Kinetics and mechanisms of the oxidation of hydrazinium ion (N2H5+) by aqueous Br2, Cl2, and BrCl. Electrophilicity scale for halogens and interhalogens

Article abstract of DOI:10.1021/ic991220k

Very rapid oxidations of N₂H₅⁺ by Br₂, Cl₂, and BrCl are measured by stopped-flow and pulsed-accelerated-flow methods in acidic solutions with excess N₂H₅⁺. Second-order rate constants (M^-1 s^-1) at 25.0 °C, μ = 1.0 M are 1.49 x 10⁷, 1.01 x 10⁸, and 5.6 x 10⁸ for the reactions with Br₂, Cl₂, and BrCl, respectively. The reactions are postulated to proceed by nucleophilic reaction of N₂H₅⁺ with XY electrophiles (XY = Br₂, Cl₂, BrCl) to form XN₂H₄⁺ with Y^- and H⁺ release in the rate-determining step. In the subsequent reactions, we propose that XN₂H₄⁺ eliminates X^- and H⁺ rapidly to form N₂H₃⁺ and diazine, N₂H₂, which is oxidized by a second Br₂, Cl₂, or BrCl to form N₂ in fast steps. The stoichiometries are measured and confirmed to be 1:2 for the Cl₂ and BrCl reactions. The relative reactivities for the oxidation of N₂H₅⁺ by halogens and interhalogens (BrCl > Cl₂ > Br₂ > ICl >> IBr >> I₂) are used to establish an electrophilicity scale (E(XY)) for this type of reaction in aqueous solution.

Full text of DOI:10.1021/ic991220k

1974
Inorg. Chem. 2000, 39, 1974-1978
+
Kinetics and Mechanisms of the Oxidation of Hydrazinium Ion (N2H5 ) by Aqueous Br2,
Cl2, and BrCl. Electrophilicity Scale for Halogens and Interhalogens
Zhongjiang Jia, Murad G. Salaita, and Dale W. Margerum*
Department of Chemistry, Purdue University, West Lafayette, Indiana 47907
ReceiVed October 18, 1999
+
Very rapid oxidations of N2H5 by Br2, Cl2, and BrCl are measured by stopped-flow and pulsed-accelerated-
flow methods in acidic solutions with excess N2H5+. Second-order rate constants (M-1 s-1) at 25.0 °C, µ ) 1.0
M are 1.49 × 107, 1.01 × 108, and 5.6 × 108 for the reactions with Br2, Cl2, and BrCl, respectively. The reactions
are postulated to proceed by nucleophilic reaction of N2H5+ with XY electrophiles (XY ) Br2, Cl2, BrCl) to form
+
XN2H4+ with Y- and H+ release in the rate-determining step. In the subsequent reactions, we propose that XN2H4
eliminates X- and H+ rapidly to form N2H3+ and diazine, N2H2, which is oxidized by a second Br2, Cl2, or BrCl
to form N2 in fast steps. The stoichiometries are measured and confirmed to be 1:2 for the Cl2 and BrCl reactions.
The relative reactivities for the oxidation of N2H5 by halogens and interhalogens (BrCl > Cl2 > Br2 > ICl .
+
IBr . I2) are used to establish an electrophilicity scale (EXY) for this type of reaction in aqueous solution.
Introduction
small amount of hydrazoic acid was formed by the action of
chlorine with hydrazine sulfate in both acid and alkaline
solutions, but none was formed with iodine reaction. They also
detected traces of azide ion (N3-) from the reaction of bromine
in alkaline solution.
