Radiolytic Reactions of Monochloramine in Aqueous Solutions

Article abstract of DOI:10.1021/jp030198k

Monochloramine reacts with hydrated electrons very rapidly, k(NH ₂Cl+e_aq^-) = (2.2 ± 0.3) × 10 ¹⁰ L mol^-1 s^-1, to produce ^.NH ₂ radicals. It reacts with ^.OH radicals more slowly, k(NH₂Cl+^.OH) = (5.2 ± 0.6) × 10⁸ L mol^-1 s^-1, to produce ^.NHCl radicals. While ^.NH₂ exhibits an absorption peak at 530 nm, with a molar absorption coefficient ε₅₃₀ = 80 L mol^-1 cm ^-1, ^.NHCl exhibits two peaks at 330 and 580 nm, ε₃₃₀ = (85 ± 30) L mol^-1 cm^-1 and ε₅₈₀ = (56 ± 30) L mol^-1 cm^-1. The ^.NHCl radical undergoes self-decay and can react also with O ₂ to form a peroxyl radical. It is suggested that the peroxyl radical exists in equilibrium NHClO₂? NHCl + O₂ with an estimated equilibrium constant of (3 ± 2) × 10 ^-3 mol L^-1. The reaction of chloramine with the carbonate radical is suggested to form a complex [CO₃NH₂Cl] ^.- with k_f = 2.5 × 105 L mol^-1 s ^-1 and k_r = 4 × 10² s^-1, and this complex decomposes with k = 7 × 10² s^-1 to form ^.NHCl.

Full text of DOI:10.1021/jp030198k

J. Phys. Chem. A 2003, 107, 7423-7428
7423
Radiolytic Reactions of Monochloramine in Aqueous Solutions
G. A. Poskrebyshev, R. E. Huie,* and P. Neta*
Physical and Chemical Properties DiVision, National Institute of Standards and Technology,
Gaithersburg, Maryland 20899-8381
ReceiVed: February 12, 2003; In Final Form: May 5, 2003
Monochloramine reacts with hydrated electrons very rapidly, k(NH2Cl+eaq-) ) (2.2 ( 0.3) × 1010 L mol-1
s-1, to produce •NH2 radicals. It reacts with •OH radicals more slowly, k(NH2Cl+•OH) ) (5.2 ( 0.6) × 108
•
•
L mol-1 s-1, to produce NHCl radicals. While NH2 exhibits an absorption peak at 530 nm, with a molar
absorption coefficient ꢀ530 ) 80 L mol-1 cm-1, NHCl exhibits two peaks at 330 and 580 nm, ꢀ330 ) (85 (
•
30) L mol-1 cm-1 and ꢀ580 ) (56 ( 30) L mol-1 cm-1. The •NHCl radical undergoes self-decay and can react
•
also with O2 to form a peroxyl radical. It is suggested that the peroxyl radical exists in equilibrium NHClO2
/ •NHCl + O2 with an estimated equilibrium constant of (3 ( 2) × 10-3 mol L-1. The reaction of chloramine
with the carbonate radical is suggested to form a complex [CO3NH2Cl]•- with kf ) 2.5 × 105 L mol-1 s-1
and kr ) 4 × 102 s-1, and this complex decomposes with k ) 7 × 102 s-1 to form NHCl.
•
Introduction
in atmospheric droplets or may be produced by AOT, e.g.
•OH, NO3 , SO3•-, SO4•-, SO5•-, and CO3•-. To evaluate the
•
Residual chlorine in discharged chlorinated water, e.g. cooling
water from power plants, is toxic to aquatic life. In the presence
of ammonia (0.07 to 4 µmol L-1 in seawater), chloramine is
formed.1-4
effect of monochloramine on the decomposition of organic
materials in water by AOT, the rate constants for reactions of
monochloramine with the primary radicals of water radiolysis
were determined.14 The rate constants for reactions of NH2Cl
were reported to be k(eaq-) ) 2.2 × 1010 L mol-1 s-1, k(H) )
Cl2 + H2O / HOCl + H+ + Cl-
NH3 + HOCl / NH2Cl + H2O
(1)
(2)
1.2 × 109 L mol-1 s-1, and k(OH) ) 2.8 × 109 L mol-1 s-1
.
To further elucidate the nature of these reactions and measure
these and other rate constants we have carried out pulse
radiolysis studies with monochloramine in aqueous solutions
under various conditions.
Reaction 2 goes to equilibrium very rapidly and its equilibrium
constant is very high1-4 (K2 ) 3.8 × 1010 L mol-1).5 Chloramine
is a relatively stable oxidant and toxic at low levels (4 µmol
L-1) to certain species of fish.6 Chloramine is also used as a
secondary weak disinfectant in water distribution systems.7,8 It
disappears from the system either by autodecomposition9,10 or
by autocatalytic reduction.11,12 The autocatalytic mechanism was
investigated in the presence of Fe(II) and was suggested to
involve formation of the •NH2 radical as intermediate.11,12 This
radical is a strong oxidant, E°(H+,•NH2/NH3) ) 2.3 V,13 and
may itself oxidize chloramine.
