Environ. Sci. Techno/. 1995, 29, 1127-1 134
Nonmetal Redox Kinetics: Mono-, Dim,and Trichloramine Reachns with Cyanide Ion Department of Chemistry, Purdue University, West Lafayette, Indiana 47907
Trichloramine reacts with excess CN- in a stepwise series of CI+ transfer reactions to give ClCN and to generate in sequence dichloramine, monochloramine, and ammonia: NC13 CN- H20 NHC12 ClCN OH-; NHC12 CNH20 NH2CI ClCN OH-; and NH2CI CN- H20 NH3 ClCN OH-. The reaction rates are first order in the corresponding concentrations of the chloramine species and in cyanide ion, with second-order rate constants (M-l s-l, 25.0 "C, p = 1.00 M) of 1.83 x lo9 for NC13, 576 for NHC12, and 1.96 x 10-2for NH2CI. The NH2CI reaction is general-acid assisted with third-order rate constants (M-2 s-l) of 4.32 x 1Olo for H3O+ and 84 for HCOs-. The reactions of NHC12 and of NC13 with CN- are not acid assisted. Since HCN ( pKa 8.95) is not reactive, the rates of the NHC12 and NC13 reactions decrease with decrease of pH (below pH 9). Off-setting effects of acid assistance and HCN formation cause the NH2CI rate to increase from pH 13 to pH 9 and to level off below pH 9. At pH 7, the relative reactivities are NC13 > HOC1 >> NH2CI >> NHC12.
+ + + + +
+
-
--t
-+ +
Hypochlorous acid has been shown (1) to react extremely rapidlywith cyanide ion to form cyanogen chloride (eq 1). HOC1
+ CN-
-
OH-
+ ClCN
(1)
The second-order rate constant for this C1+transfer reaction is 1.22 x lo9M-l s-l at 25.0 "C andp = 1.0 M, whichis only a factor of 6 smaller than the diffusion-controlled rate constant (7 x lo9 M-' s-l) (2). The hydrolysis of ClCN to give cyanate ion (eq 2) is a much slower process (3) that is base assisted (eq 3). The rate constant for the OH- path is k$& = 4.53 M-' s-l and for the H20 path is ~ C I C N= 2.58 x 10-6 s-1.
L Y N N M . SCHURTER, PAULA P . B A C H E L O R , A N D DALE W. MARGERUM*
+
Introduction
+ + +
ClCN
+ H,O - OCN- + C1- + 2H+
(2)
In the present work, we show that chloramines also react with cyanide ion via C1+transfer steps. When trichloramine is mixed with excess CN-, a stepwise series of reactions occur (eqs 4-6) that sequentially generate dichloramine, monochloramine, and ammonia. Cyanogen chloride (ClCN) is a proposed product for each of the stepwise reactions.
+ CN- + H,O -NHC1, + ClCN + OHNHC1, + CN- + H,O - NH,Cl+ ClCN + OHkl NH,Cl+ CN- + H,O - NH, + ClCN + OHNC1,
k3
k2
(4)
(5) (6)
Cyanogen chloride is an undesirable disinfection byproduct that can form in chloraminated waters and is included on the Environmental ProtectionAgency Drinking Water Priority List as a candidate for regulation (4). The threshold toxicity for aqueous ClCN with rainbow trout was calculated to be 0.08 mglL (5). Several mechanisms for ClCN formation have been proposed for the reaction of chloramines with aromatic compounds, amino acids, and peptides (6, 7),but little is known about the actual sources of trace ClCN found in water treatment. The use of chloramines as disinfectants appears to enhance ClCN formation (8,9). Knowledge of the kinetics and mechanisms of the reactions between CN- and chloramines may help to understand pathways for the formation of ClCN in water treatment processes. The NC13 reaction with CN- is extremely rapid, but the rate is suppressed in acid due to HCN formation. Thus, the rate can be measured by stopped-flow spectrophotometric methods at low pH. The present work shows that NC13 is even more reactive with CN- than is HOC1. The subsequent reactions of NHC12and of NH2C1with CN- are much slower, but the monochloramine reaction is acidcatalyzed. This work permits a comparison of the reactivities of hypochlorite ion and chloramines with CN- and with other nucleophiles and proves that ClCN is the product of the direct reaction between NHzC1 and CN-.
Experimental Section Reagents. Saturated solutions of NaOH were prepared from analytical-gradepellets. Aliquots were diluted in C02-free
0013-936Xf35/0923-1127$09.00/0
0 1995 American Chemical Society
VOL. 29, NO. 4, 1995 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
1
1127
TABLE 1
Ultraviolet Absorption Spectral Characteristics" species HOC1 OCINHzCI NHC12 NC13 ONOO-
1,nm 228 292 243 294 310 360 336 360 302 243
(max) (max) (max) (max) (max) (max)
c,
M-' em-' 120 350 461 272 195 -3 190 126 1670 770
ref
b b C
17
d 17 11 11 23 23
a The following species have negligible UV absorbance: CN-, HCN, and CICN. bGray, E. T., Jr. Ph.D. Thesis, Purdue University, West Lafayette, IN, 1977. CKurnar, K.; Day, R. A.; Margerurn, D. W. Inorg. Chem. 1986, 25, 4344-4350. This work.