Hydrazine is a very reactive reducing agent that is used in a
variety of applications such as rocket propellants, explosives,
boiler feedwater deoxygenation, and pharmaceutical synthesis.1
The oxidation of hydrazine in aqueous solution was recently
reviewed by Stanbury.2 A general mechanism for the oxidation
of hydrazine was given by Higginson3 in 1957 based on Kirk
and Browne’s classification4 of one- and two-electron oxida-
tions. In one-electron oxidations, N2H3 was proposed as an
intermediate that yields a mixture of hydrazoic acid (HN3),
dinitrogen, and ammonia. In two-electron oxidations, N2H2 was
proposed as an intermediate that leads to N2 formation. Iodine
is a two-electron oxidant, and the stoichiometry of I2 with N2H4
has been shown to be 2:1.5 The stoichiometry of the reaction
between Cl2 and N2H4 was often assumed to be 2:1.6 Olson7
determined that the number of gram atoms of Cl consumed per
mole of hydrazine was 4.15 (no precision was given). A
stoichiometric reaction
The kinetics of the reactions of hydrazine with aqueous I2
have been studied by Liu and Margerum.11 They showed that
the reaction between I2 and N2H4 proceeds by the formation of
a complex I2N2H4. This complex eliminates I- to form a steady-
state species IN2H4+, which eliminates I- and H+ to form
N2H3+. The subsequent intermediate, diazine (N2H2), is oxidized
rapidly by a second I2 to form N2. Recently, Jia and Margerum12
reported that the oxidation of hydrazinium ion (N2H5+) by
interhalogens ICl and IBr follows a similar mechanism as the
N2H5+/I2 reaction. Oxidations of N2H5+ by I2, IBr, and ICl are
all proposed to proceed by an I+ transfer process, and the relative
rates are in the order of ICl . IBr . I2.
Choi and Park13 attempted to study the reaction between
hydrazine and Br2 in sulfuric acid solutions by a chronopoten-
tiometric method. No consideration was given to the effects of
N2H4 + 2Cl2 f N2 + 4HCl
(1)
the protonated species of hydrazine, N2H5 and N2H62+, or to
+
the Br3- complex, and the reported rate constants are not valid.
The kinetics and mechanisms of the oxidation of hydrazine by
Cl2 and BrCl have not been previously studied. High acid and
halide ion concentrations are needed to suppress the reaction
rates and to prevent the hydrolysis of halogens and interhalo-
gens. In high acid concentrations, the reactions are suppressed
in eq 1 was suggested by Chuprina and co-workers8 for the
removal of chlorine from hydrochloric acid by the use of
hydrazine chloride. A similar stoichiometric reaction was used
by Kalavska and Rurikova9 for the determination of hydrazine
in its reaction with Br2. Browne and Shetterly10 reported that a
+
2+
(1) Schmidt, E. W. Hydrazine and Its DeriVatiVes; Wiley: New York,
1984; pp 714-856.
because N2H5 is less reactive than N2H4 while N2H6 is not
+
2+
reactive. The protonation constant of N2H5 to form N2H6
,
(2) Stanbury, D. M. In Progress in Inorganic Chemistry; Karlin, K. D.,
Ed.; John Wiley & Sons: New York, 1998; pp 511-561.
(3) Higginson, W. C. E. Chem. Soc., Spec. Publ. 1957, 10, 95-111.
(4) Kirk, R. E.; Browne, A. W. J. Am. Chem. Soc. 1928, 50, 337-347.
(5) Cooper, J. N.; Ramette, R. W. J. Chem. Educ. 1969, 46, 872-873.
(6) Higginson, W. C. E.; Suttton, D.; Wright, P. J. J. Chem. Soc. 1953,
1380-1386.
KP2 ) [N2H62+]/[N2H5+][H+], is 1.72 M-1 at µ ) 1.0 M and
25.0 °C.12 High concentrations of halide ions such as Br- and
Cl- also suppress the rates because Br3-, Cl3-, and BrCl2- are
not reactive. The formation equilibrium constants for these
halogen and interhalogen complexes are 16.1 M-1 for KBr3
)
[Br3-]/[Br2][Br-],14 0.18 M-1 for KCl3 ) [Cl3-]/[Cl2][Cl-],14
(7) Olson, E. C. Anal. Chem. 1960, 32, 1545-1547.
(8) Chuprina, L. F.; Shatalov, B. I.; Levinskii, M. I.; Mantulo, A. P.;
Shmatolokha, D. B. Khim. Tekhnol. (KieV) 1981, 6, 22-24.
(9) Kalavska, D.; Rurikova, D. Acta Fac. Rerum Nat. UniV. Comenianae,
Chim. 1985, 33, 35-42.
(11) Liu, R. M.; Margerum, D. W. Inorg. Chem. 1998, 37, 2531-2537.
(12) Jia, Z.; Margerum, D. W. Inorg. Chem. 1999, 38, 5374-5378.
(13) Choi, Q. W.; Park, B. B. J. Korean Chem. Soc. 1975, 19, 403-407.
(10) Browne, A. W.; Shetterly, F. F. J. Am. Chem. Soc. 1908, 30, 53-63.
