Electrochemical behavior of chloramines on the rotating platinum and gold electrodes

Article abstract of DOI:10.1149/1.1560951

Electroreduction of chloramines (mono-, di-, and trichloramine) in 1 M NH₄Cl solutions of different pH was investigated at the rotating platinum and gold electrodes. It was found that all chloramines are present in the solution in nonprotonated forms and give well-formed one-step or two-step current-potential waves. The final products of reduction are ammonium (or ammonia) and chloride ions. Monochloramine is reduced in a single two-electron irreversible wave. Hydrazine is not an intermediate in monochloramine reduction. Dichloramine reduction generally proceeds in two two-electron steps (via monochloramine). Below pH 4.3 a kinetic current due to the protonated dichloramine reduction (single four-electron wave) is in force, appearing as an increase of the height of the first step on lowering pH. Due to this process below pH 2.5 only one four-electron reduction wave is observed. Trichloramine reduction occurs in two steps: two-electron trichloramine to dichloramine reduction and four-electron dichloramine reduction. In strong acidic solutions the kinetic current due to the protonated trichloramine reduction has to be taken into account. A reaction mechanism common for all chloramines was proposed with [NXCl·] as an intermediate (X = H₂, HCl, and Cl₂ for mono-, di-, and trichloramine, respectively). The rate-determining step does not involve proton transfer.

Full text of DOI:10.1149/1.1560951

Electrochemical Behavior of Chloramines on the Rotating
Platinum and Gold Electrodes
Barbara Piela and Piotr K. Wrona
J. Electrochem. Soc. 2003, Volume 150, Issue 5, Pages E255-E265.
doi: 10.1149/1.1560951
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2003 ECS - The Electrochemical Society
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Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
E255
0
013-4651/2003/150͑5͒/E255/11/$7.00 © The Electrochemical Society, Inc.
Electrochemical Behavior of Chloramines on the Rotating
Platinum and Gold Electrodes
a
a,b,z
Barbara Piela and Piotr K. Wrona
a
Faculty of Chemistry, Warsaw University, Pasteura 1, 02-093 Warsaw, Poland
Industrial Chemistry Research Institute, Rydygiera 8, 01-793 Warsaw, Poland
b
Electroreduction of chloramines ͑mono-, di-, and trichloramine͒ in 1 M NH Cl solutions of different pH was investigated at the
4
rotating platinum and gold electrodes. It was found that all chloramines are present in the solution in nonprotonated forms and give
well-formed one-step or two-step current-potential waves. The final products of reduction are ammonium ͑or ammonia͒ and
chloride ions. Monochloramine is reduced in a single two-electron irreversible wave. Hydrazine is not an intermediate in
monochloramine reduction. Dichloramine reduction generally proceeds in two two-electron steps ͑via monochloramine͒. Below
pH 4.3 a kinetic current due to the protonated dichloramine reduction ͑single four-electron wave͒ is in force, appearing as an
increase of the height of the first step on lowering pH. Due to this process below pH 2.5 only one four-electron reduction wave
is observed. Trichloramine reduction occurs in two steps: two-electron trichloramine to dichloramine reduction and four-electron
dichloramine reduction. In strong acidic solutions the kinetic current due to the protonated trichloramine reduction has to be taken
into account. A reaction mechanism common for all chloramines was proposed with ͓NXCl•͔ as an intermediate (X ϭ H , HCl,
2
and Cl for mono-, di-, and trichloramine, respectively͒. The rate-determining step does not involve proton transfer.
2
©
2003 The Electrochemical Society. ͓DOI: 10.1149/1.1560951͔ All rights reserved.
Manuscript submitted July 24, 2002; revised manuscript received October 22, 2002. Available electronically March 18, 2003.
Chloramines are formed when molecular chlorine or hypochlo-
rite is added to aqueous solutions that contain ammonia or ammo-
nium ions. The type of chloramine formed depends on the ratio of
The current relations for the respective reduction waves in different
pH ranges are given and the diffusion coefficients of the chloram-
ines are calculated.
1
chlorine concentration to ammonia concentration and is also af-
fected by pH. Monochloramine is predominant in alkaline solu-
2
Experimental
tions. Dichloramine prevails in the neutral media, whereas
trichloramine is formed in acidic solutions.
Materials used.—The experiments were conducted in 1 M
NH Cl with addition of proper amounts of the following solutions:
4
Chloramines are important from the point of view of environ-
mental protection. For example, they can be formed during the pro-
pH Ͻ 2.5: sulfuric acid solution of different concentrations; pH 3.6-
5
.6: acetate buffer ͑x* 0.2 M acetic acid ϩ y* 0.2M sodium ac-
etate͒; pH 5.3-8.3: phosphate buffer ͑x* 0.2 M KH PO ϩ y*
3,4
cess of water chlorination. All chloramines are toxic, trichloram-
ine being the most harmful, and special care should be taken in
order to avoid its formation.
2
4
0
.2 M K HPO ); and pH 8.3-14: potassium hydroxide solution of
2 4
different concentrations.
Industrial interest in chloramines is focused mainly on formation
The x/y proportion of the buffer components was adjusted to
obtain a specific pH value. NH Cl ͑POCh, Gliwice͒, K SO ͑POCh,
5
,6
of hydrazine, which is widely used as a reducing agent, an energy-
rich compound, or hydrogen storage medium. Since
4
2
4
a
Gliwice͒, KOH ͑Chemical Institute, O s´ wi e¸ cim͒, NaOH ͑Chemical
Institute, O s´ wi e¸ cim͒, and H SO ͑POLCHEM, Toru n´ ͒ were used, all
monochloramine is an intermediate in the Raschig synthesis of hy-
drazine, understanding of chloramine chemistry is essential to the
understanding of the Raschig process. Chloramines are also used in
the synthesis of chlorine derivatives of s-triazine, cyanuric chloride,
which are used as swimming pool disinfectants, bleaches, cleansers,
sanitizers, and also for food processing.
2
4
of analytical purity quality.
The NaClO solution was prepared by passing gaseous chlorine
through 1 M NaOH solution and refrigerated ͑ϳ5°C͒ because of its
instability ͑possibility of NaClO formation͒. The concentration of
3
refrigerated sodium hypochlorite solution remained unchanged for
A complete review of the chemistry of chloramines is not avail-
able and their electrochemical behavior is not well understood.
There are only a few papers on the electrochemistry of chloramines
at solid electrodes.7 The only analytical aspect of this work was
related to the trichloramine/chlorine system. In studies on platinum
electrodes it was shown7 that all chloramines are electroactive.
They are reduced in a single irreversible peak: two-electron, four-
electron, and six-electron in the case of mono-, di-, and trichloram-
ine, respectively. The products of the reaction are ammonium ions or
ammonia, and chloride ions.
In this work the behavior of chloramines is studied by discussing
the synthesis, spectrometric detection, ranges of predominance, and
electrochemical response at rotating platinum and gold electrodes.
