. It is then expected that the effect of carbonate could be as important as the specific hydrogen ion catalyzed route to the formation of dichloramine. As found for the carbonate system, Figure 7 shows that the relative importance of silicate increases as pH increases due to a relative decrease in hydrogen ion catalysis. However the region of significant effect (R 1 O.l), bounded by a pH of 9.0 and an upper silicate concentration of 1 mM, does not extend below pH 8. This upper bound was chosen since silicate concentrations in natural waters rarely exceed this value and are more typically in the range of a tenth to several tenths millimolar. At pH 9.0, the 10% level of relative effect is met at a silicate concentration of approximately 0.2 mM while the level of effect approaches that of hydrogen ion at a concentration of 1 mM. In comparison to the carbonate system, however, typical silicate concentrations appear less likely to significantly increase the rate of general acid catalyzed monochloramine disproportionation unless the pH approaches or exceeds approximately 9. Importance of Phosphate. Although phosphate does not normally exceed several tenths of a milligram per liter in most drinking waters, phosphate has been used over a wide range in concentrations and over a large pH range to buffer chloramine solutions used for experimental purposes ( 1 , 2 , 4 , 5 ) .Figure 8 shows the relationship between total phosphate and pH for R values of 0.1, l, and 10 where R = kpHoS/k~~ and kPHos is the sum of all phosphate catalysis terms. At low pH, phosphate becomes significant only when its total concentration exceeds approximately 5 mM, which points out why concentrations in the range of 50-200 mM were used in this study at low pH to obtain easily measured effects. In the neutral to high pH range, phosphate concentrations at the sub millimolar level are expected to contribute measurably to acid-catalyzed monochloramine disproportionation primarily by the action of H2P04-. While phosphate in natural waters is not expected to significantly increase the rate of acid-catalyzed monochloramine disproportionation, phosphate should be used with caution in any study involving chloramines. Conclusions The disproportionation of monochloramine has been shown to be catdyzed by phosphate and sulfate and should be considered a general acid catalyzed reaction. Catalysis in natural waters by carbonate, and to a lessor degree by silicate, is expected. In so far as general acid catalyzed Envlron. Sci. Technol., Vol. 22, No. 6, 1988 695
1000
Table I. Summary of Measured and Predicted Specific Rate Constantsa species
P K A ~ P/Q 0.0
2.0 2.1 7.2 6.35 10.3 12.3 9.5 12.6
1/1
112
3/1 212 2/1 1/2
1/3 411 3/2
"kiunits of M-2 s-l at level indicated.
log ki PKA log ( P / Q ) measured predictedc 0.0
3.80 3.23 2.94 -0.602
4.09 i 0.80 3.00 i 0.59 2.94 i 0.54 -0.802 k 1.03 -0.127 i 0.93 -2.70 i 1.54 -3.94 2.00 -2.17 i 1.47 -4.11 i 2.07 25 O C . bReference 13. "5% confidence 1.69 2.57 7.20 6.65 10.0 11.8 10.1 12.8
.1
I,,,
, , ,, , , ,,
3
4
,, , I
5
6
7
8
~
'
I
~
9
l
(
~
11
PH
Flgure 8. Relative effect of phosphate on general acid catalyzed monochloramine disDroDottionation. Ratio R of DhoSDhate to hvdroc-n ion contributions to k&. R = k,/k- where k,, = kw,[H3P04] khpo4-[H,P0,-] and kHy= k,+[H 1.
. .
I
JP"
+
overall kinetics of chloramine decomposition could be dramatically affected by general acid catalysis, it is also quite possible that other aspects of chloramine chemistry could be influenced. Registry No. C1NH2, 10599-90-3;C12NH, 3400-09-7.
5
7
6
9
8
10
PH
Figure 6. Relative effect of carbonate on general acid catalyzed monochloramlne disproportlonatlon. Ratio R of carbonate to hydrogen lon contributions to k,. R = k,,/kw where kcm = kW[H2CO3] k,,-[HCO,-] and kHy= kH+[H+].Shaded area shows probable zone 03 importance in typical drinking waters.