Water may be purified by Advanced Oxidation Technologies
(AOT) to remove organic materials. These processes involve
radical reactions and the presence of chloramine in the water
may have an effect on these radical reactions. The radical
reactions of chloramines are important also in atmospheric
chemistry. Chloramines may be formed in atmospheric droplets
by reaction of chlorine, water, and ammonia via reactions 1
and 2. At pH <5 monochloramine disproportionates to form
dichloramine and, at lower pH, trichloramine.
Experimental Section15
The compounds used were analytical grade reagents from
various sources: NH4Cl (Aldrich), (NH4)2SO4 (Fisher), NaOCl
(10% Cl, Mallinckrodt), NaHCO3 (Fisher), Na2CO3 (Mallinck-
rodt), C6H5CO2H (Aldrich), Na2B4O7 (Aldrich), and NaOH
(50% w/w, Fisher). Water was purified with a Millipore Super-Q
system. Fresh solutions of NH2Cl were prepared from NH4Cl
and NaOCl (1.5 > NH4Cl/NaOCl > 1) at pH ≈9 (5 to 10 mmol
L-1) and were used within 8 h. The concentrations of NH2Cl
and ClO- in the mixture were determined spectrophotometrically
before each experiment (Figure 1), taking ꢀ244 ) 4.6 × 102 L
mol-1 cm-1 for NH2Cl,7,16 and ꢀ292 ) 3.5 × 102 L mol-1 cm-1
for ClO-.17 In all experiments [ClO-] < 0.01[NH2Cl]. The
concentration of NHCl2 under these conditions is negligible.18
Sodium tetraborate was used as buffer to keep the pH near 9.
Highly concentrated solutions of NaOH (free of carbonate) were
used to increase the pH. In experiments with added carbonate,
the pH was adjusted by mixing different proportions of NaHCO3
and Na2CO3. Solutions were irradiated after saturation with N2
or N2O. To study the reactions of radicals with O2, solutions
were saturated with a mixture of N2O/O2. The gas mixture was
prepared in a flow system by mixing N2O and O2. The ratio
N2O/O2 was varied from 1:1 to 1:4. Pulse radiolysis experiments
were carried out with the Linac-based apparatus described
2NH2Cl + H+ f NH4+ + NHCl2
2NHCl2 + H+ f NH3Cl+ + NCl3
NH3Cl+ + NHCl2 f NH4+ + NCl3
(3)
(4)
(5)
Reactions 3-5 are general acid-catalyzed processes.11,12 The
chloramines can react with various radicals that may be present
10.1021/jp030198k This article not subject to U.S. Copyright. Published 2003 by the American Chemical Society
Published on Web 08/22/2003

7424 J. Phys. Chem. A, Vol. 107, No. 38, 2003
Poskrebyshev et al.
Figure 1. Optical absorption spectra of chloramine (pH 9), hypochlorite
Figure 3. Rate constant for reaction of chloramine with hydroxyl
radicals, determined by competition with the reaction of OH with
benzoate ions; the absorbance of the benzoate-OH adduct was monitored
at 350 nm in N2O-saturated solutions containing various concentrations
of benzoate and chloramine at pH 9 with 2 mmol L-1 of borate.
•
(pH 9), carbonate (pH 10.8), and bicarbonate (pH 8.4), all at a
concentration of 1 mmol L-1
.
both the statistical errors ((0.1) and the estimated uncertainty
in the concentration (( 0.2). The concentration of ClO- in the
solution was estimated to be <1% of [NH2Cl] from the
absorbance at 292 nm and thus cannot have any significant
contribution to the measured rate constant. The value of
k(NH2Cl+eaq-) is identical with that reported previously14 and
comparable to k(O2+eaq-) ) 1.9 × 1010 L mol-1 s-1
,
k(BrO-+eaq-) ) 1.5 × 1010 L mol-1 s-1, and k(ClO-+eaq-) )
-
7 × 109 L mol-1 s-1 20
,
but much higher than k(NH2OH+eaq
)
) 9.2 × 108 L mol-1 s-1 20
,
suggesting that chloramine is a
much stronger oxidant than hydroxylamine.
•
-
To study the reactions of OH radicals we saturated the
solutions with N2O to convert the eaq into OH (k9 ) 9.1 ×
Figure 2. Rate constant for the decay of the eaq absorption at 600
nm as a function of [NH2Cl] in N2-saturated solutions containing 0.1
mol L-1 of methanol at pH 9.9.
-
•
109 L mol-1 s-1).20
previously,19 using a quartz cell with 2 cm optical path length.
All experiments were carried out at room temperature, 22 ( 1
°C.
N2O + eaq- (+H+) f N2 + •OH
(9)
The rate constant for reaction of chloramine with •OH radicals
was determined by competition kinetics, by using the reaction
of •OH with benzoate ion as a reference reaction and monitoring
the absorbance of the benzoate OH-adduct in the absence (A0)
and in the presence (A) of various concentrations of chloramine.
Results
The one-electron reduction and oxidation reactions of
monochloramine in aqueous solutions were studied by radiolytic
methods. The radiolysis of aqueous solutions leads to the
production of various radicals and stable products.