The existence of different forms of chloramines ͑protonated, non-
protonated, and deprotonated͒ is considered. Special attention is
paid to the mechanism of electroreduction in order to employ the
results obtained for further analytical studies.9 Overall reaction
schemes and detailed reaction mechanisms for each chloramine are
proposed. Side effects due to homogeneous chemical reactions
coupled with the electrochemical processes are taken into account.
six months.
We used triply distilled water additionally passed through a
Milli-Q system.
,8
Procedures.—Preparation and identification.—All experiments
were performed in 1 M NH Cl solution ͑see preceding section͒.
Chloramines ͑or chlorine͒ were synthesized by mixing 0.2 mL 0.5
M NaClO solution with 50 mL 1 M NH Cl solution ͑initial hy-
pochlorite concentration equal to total active chlorine concentration,
c0hyp ϭ c0Cl ϭ 2 mM) of a proper pH according to the following
equations
,8
4
4
Ϫ
Ϫ
NH ϩ ClO → NH Cl ϩ OH
͓1͔
͓2͔
͓3͔
͓4͔
3
2
ϩ
ϩ
NH4 ϩ 2HClO → NHCl ϩ H ϩ 2H O
2
2
ϩ
ϩ
3HClO ϩ NH4 → NCl3 ϩ 3H2O ϩ H
Ϫ
ϩ
HClO ϩ Cl ϩ H → Cl2 ϩ H2O
The products formed were identified spectrophotometrically on
the basis of the characteristic absorption bands ͑see Table I͒. The
obtained results are presented in Scheme 1a.
z
E-mail: pkwrona@chem.uw.edu.pl
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E256
Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
between the chlorine compounds and are dependent on the initial
hypochlorite ͑equal to the total active chlorine͒, ammonia ͑ammo-
nium ions͒, and chloride concentrations.
Table I. The values of the molar absorption coefficient „␧max) for
chloramines, hypochlorite, and chlorine.
The established domains of existence are consistent with the lit-
erature values of the mono/dichloramine and di/trichloramine equi-
librium constants (Km/d and Kd/t , respectively͒
␧max
͑dm mol cmϪ1
͒
3
Ϫ1
Reference
Compound
␭max ͑nm͒
NH Cl
245
416
445
455
457
439
13
14
15
2
2
ϩ
ϩ
4
2
2
2
2
44
45
44
44
2
NH Cl ϩ H ϭ NHCl ϩ NH
Km/d ϭ 1
2
2
ϫ 109 molϪ1 dm3 10
͓5͔
͓6͔
This work
ϩ
ϩ
4
3
NHCl ϩ H ϭ 2NCl ϩ NH
Kd/t ϭ 1.7
2
3
NHCl2
206
2100
1700
265
267
300
280
265
15
2
13
14
15
2
ϫ 104 molϪ1 dm3 2
Kd/t ϭ 1.0 ϫ 104 molϪ1 dm3 11
207
297
294
295
294
294
On the basis of the literature value of the trichloramine/chlorine
equilibrium constant (Kt/Cl) one predicts less acidic conditions for
chlorine predominance than those established by us
This work
NCl3
220
8100
6100
5000
255
270
260
285
195
199
15
2
22
13
14
15
2
22
221
220
340
336
340
337
336
336
NCl ϩ 3Cl ϩ 4H ϭ 3Cl ϩ NH Kt/Cl ϭ 10 molϪ4 dm12 2
Ϫ
ϩ
ϩ
4
3
2
͓7͔
Quantitative analysis.—The concentration of the chloramines was
determined spectrophotometrically in aqueous solutions on the basis
of the estimated ͑iodometric titration͒ molar absorption coefficients
at the maximum of the characteristic band ␧max ͑see Table I͒.
In order to estimate the molar absorption coefficients, chloram-
ines were synthesized as described in the preceding section, except
for the initial hypochlorite concentration ͑equal to total active chlo-
rine concentration͒, which in this case was equal to 4, 12, and 27
mM for mono-, di-, and trichloramine, respectively. The domains of
chloramine existence specific to such conditions are presented in
Scheme 1b.
This work
Ϫ
ClO
292
350
350
350
18
23
2
2
2
90
92
HClO
230
100
100
97
18
23
2
2
2
35
33
The molar absorption coefficient (␧max) was taken as an average
value in the range of each chloramine predominance. Its value was
checked by iodometric titration. Because the literature data show
Cl2
325
23
75
68
16
2
3
12
Ϫ
that mono- and dichloramine react with iodide quantitatively,
whereas the iodine yield with trichloramine is only 82% of the the-
Cl3
220
10000
18500
193
16
2
16
2
233
325
325
oretical amount, the results obtained for NCl were corrected assum-
3
85
ing such efficiency. Values of ␧max obtained for mono- and
2
,13,14,15
dichloramine are consistent with the literature data,
whereas
for trichloramine it is an average of 15% lower than in the
2
,13,14,15
literature.
The discrepancy may arise from the rapid decom-
position of NCl for lower NH Cl concentration ͑we have employed
3
4
1
M NH Cl) or the presence of other chlorine species.
4
The spectrophotometric method was also applied in order to de-
Ϫ
termine the ClO and Cl concentrations. The molar absorption co-
2
3
Ϫ1
Ϫ1
efficients were taken from the literature: 350 dm mol cm at
Ϫ2
3
Ϫ1
Ϫ1
16
2
90 nm for ClO and 75 dm mol cm at 325 nm for Cl2 ͑see
Table I͒.
Acid-base properties of chloramines.—Theoretically, chloramines
ϩ
can exist in simple, NY, or protonated NHY forms (Y ϭ H Cl,
2
HCl , or Cl for mono-, di-, and trichloramine, respectively͒. In the
2
3
Ϫ
case of mono- and dichloramine deprotonated forms (NHCl and
Ϫ
2
NCl ) should be also considered. Protonated and deprotonated
chloramines should have UV-visible ͑UV-vis͒ absorption spectra dif-
ferent from NY, which would introduce changes in measured absor-
bance.
There is a significant disagreement in the literature concerning
the protonation and deprotonation equilibrium constants. The acid
Scheme 1. Domains of existence of chloramines obtained for 1 M NH Cl
4
and the initial hypochlorite concentration used for the generation of chloram-
ines ͑equal to total active chlorine concentration͒ ͑a͒ equal to 2 mM and ͑b͒
equal to 4, 12, and 27 mM in the case of mono-, di-, and trichloramine,
respectively, where m is monochloramine, d dichloramine, and t stands for
trichloramine. It was assumed that in the predominance range of the indi-
vidual chlorine species its content in the mixture exceeds 95%.
Ϫ13
dissociation constant for monochloramine falls between 10
and
Ϫ18 11,17,18
10
.
This means that monochloramine can be present as
NHCl only in strongly basic solutions. Jolly estimated the respec-
tive constant for dichloramine to be 10 ; however, this was con-
tradicted experimentally by Hand and Margerum. Weil and
Morris proposed 10 for the monochloramine protonation equilib-
rium constant, whereas Wrona found 10 for monochloramine and
Ϫ
11
Ϫ7
1
0
The range of predominance for each compound was determined
assuming that its content among all chlorine species exceeds 95%. It
is important to notice here that such results are due to the equilibria
1
9
1
7
Ϫ6
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Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
0Ϫ3 for dichloramine. Wrona’s results should be taken with cau-
E257
1
tion, because they were based on electrochemical measurements in
nonbuffered solutions. It is probable that they were influenced by
local changes of pH in the vicinity of the electrode surface during
the electrode process.