10.1021/ic991220k CCC: $19.00 © 2000 American Chemical Society
Published on Web 04/06/2000

+
Oxidation of N2H5
Inorganic Chemistry, Vol. 39, No. 9, 2000 1975
Results and Discussion
and 3.8 M-1 for KBrCl2 ) [BrCl2-]/[BrCl][Cl-]15 at µ ) 1.0 M
and 25.0 °C. In this work, rate constants for oxidation of
hydrazine by Br2, Cl2, and BrCl are measured in high acid and
high halide ion concentrations and the corresponding reaction
mechanisms are proposed.
Reaction Stoichiometries. The stoichiometry of the reaction
+
between Cl2 and N2H5 was determined by mixing a known
concentration of excess hydrazine in 1.04 M HCl with known
concentrations of hypochlorite solution. The PAF instrument
was used as a precise and rapid mixer. This method overcomes
the difficulties in quantification and handling of the Cl2
solutions. When excess hydrazine in 1.04 M HCl is mixed with
Experimental Section
Reagents. Experimental details for the preparation and standardiza-
tion of solutions and “bromide-free” hydrochloric acid are given in
previous work.12 The ionic strength (µ) was adjusted with aqueous
NaClO4 that was recrystallized from water. Stock solution of hypochlo-
rite was prepared by bubbling ultrahigh-purity Cl2 gas (Matheson,
99.9%) into 0.2 M NaOH solution and stored at 5 °C in a Nalgene
bottle. The hypochlorite stock solution was standardized spectropho-
tometrically at 292 nm (ꢀOCl ) 362 M-1 cm-1).16 Chlorine solutions
were prepared by T-mixing hypochlorite solution and “Br--free”
hydrochloric acid and used immediately for kinetic studies by transfer
from syringes. The bromine solution was prepared by adding liquid
Br2 dropwise to a 0.8 M HClO4 solution to prevent hydrolysis. The
concentration of Br2 was standardized spectrophotometrically at 390
hypochlorite solution, Cl2 is generated very quickly from the
reaction20 of HOCl and HCl and reacts rapidly with N2H5
.
+
Our preliminary results show that the second-order rate constant
+
for the reaction of N2H5 and HOCl (measured at pH 3.07) is
2.66 × 103 M-1 s-1, which is much slower than the corre-
sponding Cl2 reaction. Three reactions with different concentra-
tions and molar ratios of chlorine and hydrazine were performed
on PAF. Approximately 100 mL of solution from 10 pushes
was collected for each reaction. The stoichiometry of the
reaction between BrCl and N2H5+ was determined by mixing a
known concentration of excess hydrazine with a known total
nm (ꢀBr2 ) 175 M-1 cm-1).14 Solutions of BrCl2 were prepared by
-
acidifying a stoichiometric mixture of NaBr and KBrO3 with “Br--
free” HCl:14
-
concentration of BrCl solution in ∼6 M HCl, where BrCl2
predominates.15 Data for four reactions with different concentra-
tions and molar ratios of BrCl and hydrazine were obtained.
The excess hydrazine in the reaction mixture was back-titrated
in ∼5 M HCl with 0.01 M KIO3 standard solution. The average
stoichiometries are 2.01 ( 0.03 for the Cl2/N2H5+ reaction and
BrO3- + 2Br- + 6H+ + 6Cl- f 3BrCl2- + 3H2O
(2)
The stock solutions of hydrazine were prepared from either N2H4‚HCl
or N2H4‚H2SO4 solids and standardized (in ∼5 M HCl) with 0.025 M
KIO3 solution.17
+
2.04 ( 0.02 for the BrCl/N2H5 reaction. The results are
summarized in Table 1. The precision of our results is no better
than 1-2%. Hence, trace levels of other products could be
formed. However, if appreciable amounts of N3- are formed, it
would lead to XY/N2H5+ ratios of less than 2.0, and this is not
Methodology and Instrumentation. Kinetic studies for the reaction
of hydrazine and Br2 were performed on a Applied PhotoPhysics
stopped-flow spectrophotometer (APPSF) SX.18MV (optical path
length ) 0.962 cm), which was controlled from a dedicated Acorn
RISC PC (RISC OS 3, version 3.60) with APP software (version 4.33).