- •
PhCO2- + •OH f [HOPhCO2 ]
(10)
(11)
H2O ' •OH, •H, eaq-, H+, H2O2, H2
(6)
NH2Cl + •OH f •NHCl + H2O
The adduct formed by reaction 10 was monitored at 350 nm.21
For γ or high-energy electron irradiation of dilute solutions,
the primary radiation yields are the following: G(eaq-) ) 2.8,
G(•OH) ) 2.8, G(•H) ) 0.6, and G(H2O2) ) 0.75, in units of
- •
At this wavelength [HOPhCO2
] has a molar absorption
coefficient22 of 3.6 × 103 L mol-1 cm-1 while the product of
•
10-7 mol J-1
.
reaction 11, most likely NHCl, has a relatively negligible
absorbance (see below). The rate constant of reaction 10 was
Reactions of Monochloramine with Primary Radicals of
taken to be k10 ) (6.3 ( 0.7) × 109 L mol-1 s-1 20,23
.
From the
-
Water Radiolysis. To study the reactions of eaq we used
•
•
methanol as a scavenger for OH and H (kOH ) 9.7 × 108 L
competition plot of (A0/A - 1) vs the concentration ratio [NH2-
Cl]/[PhCO2-] (Figure 3) we derive an apparent rate constant
of (5.8 ( 0.64) × 108 L mol-1 s-1. This value includes the
contribution of reaction 11 as well as that of the reaction of
mol-1 s-1, kH ) 2.6 × 106 L mol-1 s-1).20
CH3OH + •OH (•H) f •CH2OH + H2O (H2)
(7)
(8)
•
residual NH3 with OH. Residual NH3 exists, despite the high
-
equilibrium constant of reaction 2, because excess NH4Cl was
used in the preparation of NH2Cl to prevent the presence of the
highly reactive ClO-. At [NH2Cl] ) 2.8 × 10-3 mol L-1 we
determined the ratio [NH2Cl]/[NH3] ) 1.39 from the absorbance
of NH2Cl and the initial concentration of NH4Cl used in the
The rate constant for reaction of chloramine with eaq
,
NH2Cl + eaq- f •NH2 + Cl-
-
was determined by following the decay of the eaq absorption
at 600 nm as a function of [NH2Cl] in N2-saturated solutions
containing 0.1 mol L-1 of methanol at pH 9.9 (Figure 2). The
rate constant was found to be diffusion-controlled, k8 ) (2.2 (
0.3) × 1010 L mol-1 s-1, where the indicated error limit includes
preparation. By taking k(•OH+NH3) ) 9.9 × 107 L mol-1 s-1 20
,
we derive k11 ) (5.1 ( 0.6) × 108 L mol-1 s-1. This value is
lower than that derived14 from competition with thiocyanate by
a factor of 5. Although the thiocyanate competition method often

Reactions of Monochloramine in Aqueous Solutions
J. Phys. Chem. A, Vol. 107, No. 38, 2003 7425
NH2Cl, and thus obtain preliminary molar absorption coef-
ficients ꢀ330 ) (85 ( 30) L mol-1 cm-1 and ꢀ580 ) (75 ( 30)
L mol-1 cm-1. The value of ꢀ580 was then corrected for the
•
contribution of NH2 (25-30%) to give ꢀ580 ) (56 ( 30) L
mol-1 cm-1. Chlorine atom substitution thus leads to a 50-nm
red shift in the maximum of the aminyl radical, indicating
considerable interaction between the Cl atom and the unpaired
electron.
Although the second peak of •NHCl is 10 nm lower than the
peak of CO3•-, we have to consider whether part of this
absorption may be due to CO3•-, since the molar absorption
•-
coefficient of CO3 at 600 nm is 30 to 40 times higher than
•
•-
that of NHCl.29 The CO3 radicals may be formed by the
Figure 4. Transient absorption spectra of the •NH2, •NHCl, and
NHClO2• radicals. The spectrum of •NH2 (2) was monitored by pulse
radiolysis of N2O-saturated solution containing 0.1 mol L-1 of ammonia
at pH 11. The other spectra were recorded with solutions of NH2Cl
(3.5 mmol L-1) at pH 9 saturated with N2O (O) or N2O/O2 (b). The
former spectrum (O) is due mainly to the •NHCl radical, and its decay
at 590 nm is shown in the insert. The latter spectrum (b) is due mainly
to the NHClO2• radical. The solid line spectrum between 280 and 400
nm shows the absorption of NH2Cl at the same concentration as the
radicals (23 µmol L-1).
•
reactions of OH radicals with carbonate and bicarbonate that
may be present as impurities (k13 ) 3.9 × 108 L mol-1 s-1
k14 ) 8.5 × 106 L mol-1 s-1).20
,
CO32- + •OH f CO3
(13)
(14)
•-
HCO3- + •OH f CO3
•-
We remeasured the absorption due to CO3•- in a N2O-saturated
solution of 0.1 mol L-1 of carbonate and obtained ꢀ590 ) (1930
( 110) L mol-1 cm-1, in good agreement with the reported
values of ꢀ600, in units of L mol-1 cm-1, 1860,30 1910,31 1830,32
2200,33 and 1500.34 From our preparation method we estimate
that the concentration of any carbonate impurity is <0.01 mmol
L-1. From the absorbance of the solution at 200 and 244 nm
and the known spectra of chloramine2,3,16 and carbonate (Figure
1) we can only estimate that the concentration of carbonate is
<1 mmol L-1. Increasing the pH from 9 to 10.9, to convert
any residual bicarbonate (pKa(HCO3-) ) 10.3) into the more
reactive carbonate, did not affect the transient absorbance at
590 nm. Moreover, addition of 1.5 × 10-3 mol L-1 of carbonate
to the original solution at pH 9 also did not increase the
absorbance at 590 nm. This amount is much more than the
estimated upper limit of the impurity concentration. These
experiments indicate that the peak at 580 nm can be ascribed
has provided reliable rate constants for hydroxyl radical reac-
tions, it is likely that in the case of chloramine a thermal reaction
with thiocyanate before the radiolysis may have led to erroneous
results. Monochloramine is known to react with iodide, bromide,
and cyanide ions18,24,25 and is most likely to react with
thiocyanate as well. The competition with benzoate does not
suffer from such side reactions and we believe that the present
result is more reliable. Our value of k(•OH+NH2Cl) is higher
than k(•OH+NH3) by almost an order of magnitude but lower
than k(•OH+NH2OH)14,20 ) 9 × 109 L mol-1 s-1 by more than
•
an order of magnitude. The reactivity pattern toward OH is
similar to that observed for the substitution of Cl and OH on
methane in the gas phase,26 or on acetic acid in water,20 and
supports an abstraction mechanism.