The following analysis illustrates that all chloramines are present
in the studied solutions ͑for the pH range of each chloramine pre-
dominance and its mixture with other chlorine species, see Scheme
1a͒ only in simple NY forms. First, the molar absorption coeffi-
cients, estimated on the basis of iodometric titration, do not depend
on pH in the respective chloramine predominance range. This can be
regarded as evidence for the absence of protonated and deprotonated
forms in this pH range. Second, the efficiency ͑with respect to the
initial hypochlorite concentration͒ of each chloramine formation in
its predominance range was close to 100%. Mixtures of various
chlorine species were obtained also with the same total efficiency.
The calculations were based on the estimated molar absorption co-
efficients.
Experiment.—A precision digital pH meter ͑Radelkis, Hungary͒ and
a Hydromet ͑Gliwice, Poland͒ combination electrode were used for
pH measurements. The potential of the hydrogen electrode was mea-
sured instead of the direct pH measurement. pH was then obtained
from a calibration potential-pH curve ͑the pH meter was calibrated
on six buffers in the pH range from 3 to 12͒.
Electrochemical experiments were performed using a CHI elec-
trochemical system ͑USA͒. The following rotating disk electrodes
͑
RDE͒ were used: a polycrystalline platinum RDE ͑Pine Instru-
2
ments, USA͒, surface area 0.170 cm , and a polycrystalline gold
RDE ͑custom made͒, surface area 0.156 cm . A sodium chloride
2
saturated calomel electrode ͑SSCE͒ was used as the reference elec-
trode. All potentials are given vs. the potential of this electrode. The
2
counter electrode was a Pt wire with a ϳ3 cm surface area, sepa-
Figure 1. Current-potential curves obtained for chloramine reduction at
rated from the working solution by a compartment filled with the
2
͑top͒ the rotating platinum and ͑bottom͒ gold electrodes in 1 M NH Cl
4
supporting electrolyte solution. An Au wire ͑surface area ϳ2 cm ͒
solution of different pH: ͑curve 1͒ pH 7.6, pure monochloramine solution,
cm ϭ 1.8 mM; ͑curve 2͒ pH 5.1, mono- and dichloramine solution, cm
ϭ 0.6 mM, c ϭ 0.65 mM; ͑curve 3͒ pH 3.5, dichloramine solution, c
was used for electrolysis as the working electrode. During electroly-
sis the reference and the counter electrodes were the same as de-
scribed earlier. The electrolysis was performed by maintaining the
working electrode at a fixed potential and the flowing charge was
measured.
d
d
ϭ 1.0 mM; ͑curve 4͒ pH 2.9, dichloramine solution, c ϭ 1.0 mM; and
d
͑
curve 5͒ pH 1.1, di- and trichloramine solution, cd ϭ 0.7 mM, ct
ϭ 0.2 mM. w1, w2, and w3 correspond to the waves of mono-, di- and
trichloramine reduction, respectively. Rotation speed f ϭ 3000 rpm. The
currents were corrected for the supporting electrolyte background current.
Because there was no difference in the electrochemical response
Ϫ1
for scan rates in the range 0.02-0.1 V s ͑for f ϭ 3000 rpm), the
highest scan rate was used in the experiments. For lower rotation
speeds we used lower scan rates, if necessary. The results were
corrected for ohmic drop. Spectrophotometric measurements were
made with a Hewlett-Packard 8452a ͑USA͒ spectrophotometer, and
all experiments were carried out at 25 Ϯ 1°C. The working solu-
tions were deoxygenated with pure ͑5 N͒ nitrogen.
The following dependencies were obtained for the
monochloramine solution: ͑i͒ The number of electrons involved in
the reduction process is equal to 1.97 Ϯ 0.05. This finding is based
on electrolyses performed at a gold wire electrode ͑at potential equal
to Ϫ0.4 V, pH range 7.3-14͒. The results were corrected for the loss
of monochloramine concentration due to its spontaneous decompo-
sition during electrolysis. ͑ii͒ The only products of the reduction are
ammonia ͑or ammonium ions͒ and chloride ions. There was no dif-
ference between the spectra of the monochloramine solution re-
corded after the complete electrolysis and those of the supporting
electrolyte solution. ͑iii͒ i1 is linearly dependent on the
monochloramine concentration ͑Fig. 5, top͒. (iv) i1 is linearly de-
pendent on the square root of the rotation speed ͑Fig. 5, bottom͒. (v)
The diffusion coefficient, calculated from the slope of current vs.
concentration and current vs. square root of the rotation speed de-
Results and Discussion
Typical cyclic voltammograms recorded for all chloramines at
the rotating platinum and gold electrodes are presented in Fig. 1 ͑top
and bottom, respectively͒. Marks w1, w2, and w3 correspond to the
waves of heights i , i , and i , respectively, and the potential po-
1
2
3
sition and the apparent transfer coefficient are shown in Fig. 2 ͑top
and bottom͒. The UV-vis spectra corresponding to the polarization
curves are shown in Fig. 3.
As it follows from Fig. 1, all chloramines are reduced in well-
formed one-step or two-step waves. The overall reaction schemes
and the kinetic reaction mechanisms of each chloramine’s reduction
are discussed in the following two sections, respectively.
The overall reaction schemes for mono-, di-, and trichloramine
reduction are shown in Fig. 4. Detailed analysis is presented.
Ϫ5
2
Ϫ1
pendencies, is equal to (1.52 Ϯ 0.10) ϫ 10 cm s
.
Such results led us to propose the following overall
monochloramine reduction scheme:
For pH Ͻ 9.2
Ϫ
ϩ
4
Ϫ
Ϫ
NH Cl ϩ 2e ϩ 2H O → NH ϩ Cl ϩ 2OH
͑w1͒
2
2
Electroreduction of monochloramine.—Monochloramine is the
͓
8͔
only chlorine species present in 1 M NH Cl solution at pH above
4
7
.3 ͑for c0hyp ϭ c0Cl ϳ 2 mM). In such solutions a single irrevers-
For pH Ͼ 9.2
ible polarization wave w1 ͑of current height i ) is recorded ͑Fig. 1,
1
Ϫ
Ϫ
Ϫ
curve 1͒.