water and standardized against potassium hydrogen phthalate. Hypochlorite solutions were prepared by bubbling Clz(g) through NaOH solutions of known concentrations, and the OC1- content was determined spectrophotometrically at 292 nm (Table 1). Ammonia solutions were prepared by dilution of concentrated ammonia and standardized byHC1 titration. Carbonate buffer solutions were prepared by dissolving NazC03 and NaHC03 in water. Phosphate buffer solutions were prepared by dissolving mixtures of Na3P04, Na2HP04,or NaHzP04 in COz-free water. Acetate buffers were made by diluting the ap.. propriate volumes of COz-free solutions of NaC2H302and HCZH3O2. Stock solutions of CN- were prepared by dissolving analytical-grade NaCN in either NaOH or in the buffer solutions of known concentration. The CN- solutions were standardiged within 24 h by potentiometric titration against AgN03 solution with a CN- sensitive electrode. The ionic strength of all solutions was adjusted with NaC104 prepared from the recrystallized solid and standardized gravimetrically. Preparation of Chloramines. Monochloramine solutions were freshlyprepared by combining solutions of OC1with 10% excess NH3 solution through a T-mixer (double T with a flow rate of 60 mL/min) and were standardized spectrophotometrically. Trichloramine solutions were prepared by combining a 3-fold excess of HOCl (pH 7 )with an NH4+solution (pH 3-4) in a T-mixer. This mixture was allowed to age for 18-24 h in the dark to permit side reactions to go to completion. Acidic NC13 solutions are relatively stable but are volatile [thevapor pressure of pure NC13 is 150 mmHg at room temperature (10)).Therefore, NC13 solutions were stored in stoppered flasks with no headspace. Solutions of NC13 were standardized spectrophotometrically at 336 nm based on its newly determined (11)molarabsorptivityvalue of 190M-I cm-'. The amount of NC13 formed was always less than the initial NH4+ concentration due to the formation of Nz (12,13). Typical yields of NC13 were 50-90% of the initial NH4+,and the solutions had a HOCl content that was I-10% of the NC13 concentration. The amount of HOCl present in the NC13 solutions was determined by a kinetic analysis method based on the reaction of the solutions with Br- on a Durrum stopped-flow spectrophotometer. The Br- reaction with HOCl (14)is more than 100times faster than the Br- reaction with NC13 ( 1 1 ) . Although both reactions yield Brz/Br3-, the initial absorbance increase at 410 nm can be used to calculate the HOCl content of the NC13 solution. 1128 m ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 4, 1995
Dichloramine solutions were generated in situ by the reaction of NC13with excess CN- (eq 4). The NHClz yields were 100% of initial NC13 concentrations for all reaction conditions employed in this study. Instrumental Methods. Slower reactions were followed by the use of either a Perkin-Elmer Model 320 spectrophotometer interfaced to a Perkin-Elmer 3600 data station or a Lambda 9 UV-visible near-IR spectrophotometer interfaced to a Zenith 386/20 computer. When reactions were run under acid conditions, the cyanide solutionswere kept basic and in gas-tight syringes until T-mixed into the spectrophotometric cell in a hood. The cell was then tightly capped before being inserted into the instrument. A Durrum Model 110 stopped-flow spectrophotometer interfaced to a Zenith 151 PC with a MetraByte Dash-16 AID converter was used for faster reactions. Typically, six runs were obtained for each set of conditions, and pseudo-firstorder rate constants were measured from plots of ln(A-AJ versus time. All pH measurements were taken with an Orion Model SA 720 digital pH meter equipped with a Corning combination electrode. For stopped-flow reactions, the pH measurements were taken after the reactions were complete. For slow reactions, initial pH measurements were taken. All solutions were well buffered. Hydrogen ion concentrations were calculated from pH calibration data, based on electrode response to titrations of standardized NaOH with HC104 at an ionic strength of 1.00 M NaC104. We use p[H+]= -log[H+j, because [H+jvalues are expressed in molarities rather than in activities (pH = log U H ) . It is essential to use the same units of concentration for all species in rate and equilibrium constants. Membrane introduction mass spectrophotometry (MIMS) (1.5)was used to show that ClCN is the product of the reaction between NHzCI and CN-. Mass spectra were obtained with a Finnegan TSQ 4500 triple quadrupole instrument equipped with an INCOS data system. One quadrupole was used for mass analysis of the ions produced in the ion source, while the other two quadrupoles were operated in the RF mode only. Multiple ion detection was also employed. Ions were produced under 70 eV electron impact ionization conditions. The source temperature was maintained at 190 "C while the manifold temperature was set at 106 "C. A flat sheet of dimethylvinyl silicone polymer membrane with a thickness of 0.005 in. was mounted in a specidyconstructed direct insertion probe (16). The probe was inserted into the ion source of the mass spectrometer and made contact with the source block in such a way that the membrane formed one of the walls of the heated ion source of the mass spectrometer. The membrane probe was internally heated to 30 "C by a built-in heater operated by a programmable temperature controller. A peristaltic pump was used to transport the solution through the probe and across the membrane. Flow rate was maintained at 1 mL/min. Initial signal detection was typically 16 s after introduction.