The rate constant for reaction of monochloramine with
hydrogen atoms was reported14 to be k(H•+NH2Cl) ) (1.2 (
0.1) × 109 L mol-1 s-1 and was not determined in the present
study. In principle, the reaction may take place via abstraction
of hydrogen or chlorine. However, because of its high rate
•
to the NHCl radical and that the contribution of carbonate is
negligible.
The decay of •NHCl was monitored, in N2O-saturated
solutions containing two different concentrations of NH2Cl, at
340 and 590 nm (Figure 4 insert) and the kinetic curves were
fitted to a simple second-order rate law. With [NH2Cl] ) 7.2
× 10-3 mol L-1, where most of the •OH radicals are scavenged
by chloramines, but also ca. 40% of the eaq- are scavenged by
chloramine in competition with N2O, the observed rate constants
at the two wavelengths were 2k340 ) (4.1 ( 1.0) × 109 L mol-1
s-1 and 2k590 ) (2.6 ( 1.0) × 109 L mol-1 s-1, the uncertainty
ranges primarily reflecting the uncertainties in the molar
absorption coefficients. It should be pointed out, however, that
these rate constants include contributions from other radical-
radical reactions. By using [NH2Cl] ) (3.6 ( 0.2) × 10-3 mol
•
constant, by comparison to that of OH radicals, the reaction
was suggested14 to involve chlorine abstraction.
NH2Cl + H• f •NH2 + H+ + Cl-
(12)
•
Absorption Spectrum and Self-Reaction of the NHCl
Radical. The measurement of the absorption spectrum of •NHCl
and the determination of the self-reaction kinetics are compli-
cated by the relative slowness of the reaction of NH2Cl with
•OH and the rapidity of its reaction with eaq-. This leads to the
need to correct for incomplete scavenging on one hand and for
•
absorption due to NH2 on the other. To do so, first, we
L-1, where only ca. 25% of the eaq are scavenged by
-
redetermined the spectrum of •NH2 by pulse irradiating a N2O-
saturated solution containing 0.1 mol L-1 of ammonia at pH
chloramine, and monitoring the decay at 590 nm we found 2k
) 2.8 × 109 L mol-1 s-1 at pH 9; 2.7 × 109 L mol-1 s-1 at pH
10.2; and 2.2 × 109 L mol-1 s-1 at pH 10.9, i.e., independent
of pH within the above stated uncertainties. These measured
values of 2k(•NHCl) include a contribution from mixed decay
•
11. The spectrum of the NH2 radical was found to have only
one peak at λmax ) 530 nm (ꢀ530 ) 80 L mol-1 cm-1) (Figure
4, solid triangles), in agreement with previous determina-
tions.27,28 Then we determined the spectrum of •NHCl from the
pulse radiolysis of N2O-saturated solutions containing [NH2-
Cl] ) 3.5 × 10-3 mol L-1 at pH 9. Two peaks were observed,
at λmax ) 330 and 580 nm (Figure 4, open circles). We first
correct the concentration of •NHCl for the incomplete scaveng-
ing of the •OH radicals and for the partial reaction of eaq- with
with NH2 radicals and are close to the value of 2k(•NH2) )
•
(4.4 ( 1) × 109 L mol-1 s-1 27,28
.
The decay of the •NH2 radicals involves combination to form
hydrazine.28 Similarly, the decay of •NHCl radicals may proceed
via combination to form initially the unstable dichlorohydrazine

7426 J. Phys. Chem. A, Vol. 107, No. 38, 2003
Poskrebyshev et al.
(ClNHNHCl). This product may eliminate two molecules of
HCl to leave N2. Alternatively, it may eliminate only one
molecule of HCl to leave the unstable HNdNCl, which may
hydrolyze to form nitrosyl chloride (NOCl) and then nitrous
acid (HNO2) or nitrite ions. The latter products may be oxidized
by the remaining NH2Cl to form nitrate. In fact, reaction of
aqueous ammonia with excess chlorine is known to produce
-
- 35
N2 along with smaller amounts of NO2 and NO3
.