Kinetic studies for the oxidation of hydrazine by Cl2 and BrCl were
performed on a pulsed-accelerated-flow (PAF) spectrophotometer
(model IV).18 The PAF spectrophotometer uses integrating observation
during continuous flow mixing of short duration (0.4 s pulse, optical
path length ) 2.050 cm). Pseudo-first-order rate constants greater than
105 s-1 are resolved from the physical mixing processes by variation
of flow velocities under turbulent flow conditions.18 Reactions were
-
the case. In the Br2/N2H4 reaction, N3 was detected only in
base,10 whereas we have strong acid. The Br2/N2H5 reaction
+
should follow the same stoichiometry as the corresponding Cl2
and BrCl reactions. The general stoichiometric reaction is
described:
N2H5+ + 2XY f N2 + 5H+ + 2X- + 2Y-
(4)
followed by the loss of Br3 at 266 nm (ꢀBr3 ) 40 900 M-1 cm-1),14
-
+
Reaction of N2H5 and Br2. With excess total hydrazine
([N2H5+]T ) [N2H5+] + [N2H62+]) and high concentrations of
Br- and H+, the loss of total bromine ([Br2]T ) [Br2] + [Br3-]),
observed at 266 nm, follows first-order kinetics. The observed
first-order rate constants (measured on the APPSF after the
correction for mixing) are first-order in [N2H5+]T as shown in
Figure 1. The observed first-order rate constants decrease with
increasing [H+] and [Br-] (Figure 2) because of the lack of
Cl3- at 230 nm (ꢀCl3 ) 9400 M-1 cm-1),14 or BrCl2- at 232 nm (ꢀBrCl2
) 32 700 M-1 cm-1 14 with excess total hydrazine concentration under
)
pseudo-first-order conditions. The observed pseudo-first-order rate
constant is defined in
-d[XY]
T ) kobsd[XY]T
(3)
dt
2+
reactivity of N2H6 and Br3-. The data correspond to
where XY ) Br2, Cl2, and BrCl, and [XY]T ) [XY] + [XY2-]. Each
first-order rate constant, kobsd, measured on the APPSF or PAF is an
average of five trials. The kobsd values from the APPSF were corrected
for mixing limitations of this instrument by using kcorr ) kobsd/[1 -
+
2kBr2[N2H5 ]T
(1 + KBr3[Br-])(1 + KP2[H+])
kobsd
)
(5)
(kobsd/kmix)], where kmix ) 4.62 × 103 s-1 19
. All reactions were run at
25.0 ( 0.1 °C and µ ) 1.0 M. Spectrophotometric measurements were
performed on the Perkin-Elmer Lambda-9 UV/vis/NIR spectropho-
tometer interfaced to a Zenith 386/20 computer with solutions ther-
mostated to 25.0 ( 0.1 °C.
where the (1 + KBr3[Br-]) term corrects for Br3- formation and
the (1 + KP2[H+]) term corrects for N2H62+ formation. The rate
suppressions by H+ and Br- correspond only to the formation
of N2H6 and Br3 without any additional H+ and Br-
suppression (unlike the dependence found for IBr reactions12).
The resolved second-order rate constant (kBr2) for the reaction
2+
-
(14) Wang, T. X.; Kelley, M. D.; Cooper, J. N.; Beckwith, R. C.; Margerum,
D. W. Inorg. Chem. 1994, 33, 5872-5878.
(15) Liu, Q.; Margerum, D. W. Unpublished results.
(16) Furman, C. S.; Margerum, D. W. Inorg. Chem. 1998, 37, 4321-4327.
(17) Jeffrey, G. H.; Bassett, J.; Mendham, J.; Denney, R. C. Vogel’s
Textbook of QuantitatiVe Chemical Analysis, 5th ed.; Wiley & Sons:
New York, 1989; p 402.
(18) Bowers, C. P.; Fogelman, K. D.; Nagy, J. C.; Ridley, T. Y.; Wang,
Y. L.; Evetts, S. W.; Margerum, D. W. Anal. Chem. 1997, 69, 431-
438.
+
of N2H5 and Br2 is calculated from the slope in Figure 1 in
accord with eq 5, which gives a value of (1.49 ( 0.02) × 107
M-1 s-1 at 25.0 °C and µ ) 1.0 M.
Reaction of N2H5+ and Cl2. The oxidation of hydrazine by
aqueous chlorine is very fast, and rate suppression due to the
(19) Baron, C. D.; Margerum, D. W. Unpublished results.
(20) Wang, T. X.; Margerum, D. W. Inorg. Chem. 1994, 33, 1050-1055.

1976 Inorganic Chemistry, Vol. 39, No. 9, 2000
Jia et al.