Results and Discussion NC13 and CN-. The reaction between trichloramine and cyanide ion is extremely fast. This is not unexpected, because rate constants for the NC13 reactions with iodide ion (17) and with sulfite ion (18)are also verylarge (9 x lo7 and 4.5 x lo9 M-' s-l, respectively). The cyanide ion rate can be decreased to permit stopped-flow measurements by lowering the pH of the reactant solutions to give a low
16
TABLE 2
Observed Second=Order Rate Constants for Reaction between NCl3 and [CN]f lHNCl& (M)
I H C N ~ ,(M) O
10YH+l (MI
2.54 2.54 2.83 2.83 2.83 2.83
3.15 5.12 6.09 5.12 5.12 5.12
1.38 1.38 1.38 1.38 2.45 4.37
(M-' s-') 1.5 f 0.1 1.44 f 0.08 1.49 f 0.09 1.54 f 0.09 0.80 f 0.02 0.46 f 0.01
-
v1
#
i 3
a Conditions: IOAclr = 0.050 M , p = 1.0 M (NaCIO,), 25.0 f 0.1 "C, stopped-flow observations at 360 nrn.
[CN-l:[HCNl ratio. The value of pKaHCN is 8.95 at 25.0 "C andp= 1.OM (19). Asasafetyprecaution, gas-tightsyringes were used to store the low pH cyanide solutions. The disappearance of NC13 was monitored at 360 nm rather than at its absorption maximum (11,131of 336 nm in order to minimize any absorbance contribution from NHCL (Table 1). An acetatelacetic acid buffer solution (0.05M) was used to control the pH and to ensure that protontransfer steps did not limit the rate. The reactions of interest are given in eqs 7-9, and a steady-state approximation in [CN-I gives the rate expression in eq 10 where A- is the acetate ion. The initial NC13concentrationswere kept below
+ H,O
+
+ H,O+
12
8
Iyg
vo-
4
0 0
2
4
6
8
l/[H+], M ' FIGURE 1. Inverse [H+l dependence of the second-order rate constants for the reaction of NCIJwith CN- ion. The line is a linear regression with the origin included in !he data set, where the slope = LIGncn. DuFrum stopped-flow spectrophotometer, A = 360 nm, B = 1.85 cm, [OAch = 0.050 M, p = 1.00 M (NaClO,), 25.0 f 0.1 "C.
Table 2. A plot of kobsd against [H+I-' with these three points gives a line with an intercept statisticallyequivalent to zero, so the origin is also used as a data point. The plot HCN CH3COOCN- CH,COOH (8) in Figure 1 yields a slope of 2.05 k 0.07 s-l, which corresponds to a k3 value of (1.83 f 0.04) x lo9 M-' s-l. Thus, variation in both [CNITand [H+]give the same value k3 NCl, CN- H,O products (9) for k3 within experimental error; the average value for k3 is 1.83 x lo9 M-' s-l. NHClz and CN- in Basic Solution. Dichloramine, generated in situ from the rapid reaction of NC13with CN(eq 41,reacts more slowly with CN- and generates NH2C1 (eq 5). In basic solution, the reaction between NH2C1and 0.3 mM in order to maintain the inequality: k ~ [ H + l + kmCN- is much slower than the reaction between NHC12 and [HA] >+- k3[NCl31. The proton-transfer rate constant ( k d CN-, so its rate can be neglected when measuring the NHClz for the reaction between CN- and CH3COOH, where ApK, reaction. The reactions are monitored by stopped-flow = pK,(HCN) - pKa(CH3COOH)is 4.45,can be calculated (20) to be between lo9 and 1Olo M-' s-l. Since the [HA]/ methods at 310 nm to avoid NHF1 absorbance (13).With M, excellent firstexcess CN- and [OH-] = 2.51 x [NC13]ratio is always greater than 60 and k3 must be less order plots for eq 13 are obtainedwith standard deviations than the diffusion-controlledvalue of 7 x lo9 M-' s-l, the inequality is valid. It can be shown that under these conditions CN- and HCN are in equilibrium, so that eq 10 (13) can be simplified to eq 11, where KaHCN = b / k H , [CNIT=
HCN
CN-
&
+
+
+
(7)
+
-
d [NCl,] - k 3 c C N [CN] [NCl,] dt [H'l
(1 1)
+
[HCNI [CN-I, and [HCNI >> [CN-I. The reactions were observed under second-order unequal concentrations of NC13 and KNIT with variation of KNIT and [H+l concentrations (Table 21,where the observed second-order rate constant is given by eq 12. An average of the first four sets of data at 1.38 x M [H+land variable ICNITgives kobsd = (1.49 & 0.04) x lo5 M-' s-l. This corresponds to a k3 value of (1.84& 0.05) x lo9 M-' s-l. kobsd = k 3 c C N /[H']
(12)
for kobsd of fl-2%. The cyanide ion concentration is calculated from the total cyanide added less the amount that reacts rapidlywith the NC13 and with the HOC1 present in the NC13 (eq 14). The kobsd value increases with the
[CN-I, value as seen in Figure 2, where the slope is 576 f 8 M-' s-l (k2). There is a cyanide ion independent term (intercept = 0.8 f0.1s-l) that can be attributed to reaction with hydroxide ion (eq 15) where k20H = 320 M-l s-l. Previous studies (12)of NHClz decomposition reactions in NaOH gave similar kZoH values (400 & 100 M-' s-l).
A [H+]dependence can be plotted using the average of the first four data points and the last two data points in VOL. 29, NO. 4, 1995 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
1129
I
I
I
I
I
find [CNI,,,
= [CNI,
- 3[NC1,l0 - IHOCllNc,3 (17)
initial [CNI,, = [CNI, - [NCl,], - [HOCllNCl3 (18) [CNIT,, = [CNlo - 2[NCl,l, - [HOC11N,13
(19)
We used eq 19 to calculate the average [CNIT in Table 3 during the reactions. The [CN-I, concentration is calculated from eq 20. The rate expression is given by eq 13, and kobsd is defined by eq 21.