•-
Reaction of Chloramine with CO3 Radicals. To study
the reaction of monochloramine with the carbonate radical we
•-
followed the decay of CO3 in N2O-saturated solutions with
various NH2Cl concentrations. With [carbonate] ) 0.1 mol L-1
and [NH2Cl] ) (1.5 to 7.3) × 10-4 mol L-1 at pH 10.24, most
•
•-
of the OH radicals react with carbonate. The decay of CO3
took place in two steps (Figure 5a,b) and the rate of each step
was dependent on the initial [NH2Cl]. The decay curves could
not be fitted to two consecutive irreversible first-order reactions
involving an intermediate because the absorbance at the
intermediate level as well as the decay in the slower step were
dependent on [NH2Cl]. Therefore, we assume that the absor-
•-
bance at the end of the fast step is due mainly to CO3
remaining in equilibrium and that the decay must involve more
complex kinetics which can be fitted only by computer
modeling. To model the observed kinetics in this complex
system we took into account the reactions of •OH radicals with
carbonate (reaction 13, k13 ) 3.9 × 108 L mol-1 s-1 20
)
and
with chloramines (reaction 11, k11 ) 5.1 × 108 L mol-1 s-1),
the reactions of eaq with N2O (reaction 9, k9 ) 9.1 × 109 L
-
mol-1 s-1 20
)
and with chloramine (reaction 8, k8 ) 2.2 × 1010
L mol-1 s-1), and the reaction of H• atoms with NH2Cl (reaction
12, k12 ) 1.2 × 109 L mol-1 s-1), neglecting the competing
reaction of H• with OH- (k ) 2.5 × 107 L mol-1 s-1). We also
took into account the self-reaction of NH2 (2k ) 4.4 × 109 L
•
•-
mol-1 s-1),27 •NHCl (2k ) 2.4 × 109 L mol-1 s-1), and CO3
(2k determined to be 1.1 × 107 L mol-1 s-1, in agreement with
Figure 5. Kinetic traces showing the decay of the carbonate radical
absorption at 600 nm in the presence of chloramine at pH 10.2: (a
and b) 0.1 mol L-1 of carbonate and (c) 0.01 mol L-1 of carbonate.
The concentration of NH2Cl, in mmol L-1, is indicated on the right of
previous values36), and an estimated value for the radical-
radical reaction CO3 + NH2 (k ) 1.5 × 109 L mol-1 s-1).
The observed kinetics were best fit by assuming an equilibrium
reaction 15,
•-
•
each trace; the radical concentration was 1.6 × 10-6 mol L-1
.
value is about half the rate constant estimated above for the
CO3•- + NH2Cl / [CO3NH2Cl]•-
(15)
•-
•
cross reaction between CO3 and NH2 radicals.
Reaction of •NHCl with O2. To study the reaction of •NHCl
radicals with O2, we carried out pulse radiolysis experiments
with NH2Cl solutions saturated with a mixture of N2O and O2.
with k15 ) 2.5 × 105 L mol-1 s-1 and k-15 ) 4 × 102 s-1
(K15 ) 625 L mol-1 and ∆G15 ) -16.5 kJ mol-1), and
decomposition of the complex with k16 ) 7 × 102 s-1
[CO3NH2Cl]•- f HCO3- + •NHCl
.
•
The spectrum of the expected peroxyl radical, NHClO2 , was
measured with different concentrations of O2. Two peaks were
observed, at 350 and 580 nm (Figure 4, solid circles). By
correcting the concentration of •NHCl for the partial scavenging
(16)
We estimate that the uncertainties in the rate and equilibrium
constants for these reactions from this model study are better
than a factor of 2. The observed kinetics and the above model
also permit us to estimate an upper limit for the rate constant
of reaction 17.
-
of the •OH radicals by NH3 and for the partial reaction of eaq
•
with NH2Cl (see above), and assuming that all NHCl radicals
reacted with O2 (at [O2] ) 4.8 × 10-4 mol L-1), we determined
a molar absorption coefficient of ꢀ350 ) (190 ( 30) L mol-1
cm-1. It was noticed, however, that increasing the concentration
of O2 leads to an increase in the absorbance at 350 nm and a
decrease at 580 nm. This may be interpreted as an increase in
•NH2 + HCO3- f NH3 + CO3
(17)
•-
•
[NHClO2 ] accompanied by a decrease in [•NHCl] as a result
By assuming that <10% of the •NH2 radicals react with
carbonate (to fit the model), we estimate k17 < 5 × 103 L mol-1
s-1, close to an earlier estimate.37
of a fast equilibrium reaction.
•
NHClO2 / •NHCl + O2
(18)
At a lower concentration of carbonate (0.01 mol L-1), the
•OH radicals react with both carbonate and chloramine and then
the contribution of the cross-reaction CO3•- + •NHCl becomes
important. A rate constant of 6 × 108 L mol-1 s-1 was derived
from the modeling to fit the kinetic traces (Figure 5c) and is
estimated to have an uncertainty of about a factor of 2. This
The alternative explanation of competition between the forma-
tion of NHClO2• and its subsequent fast decay by reaction with
•NHCl is ruled out by the fact that the rate of decay at 350 nm
did not increase with the concentration of O2. Formation of

Reactions of Monochloramine in Aqueous Solutions
J. Phys. Chem. A, Vol. 107, No. 38, 2003 7427
TABLE 1: Comparison of Experimental and Calculated
Absorbance at 350 nm
TABLE 2: Calculated N-H Bond Dissociation Energies and
•
Rate Constants for OH Reactions
[O2] (10-4 mol L-1
)
Aexp
Acalc
compd
(S)
UHF/6-31G*
CBS-Q
k(OH+S)
(kJ mol-1
)
(kJ mol-1
)
(L mol-1 s-1
)
0
0.0043
0.0057
0.0063
0.0064
0.0065
0.0062
0.0063
0.0071
0.0081
0.0043
0.0052
0.0062
0.0064
0.0064
0.0065
0.0065
0.0072
0.0079
NH3
NH2Cl
NH2OH
337
298
265
444.5
379.8
349
8.2 × 107
5.2 × 108
9.5 × 109
1.5
2.2
2.6
2.7
2.9
3.1
4.3
5.5
∆fG°(aq NH2)13 ) 192 kJ mol-1, ∆fG°(aq OH-)38 ) -157.3
kJ mol-1, and ∆fG°(aq NH2OH)41 ) -23.4 kJ mol-1). Thus
the reduction potential of chloramine is more positive than that
of hydroxylamine by about 0.7 V, in line with the much higher
-
NH2O2• should not be important in this system because the rate
rate constant for reaction of chloramine with eaq
.