+
6
7
a
9
10
PH
Figure 7. Relative effect of silicate on general acid catalyzed monochloramlne disproportionation. Ratio R of silicate to hydrogen ion contributions to /rob&. R = ksl/kHYwhere k,, = kH,s104[H4Si041 kH,sio4-[H3Si0,-] and k , = kp[H+]. Shaded area shows probable zone of importance in typical drinking waters.
+
monochloramine disproportionation can under some conditions control overall oxidant loss, the effect of inorganics (and possibly some organic species) should be included in models used to predict chloramine speciation and fate. Catalysis could have a bearifig on the evaluation of the efficacy of chloramination disinfection, particularly the maintenance of chloramine residuals in distribution systems where a long contact time is available. Evaluation of laboratory experiments using buffers to control the pH in chloramine solutions, particularly phosphate, must be done carefully with appropriate consideration of possible effects. While it is clear that the 696
Environ. Scl. Technol., Vol. 22, No. 6,1988
Literature Cited (1) Leao, S. L.; Selleck, R. E. In Water Chlorination, Environmental Impact and Health Effects;Jolley, R. Lo,Brungs, W. A., Cortruvo, J. A., Cumming, R. B., Jacobs, V. A., Eds.; Ann Arbor Science: Ann Arbor, MI, 1983; Chapter 9. (2) Leao, S. L. Ph.D. Dissertation, University of California, Berkeley, CA, 1981. (3) Wei, I. W.; Morris, J. C. In Chemistry of Water Supply Treatment and Distribution; Rubin, A. J., Ed.; Ann Arbor Science: Ann Arbor, MI, 1974; pp 297-332. (4) Wei, I. W. Ph.D. Dissertation, Harvard University, Cambridge, MA, 1972. (5) Pressley, T. A,; Dolloff, F. B.; Roan, S. G. Environ. Sci. Technol. 1972,6, 622-628. (6) Saunier, B. M.; Selleck, R. E. J.-Am. Water Works Assoc. 1979,30, 164-169. (7) Saunier, B. M. Ph.D. Dissertation,University of California, Berkeley, CA, 1976. (8) Granstrom, L. M. Ph.D. Dissertation, Harvard University, Cambridge, MA, 1954. (9) Morris, J. C. In Principles and Applications of Water Chemistry; Faust, S. D., Hunter, J. V., Eds.; Wiley: New York, 1967; pp 23-53. (10) Gray, E. T.; Margerum, D. W.; Huffman, R. P. In Organometals and Organometaloids; Brinkman, F. F., Bellama,J. M., Eds.; American Chemical Society: Washington, DC, 1978; pp 264-277. (11) Standard Methods for the Examination of Waters and Wastewaters, 15th ed.; APHA-AWWA-WPCF: Denver, CO, 1980; pp 280-282. (12) Hand, V. C.; Margerum, D. W. Inorg. Chem. 1983, 22, 1449-1456. (13) Stumm, W.; Morgan, J. J. Aquatic Chemistry, 2nd ed.; Wiley: New York, 1981; p 135. (14) Snoeyink, V. L.; Jenkins, D. Water Chemistry;Wiley: New York, 1980; pp 446-447. (15) Moore, J. W.; Pearson, R. G. Kinetics and Mechanisms, 3rd ed.; Wiley: New York, 1981; pp 66-74. (16) Margerum, D. W.; Gray, E. T.; Huffman, R. P. In Organometals and Organometaloids; Brinkman, F. F., Bellama, J. M., Eds.; American Chemical Society: Washington, DC, 1978; pp 279-291. (17) Moore, J. W.; Pearson, R. G. Kinetics and Mechanisms, 3rd ed.; Wiley: New York, 1981; p 354. Received for review July 7, 1986. Revised manuscript received July 15, 1987. Accepted January 11, 1988.
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