0 000
I
I
0.005
0.010
I
0.015
0.020
0 025
[CNI,, M FIGURE 2. Cyanide ion dependence of the observed first-order rate constants for the reaction with NHClz in basic sotution. Linear regression on the data yields a slope of k and an intercept of k ~ ~ T 0 H - Dumm l. stopped-flow spectrophotome#er, rZ = 310 nm, Ib = 1.85 em, [ N C I h = 4.55 x lo-' M,[OH-] = 2.51 x M, [HOCllo = 3.03 x loe5 M, p = 1.00 M (NaCIO,), 25.0 f 0.1 "C.
A carbonate buffer dependence was tested to determine if the reaction experiences general-acid catalysis. At p[H+l = 9.69, with [C03]T variation from 0.025 to 0.075 M and [CN-IT,m = 1.45 x 10-3M,theresultsshowedthat carbonate buffer concentration has no effect on the k2 rate constant within statistical error. The second-order rate constant determined for k2 is 585 f 11 M-l s-l, which is in close agreement with the cyanide ion dependence data in NaOH. Thus, there is no evidence of general-acid catalysis by HC03-. NHClz and CN- in Acid Soludon. Dichloramineis again generated in situ from the reaction between NC13and excess CN- by T-mixing an acidic NC13 solution with a basic CNsolution. Reactions are monitored with the Perkin-Elmer Lambda 9 spectrophotometer and 10-cm cells. Since most of the cyanide is in the form of HCN in the p[H+] range studied (4.13-5.331, the [CNIT concentration used was increased to 10 mM to increase the reaction rate. Nevertheless, the reactions were quite slow with half-lives that varied from 0.17 to 2.2 h. To ensure an adequate absorbance change for the reaction, the NC13 concentration (and thus the NHC12 concentration) was kept at 0.2 mM. All runs fit excellent first-order kinetics over a period of more than 4 half-lives. Belowp[H+]7, the acid-catalyzed reaction between NH2C1 and CN- becomes faster than the reaction between NHCl, and CN-. The overall stoichiometry of the reaction is given in eq 16. Rate constants were evaluated under pseudo-
NHC1,
+ 2CN- + 2H20 - NH, + 2ClCN + 20H(16)
first-order conditions with [CNIT:[NCLI ratios of 40-5O:l. However, small corrections are needed for the CNconsumed by the rapid reactions with NCI, and the HOCl in the NCl3 before the NHClz reaction, for the CNconsumed by NHC12, and for the NH2C1 generated by the NHCl2 reaction. Equations 17-19 give the relationships. 1130
ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 4, 1995
Data from the p [H+l dependence are given in Table 3; the average value of the second-order rate constants is 561 f30 M-' s-l. Table 3 also contains data to test for generalacid catalysis by varying the acetate buffer concentration. No buffer dependence was found, and the average k2 value was 571 & 25 M-l ssl. Both of these values agree (within their experimental precision) with the k2 value of 576 f 8 M-' s-l measured in high base. These data clearly show that there is no specific or general-acid catalysis for the reaction between NHCl, and CN-. NH&l and OH-. In high hydroxide ion concentrations, monochloramine undergoes a slow reversible hydrolysis (eqs 22 and 23) to give hypochlorite ion. An equilibrium NH,C1
+ H,O 2 NH, + HOCl
(22)
HOCl
+ OH- 2 OC1- + H,O
(23)
constant of 3.8 x 1OloM-l has been recently evaluated (21) for [NH2C11/([NH3][HOCl]) at 25.0 "C and y = 0.50 M. A rate constant of 2.8 x lo6 M-' s-l has been measured (22) for the formation of NH2Cl from NH3 and HOCl (25.0 "C, y = 0.10 MI. We can assume that ionic strength will not greatly change this rate constant for two neutral species, so that a NHzCl hydrolysis rate constant, kh = 7.4 x s-l, can be calculated. We observed a very slow loss of NHZCl at 243 nm when 1.0 x M "$21 in loW4M NH3was reacted with 0.496 M NaOH at p = 1.0 M. Under these conditions, 9% of the NH2Cl will hydrolyze to form OC1-. The initial rate is consistent with the calculated kh value, but it is followed by a much slower reaction due to the formation of hydroxylamine (eq 24) (23). In the presence of OC1- and
+ OH- -",OH + C1klOH
"$1
(24)
0 2 (air saturated solutions), the NH20H formed in eq 24 reacts further to give peroxynitrite ion (ONOO-) with an absorbance increase at 302 nm (24). The ONOO- ion also absorbs at 243 nm (Table 1) and after correction for its contribution at this wavelength,the observed rate constant for the slower loss of NH2CI is less than s-l. (We conclude that the Anbar and Yagil measurement (23) of the rate of alkaline hydrolysis of NH2Cl to give NH20H was too large due to contributions from eqs 22 and 23.) In the presence of CN-, any OC1- formed reacts rapidly (1)to give ClCN, so there is no interference from peroxynitrite ion
TABLE 3
Second=Order Rate Constants for Reaction of NHCl2 with CN- in Acidic Solutionsa p[H+I Variation at [OAch = 0.050 M
PWI [CNh,av (mM) 1O7[CN-1m(MI 1@kob*d(8-l)
k (M-' 8-0
4.16 4.16 4.31 4.32 4.49 4.49 4.91 4.93 5.28 5.28 5.33
579 628 566 593 549 564 529 543 525 563 529 561 f 3 0
9.09 9.08 9.10 9.06 9.08 9.10 9.03 9.04 9.17 9.15 9.06
1.47 1.47 2.08 2.12 3.15 3.15 8.24 8.63 19.60 19.55 21.72
0.0854 0.0924 0.118 0.126 0.173 0.178 0.436 0.469 1.03 1.10 1.15
av
[OAclr (MI
[OAch Variation at p[H+] 4.5 f 0.1 [CNh,av 10'[CN-lmn 1@kbid p[H+I (mM) (MI (8-l)
0.050 0.050 0.061 0.061 0.118 0.118 0.150
4.49 4.49 4.45 4.47 4.43 4.45 4.58
9.08 9.10 9.11 9.08 9.10 9.06 9.14
av
3.15 3.15 2.88 3.01 2.75 2.87 3.90
0.173 0.178 0.171 0.164 0.165 0.156 0.234
0.00 0
kz (M-l s-l)
549 564 594 545 601 544 600 571 f 25
6
12
10
24
l/[OH'], M" FlGURE3. Inverse hydroxide ion dependence ofthe maction between NHzCI end CN-. The data fit the expression in eq 28 to determine k1 and klH.Perkin-Elmer Lambda 9 spectrophotometer, E, = 243 nm, b = 5.00 crn, [CNIT,~~ = 0.038 M, p = 1.00 M (NaC104),25.0 f 0.1 "C.