NH Cl ϩ 2e ϩ H O → NH ϩ Cl ϩ OH
͑w1͒ ͓9͔
2
2
3
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E258
Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
Figure 3. Chloramine absorption spectra: ͑a͒ The absorption spectrum cor-
responding to curve 1 of Fig. 1a and b: ͑ ͒ monochloramine charac-
teristic absorption band. ͑b͒ The absorption spectrum corresponding to curve
Figure 2. ͑top͒ The dependence of the half-wave potential of the chloramine
reduction waves at the rotating ͑᭡, ᭿, ᭹͒ platinum and ͑᭝, ᮀ, ᭺͒ gold
electrodes on pH. ͑bottom͒ The dependence of the apparent cathodic transfer
coefficient (␣n) on pH. Chloramine reduction at the rotating ͑᭡, ᭿, ᭹͒
2
͑
of Fig. 1a and b: ͑
• • • •
͒ monochloramine characteristic absorption band,
͒ dichloramine characteristic absorption band. ͑c͒ The absorption
spectrum corresponding to curves 3 and 4 of Fig. 1a and b: ͑
dichloramine characteristic absorption band. ͑d͒ The absorption spectrum
corresponding to curve 5 of Fig. 1a and b: ͑ ͒ dichloramine charac-
͒
platinum and ͑᭝, ᮀ, ᭺͒ gold electrodes in 1 M NH Cl: ͑᭹, ᭺͒
4
monochloramine reduction ͑w1͒; ͑᭿, ᮀ͒ dichloramine reduction ͑w2͒; ͑᭡,
᭝
͒ trichloramine and chlorine reduction ͑w3͒. ␣n and the half-wave poten-
teristic absorption band, ͑• • • •͒ trichloramine characteristic absorption
tials were obtained on the basis of the numerical fit of the theoretical equa-
tion for an irreversible reduction process to the experimental data ͑polariza-
tion wave͒. The domains of existence of chloramines ͑separated by the
band.
dashed vertical lines͒ were obtained for 1 M NH Cl and the initial hypochlo-
4
rite concentration ͑equal to total active chlorine concentration͒ equal to 2
mM. It was assumed that in the predominance range of the individual chlo-
rine species its content in the mixture exceeds 95%.
top͒. Spectroscopic investigations show that for pH 2.3-4.2
dichloramine is the only chlorine species present in solution ͑Fig.
3c͒. While no quantitative changes in the spectrum occur, the total
current also remains unchanged ͑see inset to Fig. 6, top͒. However
the current ratio i /i decreases with lowering pH ͑Fig. 7, top͒,
2
1
and the expression for i as follows
1
providing that for solutions of pH below 2.5 only a four-electron
reduction wave is present ͑Fig. 6, top͒. Some of the relations ob-
tained differ from those for more basic solutions ͑pH above 4.2͒.
Now the current ratio i2 /i1 changes with the electrode rotation
speed ͑Fig. 6, bottom͒. This indicates that some kinetic effect should
be taken into account in the acidic pH range.
2/3
i1 ϭ 2pDm cm
͓10͔
p ϭ 0.61 ϫ 10Ϫ3FA␻1/2␯Ϫ1/6 ϭ 0.61 ϫ 10Ϫ3FA͑2␲f/60͒1/2␯Ϫ1/6
͓11͔
where cm is the monochloramine concentration ͑mol dmϪ3͒, Dm is
2
Ϫ1
the monochloramine diffusion coefficient ͑cm s ͒, F the Faraday
pH Ͼ 4.2.—In the dichloramine/monochloramine solutions ͑pH
2
Ϫ1
constant, A electrode area ͑cm ͒, ␻ the angular rotation speed ͑s ͒,
range 4.2-7.3͒ two reduction waves w1 and w2 of heights i and i2
1
2
Ϫ1
are observed ͑Fig. 3, curves 2 and 3, and Fig. 6, top͒. In these curves
it is seen that i2 is proportional to dichloramine concentration and
the difference (i Ϫ i ) is proportional to monochloramine concen-
f rotation speed ͑rpm͒, and v is kinematic viscosity ͑cm s ͒.
Electroreduction of dichloramine.—Dichloramine is formed
when the monochloramine solution is acidified (pH Ͻ 7.3, Eq. 5͒.
Above pH 4.2, for an initial hypochlorite concentration close to 2
mM, dichloramine exists in mixture with monochloramine and their
concentration ratio increases with lowering pH. The electrochemical
response is the sum of dichloramine and monochloramine contribu-
tions. Dichloramine is reduced via monochloramine in two two-
electron waves ͑w1 and w2͒ of equal heights ͑Fig. 4͒. Wave w1
overlaps with that of bulk monochloramine ͑also w1͒ ͓Fig. 1 ͑curves
1
2
tration with a slope equal to that obtained in the case of pure
monochloramine solution ͑Fig. 8, top͒. Taking the dichloramine dif-
fusion coefficient Dd to be equal to (1.30 Ϯ 0.03)
Ϫ5
2
Ϫ1
ϫ 10 cm s from the i vs. dichloramine concentration depen-
2
dence one gets the number of electrons for w2 to be close to 2.
Figure 8 shows that both i and (i Ϫ i ) are linear functions of the
2
1
2
square root of the rotation speed, and in Fig. 6 ͑bottom͒ it is seen
that the i /i ratio does not depend on the rotation speed.
2
1
2-5͒, and Fig. 6, top͔. The electrochemical behavior in this pH range
Taking into account these dependencies we can write the follow-
ing overall reaction scheme
does not seem to be complicated and corresponds to a simple irre-
versible process.
The electrochemical response of the system is quite different in
more acidic solutions ͑for pH Ͻ 4.2) and is affected by pH ͑Fig. 6,
ϩ
Ϫ
Ϫ
NHCl ϩ H ϩ 2e → NH Cl ϩ Cl
͑w2͒
͓12͔
2
2
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Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
E259
Figure 4. The overall reaction schemes found for each chloramine reduction
at the rotating gold ͑platinum͒ electrode in 1 M NH Cl solution of different
4
pH. Current-potential curves were simulated for equal chloramine concen-
trations. Bold font style-overall equation, regular italic font style-reaction
steps.
ϩ
Ϫ
ϩ
Ϫ
NH Cl ϩ 2H ϩ 2e → NH4 ϩ Cl
͑w1͒
͓13͔
2
Figure 5. ͑top͒ i1 dependence on the monochloramine concentration. Inset:
log i vs. log cm dependence. Rotation speed f ϭ 3000 rpm. ͑bottom͒ i1 de-
pendence on the rotation speed of the electrode. Monochloramine concentra-
tion cm ϭ 2.1 mM. Monochloramine reduction at the rotating ͑᭹͒ platinum
and ͑᭺͒ gold electrodes in 1 M NH4Cl at pH 10.
and the expressions for i and i1
2
2/3
i2 ϭ 2pDd cd
͓14͔
͓15͔
2/3
2/3
i1 ϭ 2pDd cd ϩ 2pDm cm
Ϫ3
where c is the dichloramine concentration ͑mol dm ͒, and D is
the dichloramine diffusion coefficient ͑cm s ͒.
d
d
(vi) The i /i ratio exceeds unity and increases with lowering
2
1
2
Ϫ1
pH ͑Fig. 7, top͒.
(
vii) The total current (i ϩ i ) divided by absorbance mea-
pH Ͻ 4.2.—As stated previously, in the pH range 2.3-4.2
dichloramine is the only chlorine species present in the solution.