The rate constant for the reaction of NH2Cl with •OH radicals
is higher than that of NH3 by almost an order of magnitude
and lower than that of NH2OH by more than an order of
magnitude, which suggests a hydrogen abstraction mechanism.
The N-H bond dissociation energy of NH3 is42 455 kJ mol-1
but the values for NH2OH and NH2Cl are not known. Therefore,
we computed all three values in the gas phase by using Gaussian
9843 with two approximation methods (Table 2). Although the
CBS-Q method gives a value for NH3 that is very close to the
experimentally determined value, both methods show that the
N-H bond dissociation energy increases in the order NH2OH,
NH2Cl, NH3. The measured rate constants decrease in that order,
in line with an H-abstraction mechanism.
•
constant for the reaction of NH2 with O2 (k-18 ) 3 × 108 L
mol-1 s-1) is slower than that observed in our experiments.
From the dependence of the absorbance at 350 nm upon [O2]
in a N2O/O2 saturated solution of chloramine (3.5 × 10-3 mol
L-1), by fitting the measured absorbance (A) with the equation
•
A/(l[R]) ) {ꢀ350(•NHCl) + ꢀ350(NHClO2 )K18[O2]}/(1 + K18-
[O2]), where l is the optical path length and [R] is the initial
•
concentration of NHCl (26.9 µmol L-1), we derive an equi-
librium constant K18 ) (3 ( 2) × 10-3 mol L-1 (∆G18 ) 14.5
kJ mol-1) and a molar absorption coefficient ꢀ350(NHClO2 ) )
•
(1000 ( 600) L mol-1 cm-1. Table 1 shows the agreement
between the experimentally determined absorbance and that
calculated from the above parameters at different O2 concentra-
tions. From K18 we estimate that at [O2] ) 5 × 10-4 mol L-1
From the N-H bond dissociation energies calculated by the
CBS-Q method (Table 2), if we assume that the differences
between the respective enthalpies are the same in water as in
the gas phase and that the entropies of dissociation of these
bonds in NH3, NH2Cl, and NH2OH are similar, we can estimate
the differences between the potentials E(NH3•+/NH3), E(•NH2-
Cl+/NH2Cl), and E(•NH2OH+/NH2OH). The reduction potential
E(NH3• +/NH3) ) 2.13 V was calculated by Stanbury,13 who
chose pKa(NH3•+) ) 2.3. From this value and the above
assumptions we estimate for chloramine E(•NHCl,H+/NH2Cl)
) 1.7 V (pKa(NH2Cl•+) is unknown) and for hydroxylamine
E(•NH2OH+/NH2OH) ) 1.03 V or E(•NHOH,H+/NH2OH) )
1.28 V (by using pKa(NH2OH•+) ) 4.2). A slightly lower
reduction potential, E(•NH2OH+/NH2OH) ) 0.91 V, was
estimated before,40 by using ab initio calculation of ∆fH(NH2O),
semiempirical (AM1) calculation of S(NH2O), and assuming
that the free energy of hydration of NH2O and NH3 are the same
(-10 kJ mol-1). From the two values we suggest that E(•NH2-
OH+/NH2OH) lies between 0.9 and 1.1 V.
•
only 17% of the radicals are present as NHClO2 , and if we
assume that this radical does not absorb at 580 nm, we expect
the absorbance at this wavelength to decrease by 17% in the
presence of this concentration of O2. The decrease, in fact, was
by about 10%, in reasonable agreement within the estimated
•
uncertainties. The spectrum of NHClO2 , with λmax ) 350 nm,
•
is close to that reported for NH2O2 (λmax ) 340 nm). The decay
•
of the absorbance at 350 nm is due to the self-decay of NHClO2
•
and of NHCl and their cross reaction and the kinetic curves
were too noisy to permit clear distinction between first-order
•
and second-order rate laws. The self-decay of NHClO2 may
produce O2 and unstable products such as ClHNOONHCl or
the radical NHClO• and subsequent reactions may lead to
formation of stable products such as nitrite or nitrate.