NH,Cl+ CN-
Conditions: 25.0 "C,p = 1.0 M, [NCI~IO =Z 0.2 mM, [CNIo PZ 9.5 mM, 310 nrn, 10-cm cells. a
kobsd
klH + H30+NH, + ClCN + H,O
= kh -k (k1 -k k1H[H30+l)[CN-l
(26)
(27)
TABLE 4
(28)
Inverse Hydroxide Ion Dependence for Reaction between NHzCl and CN- a lDZ[OH-] (M)
10(kobrd(8-l)
4.91 9.93 14.8 19.8 25.3 29.8 34.2 38.6
17.2f 0.1 12.4f 0.1 11.1 f 0.1 10.7f 0.1 9.93f 0.2 9.57f 0.2 9.53 f 0.1 9.43f 0.10
1Wkobrd - kh)/[CN-] (M-'
8-l)
4.33f 0.03 3.07 f 0.03 2.73 f 0.03 2.62 f 0.03 2.42f 0.05 2.32 f 0.05 2.31 f 0.03 2.29 f 0.03
= Conditions: [CNIT,~" = 0.038 M, [NHzCI] = 2.41 x lo-' M, 2.4 x M; 5.00 cm cells, A = 243 nm,y = 1.0 M, 25.0 f 0.1 O C .
formation. At lower OH- concentrations, the hydroxylamine pathway for loss of NH2C1 in eq 24 can be neglected relative to the hydrolysis pathway in eq 22. NHzCl and CN-, Inverse Hydroxide Ion Dependence. The reaction between NHzC1 and CN- in basic solution is much slower than the NC13 or NHCL reactions with CN-. The disappearance of NHzC1 is monitored at its Amax of 243 nm (Table 1). The loss of NH2Cl is first-order over 4 or more half-lives (eq 25). The reaction between (25) monochloramine and cyanide ion (eq 6) has an inverse hydroxide ion dependence (Table 4) that can be attributed to contributions from an acid-assisted pathway (eq 26). The observed rate constant for the loss of NHzCl from eqs 6, 22, and 26 is given by eq 27. Rearrangement of this equation in terms of [OH-] (eq 28) and the use of the
calculated kh value of 7.4 x s-l permits evaluation of kl and klH. Data plotted in Figure 3 at constant [CN-1 = 0.038 M and variable [OH-] give klH = (4.6 f 0.1) x 1Olo M-, s-l, based on K, = 10-13.60at 25.0 "C and p = 1.0 M (23, and kl = (2 & 2) x lo-, M-l s-l. If the value for kh is not assigned, it can be determined from a nonlinear regression analysis program (Sigma Plot) (267,which uses a Marquardt-Levenberg algorithm to solve for kh as well as kl and klH. This gives kh = 6.6 x s-l, but the uncertainty is 10 times the value. "$1 and CN-, Cyanide Ion Dependence. Rate constants were measured with variation of [CN-IT concentrations from 0.627 x lo-, to 5.40 x lo-, M at three different concentrations of NaOH (from 0.030 to 0.468 M). These combined data were fit to the kobsd expression in eq 29 and were analyzed by nonlinear regression (26). Values
obtained are as follows: kh = (7.4 f 1.4) x 10-5 s-1, kl = (1.96 0.04) x lo-, M-' s-l, and klH= (4.32 f 0.07) x 1Olo M-, s-l. Figure 4 shows a plot of the data at each OHconcentration. The results of this analysis give a value for kh that agrees with our prediction and our estimate from the hydroxide dependence data. The kl and klH values are in agreement with the results obtained from the [OH-] dependence data at constant [CN-] concentration, but are determined with greater precision. Hydrogen Carbonate Dependence. In order to test for general-acid catalysis as shown in eq 30, a buffer depen-
*
VOL. 29, NO. 4, 1995 /ENVIRONMENTAL SCIENCE &TECHNOLOGY
1131
0.0015
9 TABLE 5
Carbonate B a r Dependence Data for Reaction between NHzCl and CN- a [COdT.nn (MIb
p[H+I
0.050 0.100 0.150 0.200 0.250 0.300 0.350 0.400
9.76 9.76 9.76 9.76 9.76 9.77 9.77 9.77
lgkobsd kl) 4.05 4.93 5.68 6.30 6.88 7.47 8.07 8.52
a Conditions: [ C N I T ,= ~ 5.0 ~ ~ rnM, INH~CII,,, = 0.478 rnM, INH&n = 0.05 m M ; p = 1.0 M, 25.0 f 0.1 "C, stopped-flow observations at 243 nrn. b T o calculate [HCO3-1,the pK, = 9.57, y = 1.0 M, 25.0 i 0.1 "C is used. Frydrnan, N.; Nilsson, G.; Rengerno, T.; Sillen, L. G. Acta Chern. Scand. 1958, 12, 878-884.