Two reduction waves, w1 and w2 ͑for pH below 2.5 only one w2͒ of
heights i and i , respectively, are present ͑Fig. 6 top͒. The follow-
1 2
sured at 294 nm ͑maximum of the dichloramine characteristic ab-
sorption band͒ does not depend on pH ͑inset to Fig. 6, top͒.
(
viii) The i /i ratio is greater for smaller rotation speeds of the
2 1
1
2
electrode ͑Fig. 6, bottom͒. The effect is more pronounced at lower
pH.
From these findings, except the last two, it can be deduced that
dichloramine undergoes protonation ͑Eq. 16͒ and the resulting pro-
tonated form is further reduced in one four-electron irreversible
wave ͑Eq. 17͒
ing dependencies were observed.
i͒ The number of electrons involved in the overall dichloramine
͑
reduction process is equal to 4.08 Ϯ 0.15. This was found from
electrolysis performed at the gold wire electrode (E
ϭϪ0.4 V, pH 2.3-4.2). The results were corrected for the loss of
dichloramine concentration due to its spontaneous decomposition
during electrolysis.
k
1
ϩ
ϩ
2
NHCl ϩ H ꢀ NH Cl
͓16͔
͑
ii͒ The only products of the reduction are ammonium and chlo-
2
2
k
2
ride ions. There was no difference between the spectra of the
dichloramine solution recorded after complete electrolysis and the
spectrum of the supporting electrolyte solution.
ϩ
2
ϩ
Ϫ
ϩ
Ϫ
NH Cl ϩ 2H ϩ 4e → NH4 ϩ 2Cl
͓17͔
͓18͔
2
͑
iii͒ The total current (i ϩ i ) is linearly dependent on the
K ϭ K /K
ϭ cpd /͑cdcHϩ͒ ϭ k /k
H O 1
1
2
p
b
2
2
dichloramine concentration ͑Fig. 9, top͒.
(
iv) The total current (i ϩ i ) is linearly dependent on the
cd ϭ cpd ϩ cnd
͓19͔
1
2
square root of the rotation speed ͑Fig. 9, bottom͒.
v) The diffusion coefficient, calculated from the slope of the
i ϩ i ) vs. concentration and (i ϩ i ) vs. square root of the ro-
where k1 is the dichloramine protonation rate constant ͑mol dm3
Ϫ1
(
Ϫ1
Ϫ1
(
s ͒; k2 is the dichloramine deprotonation rate constant ͑s ͒; Kp is
1
2
1
2
Ϫ1
3
tation speed dependencies, assuming a four-electron process, is
the dichloramine protonation equilibrium constant ͑mol dm ͒; Kb
is the dichloramine basic ionization constant; K is the activity
Ϫ5
2
Ϫ1
equal to (1.30 Ϯ 0.15) ϫ 10 cm s
.
H O
2
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E260
Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
Figure 7. ͑top͒ The current ratio (i /i ) vs. pH plot. ͑bottom͒ i /(i
1
2
1
2
ϩ i ) dependence on pH. ͑᭹͒ Data obtained for dichloramine or, for pH
2
above 4.2, its mixture with monochloramine, reduction at the rotating plati-
num electrode. ͑᭺͒ Data obtained for dichloramine or, for pH above 4.2, its
mixture with monochloramine, reduction at the rotating gold electrode.
͑
͒ The numerical fit of Eq. 27 to the experimental data obtained for
pure dichloramine reduction at the rotating platinum electrode. ͑• • • •͒ The
numerical fit of Eq. 27 to the experimental data obtained for pure dichloram-
ine reduction at the rotating gold electrode. Rotation speed f ϭ 3000 rpm.
Initial hypochlorite concentration, c0hyp ϭ 2 mM.
Figure 6. ͑top͒ Voltammograms for pure dichloramine ͑pH 2.4, 2.9, 3.5, 4.1,
thick lines͒ and dichloramine/monochloramine mixture ͑pH 4.5 and 5.1, thin
lines͒ reduction at the rotating gold disk electrode in NH Cl solution. Initial
4
hypochlorite concentration c0hyp ϭ 2 mM. Rotation speed f ϭ 3000 rpm.
The currents were corrected for the supporting electrolyte background cur-
dichloramine should occur at lower wavelengths compared to the
nonprotonated form. As stated previously, no qualitative ͑and for a
constant total current also quantitative, see inset to Fig. 6, top͒
changes with pH in the dichloramine UV-vis spectrum were ob-
served.
rent. Inset: the influence of pH on the total current (i1 ϩ i ) divided by
2
absorbance measured at 294 nm ͑maximum of dichloramine characteristic
adsorption band͒. Dichloramine predominance range. ͑bottom͒ The current
ratio i /i vs. rotation speed of the electrode. Experimental conditions ͑ex-
2
1
cept the rotation speed͒ are the same as in the top figure.
The i /i ratio dependence on the rotation speed indicates that
2
1
kinetic effect should be considered here. It is probable that the pro-
tonated dichloramine reduction current could not be proportional to
the bulk concentration of this form. This can happen when during
the reduction process the decrease of the concentration of this form
of dichloramine pushes the equilibrium of Reaction 16 to the right.
At pH 2.3-4.2, the chemical equilibrium is probably shifted largely
product of water ͑mol dmϪ6͒; cnd is the nonprotonated dichloram-
2
Ϫ3
ine concentration ͑mol dm ͒; cpd is the protonated dichloramine
Ϫ3
concentration ͑mol dm ͒; and c is the total dichloramine concen-
d
Ϫ3
tration ͑mol dm ͒. The four-electron wave overlaps with the wave
corresponding to the first step of nonprotonated dichloramine reduc-
tion ͑all w2͒.
The electrochemical response in such a case is the sum of the
contributions of the two forms of dichloramine and may be de-
scribed by the following equations
to the left ͑small K value͒, so that practically only nonprotonated
p
dichloramine is present in solution ͑no changes in UV-vis spectrum͒.
ϩ
2
However, the NH Cl reduction current ͑kinetic current, i ) can be
2
k
observed, provided that the magnitude of the protonation rate con-
stant k1 is large. This effect is stronger ͑greater value of current
plateau͒ in more acidic solutions ͑the rate of the protonation reaction
is pH-dependent͒ and for lower electrode rotation speed ͑see Fig. 6,
bottom͒.