Discussion
-
The rate constant for the reaction of NH2Cl with eaq is
comparable to the rate constants20 for reactions of O2 and BrO-
but much higher than that of NH2OH, indicating that mono-
chloramine is a much stronger oxidant than hydroxylamine. The
one-electron reduction potential of monochloramine E(NH2Cl/
•NH2,Cl-) can be estimated from ∆fG(Cl-) (-131.2 kJ mol-1),38
In summary, chloramine reacts very rapidly with solvated
•
electrons to form NH2 and somewhat more slowly with the
•
hydroxyl radical to form NHCl. Chloramine reacts with the
carbonate radical to form an intermediate complex and then
•NHCl. On the basis of the results with carbonate it is expected
that other strong oxidants also will react with chloramine to
∆fG(•NH2) (192 kJ mol-1),13 and ∆fG(NH2Cl) (70.3 kJ mol-1
)
(this latter value is derived from the most recent publication5
and a measure of its uncertainty is reflected in the earlier
reported values: 84,17 77.4,39 and 66.93 kJ mol-1). From these
values we calculate E(NH2Cl/•NH2,Cl-) ) 0.1 V. This potential
is close to E(HOBr/HOBr -) ) 0.14 V and higher than E(H2O2/
OH,OH-) ) -0.03 V, E(HO2-/O-,OH-) ) -0.04 V, E(O2/
O2-) ) -0.16 V, and E(BrO-/Br-,O-) ) -0.24 V; all of these
oxidants react with hydrated electrons with practically diffusion-
controlled rate constants. Hydroxylamine reacts more slowly
•
•
•
produce NHCl. Both NH2 and NHCl radicals decay via
radical-radical reactions and can react with O2 to form peroxyl
•
radicals. The NH2O2 radical has been shown before to eliminate
•
water and produce NO.27 The fate of the NHClO2 radical is
•
unclear because it is present in a mixture with NHCl and the
decay involves various reactions.
References and Notes
(1) Weil, I.; Morris, J. C. J. Am. Chem. Soc. 1949, 71, 1664.
(2) Margerum, D. W.; Gray, E. T., Jr.; Huffman, R. P. In Organometals
and Organometalloids, Occurrence and Fate in the EnVironment; ACS
Symp. Ser. No. 82; Brinckman, F. E., Bellama, J. M., Eds.; American
Chemical Society: Washington, DC, 1978; pp 278-291.
(3) Gray, E. T., Jr.; Margerum, D. W.; Huffman, R. P. In Organometals
and Organometalloids, Occurrence and Fate in the EnVironment; ACS
with eaq- because E°(NH2OH/NH2 ,OH-) is more negative than
•
these values. Although E°(NH2OH/NH2 ,OH-) was estimated
•
to be -0.12 V,13 the value of ∆fG(NH2OH(aq)) used in that
estimation has been corrected,40 and this correction leads to
E°(NH2OH/NH2 ,OH-) ≈ -0.6 V (based on the value
•

7428 J. Phys. Chem. A, Vol. 107, No. 38, 2003
Poskrebyshev et al.
Symp. Ser. No. 82; Brinckman, F. E., Bellama, J. M., Eds.; American
Chemical Society: Washington, DC, 1978; pp 264-277.
(4) Armesto, X. L.; Canle, M.; Garcia, M. V.; Santaballa, J. A. Chem.
Soc. ReV. 1998, 27, 453-460.
(28) Pagsberg, P. B. InVestigation of the NH2 radical produced by pulse
radiolysis of ammonia in aqueous solution; RISO Report No. 256, 1972; p
209.
(29) Carbonate impurities have led investigators to suggest the formation
of borate (e.g.: Padmaja, S.; Ramakrishnan, V.; Rajaram, J.; Kuriacose, J.
C. J. Chem. Soc., Faraday Trans. 1 1989, 85, 2249. Padmaja, S.; Rajaram,
J.; Ramakrishnan, V.; Kuriacose, J. C. J. Photochem. Photobiol. A 1990,
52, 1451) and silicate radicals (Kuz’min, V. A.; Chibisov, A. K. High Energy
(5) Margerum, D. W.; Schurter, L. M.; Hobson, J.; Moore, E. E.
EnViron. Sci. Technol. 1994, 28, 331-337.
(6) Farrell, A. P.; Kennedy, C.; Cheng, W.; Lemke, M. A. Water Qual.
Res. J. Can. 2001, 36, 133-149.
•-
Chem. 1970, 84, 1146) upon reaction of SO4 in basic solutions, with
(7) Norton, C. D.; LeChevallier, M. W. J. Am. Water Works Assoc.
1977, 89, 66-77.
absorptions near 600 nm. We have subsequently discovered (unpublished
results) that the reported absorptions were due solely to carbonate radicals
and that the other radicals were not formed.
(30) Behar, D.; Czapski, G.; Duchovny, I. J. Phys. Chem. 1970, 74,
2206.
(31) Eriksen, T. E.; Lind, J.; Merenyi, G. Radiat. Phys. Chem. 1985,
26, 197.
(32) Chen, S.; Cope, V. W.; Hoffman, M. Z. J. Phys. Chem. 1973, 77,
1111.
(33) Czapski, G.; Holcman, J.; Bielski, B. H. J. J. Am. Chem. Soc. 1994,
116, 11465.
(34) Huie, R. E.; Clifton, C. L. J. Phys. Chem. 1990, 94, 8561.
(35) Pressley, T. A.; Bishop, D. F.; Roan, S. G. EnViron. Sci. Technol.
1972, 6, 622.