0 0000 0 00
0 02
0 04
0 06 I
I
I
I
[CN.], M"
i/
FIGURE4. Cyanide ion dependence in base for the reaction between NHzCIand CN-. Perkin-Elmer Lambda 9 spectrophotometer, 1 = 243 M, [NHs] = 0.1 [ N H & , nm, b = 5.00 cm, [NHSlb = (4.1-4.9) x p = 1.00 M (NaC103.25.0 f 0.1 "C. (0)[OH-] = 0.030 M, (A) [OH-] = 0.100 M, and (v)[NaOHl = 0.468 M.
NH,Cl
?
+ CN- + HA -NH, + ClCN + A-
I
(30)
dence was conducted over a concentration range of [ C 0 3 1 ~ = [HC03-] [C032-l = 0.050-0.400 M with constant ionic strength maintained at 1.0 M. The pseudo-first-order data for the carbonate buffer dependence are given in Table 5 and are plotted in Figure 5. The data follow the relationship in eq 31 with the slope equal to klm. The kl term is small but not negligible under these reaction conditions. Regression on the data yields a slope of klHA= (84 & 2) M-2 s-l for HC03-, which clearly shows general acid assistance by HC03-.
+
Brsnsted-Pedersen Relationship. The acid dissociation equilibrium constants (K,) and the third-order rate constants (klHA)for the general acid-assisted reactions and CN- are correlated by the Bronstedbetween "$21 Pedersen relationship (27)in eq 32, where Gais the Bronsted
r:)
=logGa+ alog -
(32)
proportionality constant. Table 6 summarizes the Ka and kl*values for HA equal to H20,HCO3-, and H3O+ and the assigned values for p (number of equivalent acid sites on the acid form of the species) and 9 (number of equivalent basic sites on the conjugate base). An a value of 0.83 f 0.04 is determined from the slope of log ( k l H A l pvs ) log (Kaqlp),as shown in Figure 6. This a value indicates that there is a high degree of proton transfer in the transition state as shown in Scheme 1. It should be noted that preequilibration of NH2Cl with general acids to form NH3Cl+, pK, 1.44 (221,followed by a rate-determining "$3 reaction with CN- is not an acceptable mechanism. Such a preequilibration step would make the process a specific 1132
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 29,
NO. 4,
1995
0 0 00
0.03
0 12
0.09
006
0.15
0 18
[HCO,'], M FIGURE 5. Determinationof HCOs- catalysis of the reaction between NHzCl and CN- ion. Linear regression was performed on the data using eq 31, which was rearranged to give an intercept of zero and Durrum stopped-flow spectropkotometer. 1= 243 a slope of knA. nm, b= 1.85 cm, [COJ~-I= 2[HC03-],p = 1.00 M (NaClO4 and NaOH), 25.0 f 0.1 "C. TABLE 6
Thinl=Order Rate Constants for General Acid=Assisted Reactions with NH$I and CM- Used for Brensted-Pedersen Determination of a Valuea species
PKa
P
9
klHAiWZs-l)
H20 H30' HCOj-
15.34b -1.72
2 3
3 2
3.53 1 0 - 4 c 4.32 x 1O'O
9.57*
1
3
84
a Conditions: p = 1.0 M, 25.0 f 0.1 "C, pK,= 13.60. -log (KJ55.5). kd55.5. Frydrnan, N.; Nilsson, G.; Rengemo, T.; Sillen, L. G. Acta Chern. Scand. 1958, 12, 878-884.