+
Table 2. Reactions of N2H5 and Cl2 under Second-Order Unequal
Concentration Conditions on the PAFa
+
104[N2H5
M
]
T
104[Cl2]T
M
10-7kr
[N2H5+]T/[Cl2]T
M-1 s-1
2.08
2.77
4.16
5.54
5.54
6.93
8.31
8.31
1.45
2.00
2.56
3.12
1.45
2.00
3.12
2.56
1.43
1.39
1.63
1.78
3.82
3.47
2.66
3.25
6.3 ( 0.1
7.1 ( 0.1
6.11 ( 0.08
6.9 ( 0.2
5.9 ( 0.4
6.5 ( 0.2
7.2 ( 0.2
7.5 ( 0.4
ave: 6.9 ( 0.5
a Conditions: [HCl] ) 1.00 M, 25.0 °C, λ ) 230 nm.
Figure 1. Dependence of the observed first-order rate constants on
[N2H5+]T for the oxidation of N2H5+ by Br2 at 25.0 °C and µ ) 1.0 M
on the APPSF. Conditions are the following: [Br2]T ) (0.37 to 1.9) ×
10-5 M, [H+] ) 0.246 M, [Br-] ) 0.7716 M, λ ) 266 nm, slope )
Bromide ion can interfere with the reaction of chlorine and
- 13
hydrazine by forming BrCl and BrCl2
,
Cl2 + Br- a BrCl2
(7)
(8)
-
(1.56 ( 0.02) × 106 M-1 s-1
.
BrCl + Cl- {KBrCl2} BrCl2
-
and BrCl is very reactive toward N2H5+. The formation
-
equilibrium constant of BrCl2 from the reaction of Cl2 and
Br- (eq 7) is calculated to be 1.2 × 107 M-1 by using related
equilibrium constants14 and E° values of Cl2/Cl- and Br2/Br-.
Most of the Br- (1.0 × 10-5 M) in 1.0 M HCl (without
purification) will be converted by Cl2 to BrCl and BrCl2-, and
its concentration can be as much as 10% of the chlorine solution
that is used. Bromide formed from the reaction of BrCl and
+
N2H5 will regenerate BrCl from its reaction with Cl2 (eq 7).
To obtain accurate kinetic measurements, “Br--free” HCl is used
+
for the reaction of N2H5 and Cl2.
Figure 2. Observed first-order rate constants as a function of [H+]
(O) and [Br-] (b) for the oxidation of N2H5+ by Br2 at 25.0 °C and µ
Rate constants were measured (Table 2) by the PAF technique
for a series of reactions under second-order unequal concentra-
tion conditions with different [N2H5+]T/[Cl2]T molar ratios in
1.00 M HCl. The resolved second-order rate constant, kCl2, is
calculated from
) 1.0 M on the APPSF: (O) [N2H5
]
T
) 7.63 × 10-5 M, [Br2]T
)
+
1.20 × 10-5 M, [Br-] ) 0.3585 M, λ ) 266 nm; (b) [N2H5+]T ) 7.63
× 10-5 M, [Br2]T ) 1.14 × 10-5 M, [H+] ) 0.125 M, λ )
266 nm.
+
Table 1. Measurements of Stoichiometry for the Cl2/N2H5 and the
BrCl/N2H5 Reactions
kr(1 + KCl3[Cl-])(1 + KP2[H+])
+
kCl2
)
(9)
stoichiometry
2
+
[N2H5
M
]
[Cl2]T or [N2H5+]T M,
[BrCl]T M unreacted
Cl2 (or
T
+
reaction
BrCl)/N2H5
in which kr is the observed second-order rate constant, KCl3
)
0.18 M-1, KP2 ) 1.72 M-1, and [HCl] ) 1.00 M. For the
+
Cl2 + N2H5
0.009 94 0.005 49
0.014 91 0.010 98
0.019 88 0.016 47
0.007 16
0.009 54
0.011 70
1.97 ( 0.02
2.04 ( 0.01
+
reaction of N2H5 and Cl2, the average second-order rate
constant kCl2 is (1.11 ( 0.08) × 108 M-1 s-1 at 25.0 °C and µ
2.01 ( 0.01
ave: 2.01 ( 0.03
2.06 ( 0.01
) 1.00 M.