(36) Neta, P.; Huie, R. E.; Ross, A. B. J. Phys. Chem. Ref. Data 1988,
17, 1027. NDRL/NIST Solution Kinetics Database, NIST Standard Refer-
ence Database 1040, 1995. http://kinetics.nist.gov/solution/index.php.
(37) Draganic, Z. D.; Negron-Mendoza, A.; Sehested, K.; Vujosevic,
S. I.; Navarro-Gonzales, R.; Albarran-Sanchez, M. G.; Draganic, I. G.
Radiat. Phys. Chem. 1991, 38, 317-321.
(38) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.;
Halow, I.; Bailey, S. M.; Churney, K. L.; Nutall, R. L. J. Phys. Chem. Ref.
Data 1982, 11, Suppl. No. 1.
(8) Wolf, R. L.; Ward, N. R.; Olson, B. H. J. Am. Water Works Assoc.
1984, 75, 74-88.
(9) Valentine, R. L.; Ozekin, K.; Vikesland, P. J. Chloramine Decom-
position in Distribution System and Model Waters; AWWARF and
AWWA: Denver, CO, 1998.
(10) Vikesland, P. J.; Ozekin, K.; Valentine, R. L. Water Res. 2001,
35, 1766-1776.
(11) Vikesland, P. J.; Valentine, R. L. EnViron. Sci. Technol. 2000, 34,
83-90.
(12) Vikesland, P. J.; Valentine, R. L. EnViron. Sci. Technol. 2002, 36,
662-668.
(13) Stanbury, D. M. AdV. Inorg. Chem. 1989, 33, 69.
(14) Johnson, H. D.; Cooper, W. J.; Mezyk, S. P.; Bartels, D. M. Radiat.
Phys. Chem. 2002, 65, 317-326.
(15) The mention of commercial equipment or material does not imply
recognition or endorsement by the National Institute of Standards and
Technology, nor does it imply that the material or equipment identified are
necessarily the best available for the purpose.
(16) Yiin, B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 2135-2141.
(17) Soulard, M.; Block, F.; Hatterer, A. J. Chem. Soc., Dalton Trans.
1981, 2300.
(18) White, G. C. The Handbook of Chlorination and AlternatiVe
Disinfectants, 4th ed.; Wiley: New York, 1999.
(19) Grodkowski, J.; Neta, P. J. Phys. Chem. A 2000, 104, 4475.
(20) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B.
J. Phys. Chem. Ref. Data 1988, 17, 513. NDRL/NIST Solution
Kinetics Database, NIST Standard Reference Database 540, 1995.
http://kinetics.nist.gov/solution/index.php.
(21) Neta, P.; Dorfman, L. M. AdV. Chem. Ser. 1968, 81, 222.
(22) Wander, R.; Neta, P.; Dorfman, L. M. J. Phys. Chem. 1968, 72,
2946.
(39) Jolly, W. L. J. Phys. Chem. 1956, 60, 507.
(40) Hung, M.-L.; McKee, M. L.; Stanbury, D. M. Inorg. Chem. 1994,
33, 5108; correction in Inorg. Chem. 2001, 5140, 2206.
(41) Latimer, W. M. The Oxidation States of the Elements and Their
Potentials in Aqueous Solutions, 2nd ed.; Prentice Hall: Englewood Cliffs,
NJ, 1964.
(42) Gurvich, L. V.; Veyts, I. V.; Alcock, C. B. Thermodynamic
Properties of IndiVidual Substances; Hemisphere Publishing: New York,
1991.
(43) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb,
M. A.; Cheeseman, J. R.; Zakrzewski, V. G.; Montgomery, J. A., Jr.;
Stratmann, R. E.; Burant, J. C.; Dapprich, S.; Millam, J. M.; Daniels, A.
D.; Kudin, K. N.; Strain, M. C.; Farkas, O.; Tomasi, J.; Barone, V.; Cossi,
M.; Cammi, R.; Mennucci, B.; Pomelli, C.; Adamo, C.; Clifford, S.;
Ochterski, J.; Petersson, G. A.; Ayala, P. Y.; Cui, Q.; Morokuma, K.; Malick,
D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Cioslowski, J.;
Ortiz, J. V.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi,
I.; Gomperts, R.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.;
Peng, C. Y.; Nanayakkara, A.; Gonzalez, C.; Challacombe, M.; Gill, P. M.
W.; Johnson, B.; Chen, W.; Wong, M. W.; Andres, J. L.; Gonzalez, C.;
Head-Gordon, M.; Replogle, E. S.; Pople, J. A. Gaussian 98; Gaussian,
Inc.: Pittsburgh, PA, 1998.
(23) Poskrebyshev, G. A.; Neta, P.; Huie, R. E. J. Phys. Chem. A 2002,
106, 11488.
(24) Trofe, T. W.; Johnson, J. D.; Inman, G. W., Jr. The Kinetics of
Monochloramine Decomposition in the Presence of Bromide; Nuclear
Regulatory Commission: Washington, DC, Report No. NUREG/
CR-1116RE, 1980.
(25) Schurter, L. M.; Bachelor, P. P.; Margerum, D. W. EnViron. Sci.
Technol. 1995, 29, 1127-1134.
(26) NIST Standard Reference Database 17, http://kinetics.nist.gov/
index.php.
(27) Laszlo, B.; Alfassi, Z. B.; Neta, P.; Huie, R. E. J. Phys. Chem. A
1998, 102, 8498.