2
d
/3
2/3
2/3
2/3
i2 ϭ 2pD cnd ϩ 4pDd cpd ϭ 4pDd cd Ϫ 2pDd cnd ͓20͔
i1 ϭ 2pD /3cnd
2
͓21͔
d
The kinetic current effect has been described in the literature.20
A
It was assumed that the dichloramine diffusion coefficient (D ) is
identical for both forms.
typical example of this type of process is the electrochemical reduc-
tion of formaldehyde in buffered solutions. In aqueous solutions
formaldehyde exists mostly in the hydrated form, which is not re-
duced at the mercury electrode. Only the anhydrous form is electro-
active, and therefore at the beginning of the process, when the equi-
librium between those two forms is disturbed at the surface,
formation of the anhydrous form takes place and current flows.
d
However, such considerations are not in agreement with the i /i
2
1
ratio dependence on the rotation speed, which is not predicted by
Eq. 20 and 21. Also, the spectroscopic investigations exclude the
possibility of existence of protonated dichloramine in the examined
pH range. The characteristic absorption band of that form of
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Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
E261
Figure 8. ͑top͒ i2 ͑᭹, ᭺͒ and (i1 Ϫ i ) ͑᭿, ᮀ͒ dependence on the
2
dichloramine and monochloramine concentration, respectively. Rotation
speed f ϭ 3000 rpm. ͑᭹, ᭿͒ The rotating platinum electrode; ͑᭺, ᮀ͒ the
rotating gold electrode. bottom: i2 ͑□͒ and (i1 Ϫ i ) ͑᭹͒ dependence on the
2
Figure 9. ͑top͒ The total current (i1 ϩ i ) dependence on dichloramine
concentration. Rotation speed f ϭ 3000 rpm. ͑᭹, ᭡, ᭢, ᭿͒ the rotating
platinum electrode; ͑᭺, ᭝, ᭞, ᮀ͒ the rotating gold electrode. ͑bottom͒ i1
rotation speed of the gold electrode. Monochloramine concentration cm
ϭ 1.7 mM, dichloramine concentration c ϭ 0.3 mM. 1 M NH Cl at pH 6.
2
d
4
͑
᭺͒, i2 ͑᭹͒, and i1 ϩ i2 ͑ᮀ͒ dependence on the rotation speed of the gold
electrode. Dichloramine concentration c ϭ 1 mM, 1 M NH Cl at pH 3.0.
d
4
Dichloramine protonation-Estimation of kinetic parameters.—The
reaction discussed in Eq. 16 and 17 can be considered as an elec-
trode process preceded by a pseudo first-order chemical reaction ͑for
constant hydrogen ion concentration͒. The kinetic equation can be
written as follows20
where i2a denotes the current of the first step of nonprotonated
dichloramine reduction.
The estimation of K and k could be done on the basis of the
ik /␻ vs. ik dependence. Such dependencies obtained for pure
dichloramine solution of different pH are presented in Table II and
p
1
i ␻1 ϭ ͑i ϩ i ͒/␻ Ϫ i b/͓K c ͑k1c
/2
1/2
ϩ k2͒1/2͔
1/2
k
1
2
k
p
2
Hϩ
Hϩ
ϭ ͑i ϩ i ͒/␻ Ϫ i b/͓K k c ϩ k K c ͔1/2
1/2
3
3
2
1
2
k
p
1
Hϩ
1
p Hϩ
͓22͔
b ϭ ͑D /␯͒1/6/1.61
͓23͔
d
Table II. Studies on dichloramine protonation. The results of the
1Õ2
numerical fit of Eq. 22 „i Õ␻ vs. ik dependence… to the experi-
k
mental data,a and dichloramine reduction at the rotating gold
electrode.
where ik denotes the kinetic current due to the protonated
dichloramine reduction ͑amperes͒. Because the electroreduction po-
tential of the protonated dichloramine reduction wave is similar to
the potential of the first wave of nonprotonated dichloramine reduc-
tion, only one total wave is observed experimentally ͑w2͒. The
Range of the rotation
Dichloramine
conc. ͑mM͒
Slope ϫ 10Ϫ3
pH
speed ͑rpm͒
2.7
2.9
3.1
3.3
3.5
3.8
4.1
500-5000
200-5500
800-5500
400-5500
200-5000
300-6000
200-5000
0.9
1.0
0.8
0.9
0.9
0.8
0.8
10.6 Ϯ 0.4
19.6 Ϯ 1.0
26.3 Ϯ 1.9
40.1 Ϯ 1.9
72.3 Ϯ 5.5
105.8 Ϯ 9.7
280.7 Ϯ 28.8
height of this wave (i ) is equal to the sum of currents due to
individual processes
2
i ϭ i ϩ i ϭ i ϩ i1
͓24͔
2
k
2a
k
or
i ϭ i Ϫ i
͓25͔
a
k
2
1
See also Fig. 10.
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E262
Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
Figure 11. Current-potential curves obtained for trichloramine/dichloramine
mixture ͑pH 0.9, 1.3͒, chlorine/trichloramine/dichloramine mixture ͑pH
Ϫ0.2, 0.3͒ and chlorine ͑pH Ϫ0.5͒ reduction at the rotating gold disk elec-
trode in 1 M NH Cl solution. Initial hypochlorite concentration c
4
0hyp
ϭ 2 mM. Rotation speed f ϭ 3000 rpm. The currents were corrected for the
supporting electrolyte, background current. Inset: trichloramine (ct
ϭ 6 mM,
in 1 M NH Cl solution.
͒ and chlorine (cC12 ϭ 15 mM, • • • •͒ absorption spectra
4
1
/2
Figure 10. ͑top͒ i /␻ vs. ik dependence ͑see Eq. 22͒. ͑bottom͒ log B (B
k
2
3
2
Hϩ
ϭ K k c ϩ k K c ) as a function of pH ͑see Eq. 26͒. Dichloramine ͑1
trichloramine in the examined solution (1 M NH Cl,c ϭ c0hyp
p
1
Hϩ
1
p
4
0Cl
M NH Cl solutions of different pH͒ reduction at the rotating gold electrode.
4
ϭ 2 mM). Spectroscopic investigations show that NCl exists here
3
only in the mixture with dichloramine ͑less acidic solutions͒ and/or
chlorine ͑more acidic solutions, see Scheme 1a͒. The chlorine ab-
sorption band lies near the trichloramine band ͑inset to Fig. 11͒,
which makes the simultaneous determination of chlorine and
trichloramine difficult. Because of fast trichloramine decomposition
it is also hard to investigate the rotation speed influence on the
observed current response.
in Fig. 10, top. The reciprocal of the slope multiplied by b ͑see Eq.
2
3
2
2
3͒ and squared is equal to K k c ϩ k K c
and is further
p
1
Hϩ
1
p
Hϩ
denoted as B.
Figure 10 ͑bottom͒ presents log B as a function of pH. The de-
pendence is linear with the slope close to 2 (1.9 Ϯ 0.1). It can be
2
3
concluded that in the studied pH range the K k c factor is much
p
1
Hϩ
pH 0.5-2.3.—In solutions of pH 0.5-2.5 trichloramine exists in a
mixture with dichloramine and their concentration ratio increases
with lowering pH. Electrochemical experiments show the presence
of two waves, w2 and w3. The following dependencies were ob-
tained in this case: ͑i͒ The current ratio i /i increases with higher
2
smaller than K k c
and can be neglected. The dependence sim-
p
1
Hϩ
plifies then to the form
log B ϭ log͑k K ͒ Ϫ 2pH
͓26͔
1
p
3
2
acidity. For pH ϳ0.5 it reaches 0.4 ͑Fig. 12, top͒. ͑ii͒ The i3 vs.
trichloramine concentration dependence is linear ͑Fig. 12, bottom͒.
From the intercept of log B vs. pH dependence one extracts the k K
1
p
7
Ϫ2
6
Ϫ1
value equal to (6.2 Ϯ 4.8) ϫ 10 mol dm s
.