acid-assisted mechanism (dependent on [H+I only) as opposed to the observed general acid-assisted mechanism, which is dependent on the concentrations of allacids (H20, HCOs-, H30+). The proton transfer must take place in the transition state in order to have general-acid participation. MIMS Analysis of the Reaction Products. In eqs 4-6, ClCN is proposed as a major product of the reactions. To
I
-16
-12
I
I
I
-8
-4
0
/
1% (K,dP) FIGURE 6. Brensted-Pedersen plot (eq 32) for general acid assistance of the reaction between "$1 and CN-, where the slope = a = 0.83 f 0.04. SCHEME 1 AH
+ NH&l + CNr
50
70
60
80
dz FIGURE 7. MIMS spectrum of the reaction product of NHtCl and an equivalent amount of NaCN. [NHtCllo = 3 mY, p[H+l,. = 10; after quenching with HClO, for MIMS,the p[H+] = 5.5. 10
-
I
I
I
/
I
.-'-'-'-'-'
NCI,
8 U
. )
1
40
4
H'
A-
+
NH3
+ ClCN
6
I
prove the existence of this reaction product, a 3 mM solution of NHzC1 at p[H+l 10 was reacted with 1equivalent of solid NaCN. The reaction solution was quenched with HCIOl to give p [H+l5.5in order to decrease the rate ofbase hydrolysis of ClCN. The quenched solution was passed over the membrane of the MIMS insertion probe while the instrument scanned over the range of 45- 150 m u . The spectrum shown in Figure 7 exhibits peaks at mlz ratios of 61 and 63. These peaks correspond to 35C1CN and 37C1CN,in the proper isotopic ratio. All of the NH2Cl has been consumed in the reaction, as no peaks appear at mlz values of 51 (NHz35C1+)or 53 (NH237C1+). Mechanistic Comparisons. The reactions of CN- with the chloramines follow the sequential reaction path given in eqs 4-6. NC13and NHC12 show no acid catalysis,whereas NH2C1exhibits both specific and general acid catalysis. Thus, proton transfer does not occur in the NC13 and NHClz transition states. Only after the N-Cl bond is broken and ClCN forms does a proton add to the nitrogen. Scheme 2 depicts the mechanism for the reaction between NC13 and CN-. The reaction between NHC12 and CN- follows an analogous mechanism, where C1+ transfer from the chloramine to the CN- occurs in the transition state. A plot of predicted second-order rate constants for NC13, NHC12, "$1, and OC1- vs p[H+]is shown in Figure 8. The decrease in rate constants with decreasing p[H+l for the reactions of NC13, NHC12, and HOCl is due to protonation of CN- to form the unreactive HCN. The reaction rate for OC1- also drops at high p[H+],because OC1- is less reactive than HOCl (pK, = 7.47) (I)with CN-. The "$21 reaction rate increases as the p[H+]decreases from 12 to 9 because the reaction is acid-catalyzed. As the p[H+]decreases below
.$ 00
-
4
\
0
-------
2
:
, . . ' ........................ NHCI,
\
\
\ 0
\
\
NH,CI \
-2 10
8
6
4
12
14
PF+I FIGURE 8. Plot of the predicted second-order rate constants for NCls NHCh NHzCI, and HOCVOCI- reactions with [CNh ion as a function of p[H+I. These rate constants are based on experimental data with NCl3 at p[H+l3.6-3.9, with NHCIt at p[H+14.2-11.0, with NHtCl at p[H+l 9.8-13.3, and with OCI- at p[H+l 11.8-13.3. SCHEME 2
NC13
+ CN-
-
CI,
,.. (-)
'1
N- -
$'
CI NHClp
+ ClCN + OH-
pH 8, the NHZCl reaction rate levels off because of counterbalancing effects of specific acid catalysis and protonation of CN-. Rate constants for the reactions of NC13, NHC12,"$3, HOC1, and OC1- with CN-, S 0 3 2 - , I-, and Br- are compared VOL. 29, NO. 4, 1995 I ENVIRONMENTAL SCIENCE &TECHNOLOGY
1133
TABLE 7
Acknowledgments
Rate Constants for Reactions between Active Chlorine Electrophiles and Various Nucleophiles (CI+ Transfer Processes)a
This work was supported by National Science Foundation Grant CHE-9024291. We wish to acknowledge preliminary studies by Jenna T. Nguyen and a literature review of ClCN by Erik Pederson 111. We are grateful to Lindy E. Dejarme and R. Graham Cooks for their assistance in the MIMS experiments.
rate constants’ for nucleophiles (n) electrophiles SO& (5.10)
OCIHOC1 NHpCl H+ “$21 NHClp H+ NHC12 NC13
+ +
CN- (5.10)
2.3 x IO4 3.1 x IO2 7.6 x I O 8 * 1.2 x IOsc 7.7‘ 1.96 x 8.0 x 1O’O 4.3 x 10’0 5.8 x IO6 I 5.76 x IO2 none none 4.5 1091 1.8 1098
I- (5.04)
Br- (3.89)
not detn 1.4x I O E e not detn 2.4 x 1Olo e
0.9 x 1.55 x IO3“ not detn 8.9 x IO4 not detn not detn 12~
0.56e
9.3 x IO5 e 9 107)
“The nucleophilicity value ( n ) of each nucleophile is given in parentheses. M - l s - ’ for second-order reactions; M-Z s-’ for H+-assisted third-order reactions. * Fogelman, K. D.; Walker, D. M.; Margerum, D. W. Inorg. Cbem. 1989,28,986-993. Ref 1. Ref 13.e Kumar, K.; Day, R. A.; Margerum, D. W. Inorg. Cbem. 1986,25,4344-4350.’Yiin, B. S.; Walker, D. M.; Margerum, D. W. Inorg. Cbem. 1987, 26,3435-3441. gThis work. *Gazda, M.; Margerum, D. W. To be submitted for publication. ’ Ref 17. Ref 16. Ref 1 1 .