͑
iii͒ The slopes of i vs. trichloramine concentration dependence for
3
Similar results can be obtained from a numerical fit of the func-
tion given by Eq. 27 to the experimental data presented in Fig. 7,
bottom. By combining Eq. 22 and 25 the following relation is ob-
tained
both electrodes correspond to a two-electron process ͑assuming the
trichloramine diffusion coefficient D ϭ (1.15 Ϯ 0.07)
ϫ 10 cm s ). (iv) The difference (i Ϫ i ) was found to be a
t
Ϫ5
2
Ϫ1
2
3
linear function of dichloramine concentration with the same slope as
the (i1 ϩ i2) vs. cd dependence in the pH range 2.3-4.2 ͑Fig. 12,
bottom͒.
Considering the results for this range of pH we can write the
overall reaction scheme as follows
i /͑i ϩ i ͒ ϭ 0.5 ϩ 0.5/
͕
1 ϩ ␻1/2b/͓KpcHϩ͑k1cHϩ ϩ k2͒1/2͔
͖
2
1
2
͓27͔
The analysis performed for the pH range of pure dichloramine ex-
7
Ϫ2
6
Ϫ1
istence gives k K equal to (6.7 Ϯ 0.2) ϫ 10 mol dm s for
1
p
ϩ
Ϫ
Ϫ
7
Ϫ2
6
Ϫ1
NCl ϩ H ϩ 2e → NHCl ϩ Cl
͑w3͒
͓28͔
͓29͔
the gold and (4.3 Ϯ 0.2) ϫ 10 mol dm s for the platinum
3
2
electrode.
ϩ
ϩ
Electroreduction of trichloramine.—Trichloramine forms in the
most acidic solutions (pH Ͻ 2.3) and simultaneously a third wave
w3 forms on the current-potential curves ͑Fig. 11͒. It is impossible
to record an electrochemical response corresponding to pure
NHCl ϩ H ϭ NH Cl
2 2
2
ϩ
2
ϩ
Ϫ
ϩ
Ϫ
NH Cl ϩ 2H ϩ 4e → NH4 ϩ 2Cl
͑w2͒ ͓30͔
2
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Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
E263
tive analysis difficult. The second phenomena is that small amounts
of dichloramine are present in solution, providing for an increase of
i2 . Third, the kinetic effect due to trichloramine protonation ͑as in
the case of dichloramine͒
ϩ
ϩ
NCl ϩ H ϭ NHCl3
͓34͔
͓35͔
3
ϩ
ϩ
Ϫ
ϩ
Ϫ
NHCl3 ϩ 3H ϩ 6e ϭ NH4 ϩ 3Cl
The kinetic current due to the protonated trichloramine reduction,
ik1 , may increase the height of wave w3.
Taking into account all three possibilities we can propose the
following relations for i and i
2
3
2
/3
2/3
i3 ϭ 2pDt ct ϩ 2pDCl2cCl2 ϩ ik1
͑w3͒
͑w2͒
͓36͔
͓37͔
2
/3
2/3
i2 ϭ 4pD cd ϩ 4pDt ct Ϫ ik1
d
It is difficult to distinguish whether the observed increase of the
current ratio i /i with decreasing pH ͑above 0.5͒ is due to the
3
2
change of the trichloramine reduction mechanism or to additional
chlorine reduction. It is also probable that both effects are present.
Kinetic reaction mechanisms.—It was found that the following
kinetic reaction mechanism is obeyed for each chloramine ͑NXCl͒
reduction
Ϫ
Ϫ
NXCl ϩ e → ͓NX•͔ ϩ Cl
rate-determining step
͓
38͔
ϩ
ϩ
͓
NX•͔ ϩ H → ͓NXH ͔
͓39͔
͓40͔
͓41͔
ϩ
Ϫ
͓
NXH ͔ ϩ e → NXH
NXCl ϩ H ϩ 2e → NXH ϩ Cl
ϩ
Ϫ
Ϫ
Figure 12. ͑top͒ The current ratio i /i2 vs. pH plot. Initial hypochlorite
where X denotes H , HCl, and Cl for mono-, di-, and trichloram-
2 2
3
concentration c0hyp ϭ 2 mM. ͑bottom͒ i3 ͑᭹, ᭺͒ and i2 Ϫ 2i3 ͑᭿, ᮀ͒ de-
pendence on trichloramine and dichloramine concentration, respectively,
pH ϳ 1. ͑᭹, ᭿͒. The rotating platinum electrode; ͑᭺, ᮀ͒ the rotating gold
ine, respectively.
In the case of di- (X ϭ HCl) and trichloramine (X ϭ Cl ) a sub-
strate adsorption step is additionally involved
2
electrode. 1 M NH Cl. Rotation speed f ϭ 3000 rpm.
4
NXCl → ͓NXCl͔ads
͓42͔
Trichloramine is reduced in two steps ͑waves w3 and w2͒ via
dichloramine. The second step overlaps with that accounting for
bulk dichloramine reduction ͑all w2͒. The following equations for i3
and the adsorbed intermediates ͓NXCl͔ads
,
͓NX•͔ads
,
and
ϩ
͓
NXH ͔ further participate in Reactions 38-40.
ads
Detailed analysis of the kinetic mechanism for each chloramine
and i hold
reduction is presented in the following section.
2
2
/3
Monochloramine reduction mechanism.—On the basis of the overall
reaction scheme presented in the Electroreduction of monochloram-
ine section, the following results were obtained: ͑i͒ The reaction is
i3 ϭ 2pDt ct
͓31͔
͓32͔
2/3
2/3
i2 ϭ 4pDd cd ϩ 4pDt ct
electrochemically irreversible, the slope ץE/ץ log(i Ϫ i )/i is close
g
where c denotes trichloramine concentration ͑M͒.
to 200 mV. ͑ii͒ The apparent transfer coefficient (␣n), calculated
from the numerical fit of the i vs. E equation for an irreversible
process to the experimental data, is close to 0.3 for both electrodes
t
pH-0.5 to 0.5.—In this pH range a mixture of di-, trichloramine, and
chlorine exists. The electrochemical response of the system does not
qualitatively change compared to that obtained for the dichloramine/
trichloramine mixture. Two waves w2 and w3 form and the i /i
current ratio increases with higher acidity ͑Fig. 12, top͒.
Apart from trichloramine reduction, one should now take into
account three additional phenomena. The first is formation of mo-
͑Fig. 3, bottom͒. ͑iii͒ The half-wave potential (E1/2) for the platinum
disk is the same as for the gold one ͑ϳ Ϫ0.2 V, Fig. 3, top͒. (iv)
The changes of E1/2 with pH are small. For pH Ͻ 9.5 a 10 mV shift
of E1/2 toward more positive values per unit increment of pH was
observed; for pH Ͼ 9.5 the dependence was the opposite. (v) The
reaction order with respect to monochloramine is unity
3
2
lecular chlorine from NCl decomposition
3
(
ץ log i/ץ log cmon ϭ 1, inset to Fig. 5, top͒.
Ϫ
ϩ
ϩ
NCl ϩ 3Cl ϩ 4H ϭ 3Cl ϩ NH4
͓33͔
3
2
We can propose two probable reaction mechanisms for
monochloramine reduction:
Its presence does not qualitatively alter the electrochemical image of
trichloramine reduction. Chlorine is reduced in a two-electron wave
at the same potential as wave 3 ͑see Fig. 11͒ and increases the i3
value. The amount of chlorine increases with decreasing pH which
Ϫ
Ϫ
I. NH Cl ϩ e → ͓NH •͔ ϩ Cl
rate-determining step ͓43͔
2
2
ϩ
Ϫ
͓
NH •͔ ϩ H O → ͓NH ͔ ϩ OH
͓44͔
͓45͔
͓46͔
2
2
3
may be one of the reasons for the i /i2 increase. The chlorine ab-
ϩ
Ϫ
3
͓NH ͔ ϩ e → NH3
3
sorption band ͑maximum at 325 nm͒ occurs near the trichloramine
band ͑maximum at 336 nm, inset to Fig. 11͒. This makes quantita-
Ϫ
Ϫ
Ϫ
NH Cl ϩ H O ϩ 2e → NH ϩ Cl ϩ OH
2
2
3
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E264
or
Journal of The Electrochemical Society, 150 ͑5͒ E255-E265 ͑2003͒
in a mixture with dichloramine. ͑ii͒ Dichloramine dominates in the
pH range 2.3-4.2. ͑iii͒ Trichloramine is present only in mixture with
dichloramine ͑pH 0.5-2.3͒ or dichloramine and chlorine ͑pH Ϫ0.5-
0.5͒.
Ϫ
ϩ
Ϫ
Ϫ
NH Cl ϩ 2H O ϩ 2e → NH4 ϩ Cl ϩ 2 OH
͓47͔
2
2
Ϫ
Ϫ
II. NH Cl ϩ e → ͓NH •͔ ϩ Cl
rate-determining step ͓48͔
2
2
It was found that all chloramines are electroactive and give well-
formed one- or two-step current-potential waves at the rotating plati-
num and gold electrodes. The overall number of electrons for the
total process is equal 2, 4, and 6 in the case of mono-, di-, and
trichloramine reduction, respectively. The final products of reduction
are ammonium ͑or ammonia͒ and chloride ions. Monochloramine is
reduced in a single two-electron wave, and dichloramine reduction
proceeds in two two-electron steps ͑via monochloramine͒. In more
acidic solutions ͑pH below 4.2͒ the kinetic current due to the proto-
nated dichloramine reduction ͑single four-electron wave͒ is ob-
served, providing the increase of the height of the first step. In
solutions below pH 2.5, only one-step four-electron reduction is
recorded. The trichloramine reduction consists of two steps: first is
two-electron dichloramine formation followed by four-electron
dichloramine reduction. Trichloramine reduction seems to be more
complicated in more acidic solutions. It is probable that the kinetic
current due to the protonated trichloramine reduction has to be taken
into account.
2
͓NH •͔ → NH NH
͓49͔
2
2
2
1
Ϫ
Ϫ
•
͑NH NH ϩ 2H O ϩ 2e → 2NH ϩ 2 OH ͒ ͓50͔
2
2
2
3
2
Ϫ
Ϫ
Ϫ
NH Cl ϩ H O ϩ 2e → NH ϩ Cl ϩ OH
͓51͔
2
2
3
or
Ϫ
ϩ
Ϫ
Ϫ
NH Cl ϩ 2H O ϩ 2e → NH4 ϩ Cl ϩ 2 OH
͓52͔
2
2
The assumption that Reaction 43 ͑or 48͒ is the rate-determining
step is in agreement with ␣n and the slight changes of E1/2 with pH.
Because hydrazine can be oxidized at the platinum electrode ͑at
about Ϫ1.0 V in basic solutions2 ͒, in order to check if the second
mechanism is operative ring-disk experiments were performed.
The disk electrode was kept at a fixed potential ͑taken from the
current plateau or the rising part of the monochloramine reduction
i-E curve͒ while the voltammogram for the ring was recorded over a
wide range of potentials. The collection efficiency for the
Fe(CN)6 /Fe(CN)6 system was determined to be 0.24. The ring-
disk experiments were performed over a wide range of pH, for dif-
ferent disk potentials, different rotation speeds ͑up to 5000 rpm͒,
and temperatures ͑down to Ϫ10°C͒. No oxidation current was ob-
served except the oxide layer formation and ammonia oxidation cur-
rents. Because it is known that hydrazine decomposition is catalyzed
4-26
Chloramine reduction follows an irreversible pathway. The
kinetic parameters for monochloramine → ammonium, di
→ monochloramine, and tri → dichloramine two-electron reduction
are almost the same: the apparent transfer coefficient is close to 0.5,
and the half-wave potential does not change considerably with pH.
The kinetic mechanism with the following rate-determining step was
proposed:
3
Ϫ
4Ϫ
Ϫ
Ϫ
NXCl ϩ e → ͓NX•͔ ϩ Cl
͓54͔
6
by traces of copper or iron ions, ethylenediaminetetraacetic acid
͑
EDTA, final concentration 0.005 M͒ was added to the
where X denotes H , HCl, and Cl for mono-, di-, and trichloram-
2
2
monochloramine solution: the current response remained un-
changed.
From these facts it can be concluded that the first mechanism is
more probable. That mechanism resembles the one proposed by
ine, respectively. In the case of di- and trichloramine a substrate
adsorption step can be involved in the mechanism, because strong
dependence of the half-wave potential on the electrode material is
observed. Dimerization of ͓NH •͔ in monochloramine reduction to
2
1
2
Ward et al. for NHF reduction on a mercury electrode.
2
form hydrazine was excluded by the ring-disk experiments.
Dichloramine reduction mechanism.—The similar behavior ͑the ki-
netic relations, kinetic parameters, and the overall number of elec-
trons͒ of monochloramine → ammonia and dichloramine
Acknowledgments
Helpful discussions with Professor Zbigniew Galus are gratefully
acknowledged. We are also grateful to Dr. Piotr Piela, who read and
corrected the manuscript. Support from the Polish Committee for
Scientific Research through grants 7-T09A-053-21 and 3-T09A-
→
monochloramine reduction ͑waves w1 and w2, respectively͒ sug-
gests that an analogous mechanism accounts for w2 ͑Eq. 43-45͒
with ͓NHCl•͔ as an intermediate. The only difference may be the
substrate adsorption step ͑Eq. 53͒ that is additionally involved, be-
cause in the case of dichloramine a strong dependence of the half-
wave potential on the electrode material of different catalytic prop-
erties was observed.
0
46-18 is also acknowledged.
Warsaw University assisted in meeting the publication costs of this ar-
ticle.
References
NHCl → ͓NHCl ͔
͓53͔
2
2 ads
1
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2
3
͑
two-electron process, similar kinetic dependencies, and parameter
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The behavior of chloramines was studied by discussing the syn-
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1
1
1
1
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4
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