in Table 7. All of these reactions proceed by C1+ transfer mechanisms. The magnitude of the rate constants are very sensitive to the nucleophilicity (n) (28) values for S032-, CN-, I-, and Br-, as shown previously for the HOCl reactions. The nucleophilicities of s03’- and CN- are the same, and I- is only slightly less, but Br- is a much weaker nucleophile. This is reflected in much smaller rate constants for Br- with OC1-, HOC1, and NC13. The huge effect of nucleophilicity on the reaction rates can be seen for the NC13 reactions with SO3*- and with Br-, where the ratio of the rate constants is 3.8 x lo8. None of the NC13 reactions are acid assisted, whereas all of the nucleophiles have acid-assisted pathways with OC1and NH2Cl. Dichloramine shows no acid assistance in its reactions with S032- or CN-, but it displays an imporant acid-assisted path in its reactions with I-. At pH 7, NHCl2 is the least reactive electrophile with CN- and I-, whereas NH2C1is the least reactive with S032-. The ratio of rate constants for S032- vs CN- and for CN- vs I- vary from 0.6 to 20 for the various electrophileswith the notable exception of NHC12. In this case, the ratios are 1 x lo4for S032-/CNand 1 x lo3 for CN-/I-. It is not clear why the reactivity of NHCl2 is more sensitive to the nature of these nucleophiles. At pH 7, the relative reactivities with [CNIrare NC13 > HOCl >> NHZCl>> NHC12. The CN-/HCN reactions appear to be quantitative with allthese active reactive chlorine species. The rate constants shown in Figure 8 indicate that the composition of mixtures of hypochlorite and chloramines could be determined by kinetic analysis. For example, at pH 10, the relative reactivitiesare NCl3 >> OC1- >> NHCl2 >> NH2Cl. Adjustment of pH, level of CN- concentration, and the wavelength for spectrophotometric observation could be used for this type of kinetic analysis.
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ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 4,1995
literature Cited (1) Gerritsen, C. M.; Margerum, D. W. Inorg. Chem. 1990,29,27572762. (2) Caldin, E. F. Fast Reactions in Solution; Wiley: London, 1964; pp 10-12. (3) Bailey, P. L.; Bishop, E. J. Chem. SOC., Dalton Trans. 1973, 9, 912-916. (4) (a) US. EPA.Fed. Regist. 1988,53,1892-1902. (b) U.S. EPA.Fed. Regist. 1991,56,1470-1474. (c) U.S. EPA.Fed. Regist. 1994,59, 38471. (5) Allen, L.A,;Blezard, N.;Wheatland, A. B.J. Hyg, 1948,46,184193. (6) (a) Ohya, T.; Kanno, S. Chem. Plzarm. Bull. 1988,36,4095-4102. (b) Ohya, T.; Kanno, S. Chemosphere 1989,19,1835-1842. (7) Hirose, Y.; Maeda, N.; Ohya, T.; Nojima, K.; Kanno, S. Chemosphere 1988,17,865-873. (8) Krasner, S. W.; McGuire, M. J.; Jacangelo, J. G.; Patania, N. L.; Reagan, K. M.; Aieta, E. M. J. Am. Water Works Assoc. 1989,80, 41-53. (9) Krasner, S.W.; Hwang, C. J.; Lier, T. K.; West, M. 1. Proc. Water Qual. Technol. ConJ 1991 (publ. 1992),1207-18. (10) TheMerckIndex, 9th ed.; Windholz, M. Ed.; Merck: Rahway, NJ, 1976;p 858. (11) Gazda, M.; Kumar, K.; Margerum, D. W. Inorg. Chem. Submitted for publication. (12) Hand, V. C.; Margerum, D. W. Inorg. Chem. 1983,22, 14491456.
(13) Yiin, B. S.;Margerum, D. W. Inorg. Chem. 1990,29,2135-2141. (14) Kumar, K.; Margerum, D. W. Inorg. Chem. 1987,26,2706-2711. (15) Kotiaho, T.; Lauritsen, F. R.; Choudhury, T. K.; Cooks, R. G.; Tsao, G. T. Anal. Chem. 1991,63,875A-883A. (16)Bier, M. E.; Kotiaho, T.; Cooks, R. G.Anal. Chim.Acta 1990,231, 175- 190. (17) Nagy, J. C.; Kumar, K.; Margerum, D. W. Inorg. Chem. 1988,27, 2773-2780. (18) Yiin, B. S.; Margerum, D. W. Inorg. Chem. 1990,29,1942-1948. (19)Giibeli, A. 0.; CBt6, P. A. Can. J. Chem. 1972,50, 1144-1148. (20) Eigen, M. Angew. Chem., Int. Ed. Engl. 1964,3, 1-72. (21) Margerum, D. W.; Schurter, L. M.; Hobson, J.; Moore, E. E. Environ. Sci. Technol. 1994,28, 331-337. (22)Margerum, D. W.; Gray, E. T., Jr.; Huffman, R. P. Organometals and Organometalloids: Occurrenceand Fate in theEnvironment; Brinckman, F . E., Bellama, J. M., Eds.;ACS Symposium Series 82;American Chemical Society: Washington, DC, 1978;p p 278291. (23) Anbar, M; Yagil, G. J. Am. Chem. SOC. 1962,84, 1790-1796. (24) Hughes, M. N.; Nicklin, H. G. J. Chem. SOC. (A] 1971,164-168. (25)Molina, M.; Melios, C.; Tognolli, J. 0.;Luchiari, L. C.; Jafelicci, M., Jr. I. Electroanal. Chem. Interfacial Electrochem. 1979,105, 237-246. (26) Jandel Scientific, P.O. Box 700S,San Rafael, CA 94912-7005. (27) Bell, R. P. The Proton in Chemistry, 2nd ed.; Cornell University Press: Ithaca, NY, 1973;p 198. (28) Swain, C. G.; Scott, C. B. 1.Am. Chem. SOC. 1953, 75,141-147.
Received for review October 10, 1994. Revised manuscript received December 20, 1994. Accepted December 21, 1994.@ ES9406290 @
Abstract published in Advance ACS Abstracts, February 1, 1995.