+
BrCl + N2H5 0.009 85 0.011 25
0.004 535
+
The reactions of N2H5 and Cl2 under pseudo-first-order
0.012 48 0.009 375 0.007 838
0.017 11 0.010 71 0.011 88
0.018 63 0.007 000 0.015 21
2.02 ( 0.01
conditions were studied with increasing total hydrazine con-
centrations in 1.00 M HCl. The observed pseudo-first-order rate
constants are first-order in total hydrazine concentrations (Figure
3), where the slope kr is (6.3 ( 0.2) × 107 M-1 s-1 for the
2.05 ( 0.01
2.05 ( 0.01
ave: 2.04 ( 0.02
-
+
formation of Cl3 is small because KCl3 is only 0.18 M-1
:
reaction of N2H5 and Cl2. The resolved second-order rate
constant kCl2 is calculated to be (1.01 ( 0.03) × 108 M-1 s-1
at 25.0 °C and µ ) 1.00 M. This second-order rate constant
(obtained under the pseudo-first-order conditions) is in good
agreement with the rate constant obtained under second-order
unequal concentration conditions.
KCl3
Cl2 + Cl- { } Cl3
(6)
-
Because of the small fraction of Cl3-, high concentrations of
total chlorine solution are needed to reach reasonable absor-
bance. The reactions were studied by the PAF technique under
both second-order unequal concentration conditions and pseudo-
first-order conditions. In the second-order unequal concentration
reactions, the initial total concentration of chlorine was calcu-
lated from the initial absorbance of chlorine solution and the
calibration curve on PAF under the same conditions.
Reaction of N2H5+ and BrCl. The plot of the observed first-
order rate constants, kobsd, vs [N2H5+]T is linear as shown in
Figure 3. The observed second-order rate constant kr is the slope
in Figure 3, which is (8.6 ( 0.8) × 107 M-1 s-1 for the reaction
of N2H5+ and BrCl in 1.00 M HCl. The resolved second-order
rate constant, kBrCl, is calculated to be (5.6 ( 0.5) × 108 M-1
s-1 at 25.0 °C and µ ) 1.00 M according to eq 9, where KBrCl2

+
Oxidation of N2H5
Inorganic Chemistry, Vol. 39, No. 9, 2000 1977
+
Table 3. Rate Constants (kXY) of Oxidation of N2H5 by XY at
25.0 °C and Electrophilicities (EXY
)
kXY
EXY
EXY
XY
I2
IBr
ICl
Br2
Cl2
BrCl
ClF
M-1 s-1
aqueous-phased
gas-phasee
0.70a
0
7.6 × 104 b
4.12 × 106 b
1.49 × 107 c
1.01 × 108 c
5.6 × 108 c
5.0
6.8
7.3
8.2
8.9
7.4
5.1
9.0
9.8
a Reference 11, µ ) 0.5 M. b Reference 12, µ ) 1.0 M. c This work,
µ ) 1.0 M. d This work. e Reference 25.
rapidly by a second XY to form N2 in multiple steps (eq 16).
Therefore, the stoichiometric ratio value of 2 in eqs 5, 9, and
10 is taken into account for the consumption of the second XY.
Figure 3. Dependence of the observed first-order rate constants on
+
[N2H5+]T for the oxidation of N2H5 by Cl2 (4) and BrCl (9) at 25.0
°C and µ ) 1.00 M on the PAF: (4) 10-5kobsd vs 103[N2H5+]T, [Cl2]T
KP2
N2H5+ + H+ { } N2H6
(fast)
(fast)
(11)
(12)
2+
) (0.51 to 3.30) × 10-4 M, [HCl] ) 1.00 M, λ ) 230 nm, slope )
(6.3 ( 0.2) × 107 M-1 s-1; (9) 10-4kobsd vs 104[N2H5+]T, [BrCl]T
)
(0.469 to 2.344) × 10-5 M, [HCl] ) 1.00 M, λ ) 232 nm, slope )
KXY2
-
XY + Y- { } XY2
(8.6 ( 0.8) × 107 M-1 s-1
.
+
Scheme 1. Transition State for the Reactions of N2H5 and
XY in Eq 13, Where XY ) Br2, Cl2, and BrCl
kXY
N2H5+ + XY 8 XN2H4+ + Y- + H+
(rds) (13)
XN2H4+ f N2H3+ + X- + H+
(fast)
(fast)
(14)
N2H3+ a N2H2 + H+
N2H2 + XY f N2 + 2H+ + X- + Y-
(15)
) 3.8 M-1 is used instead of KCl3. The observed first-order rate
constants change as a function of [H+] (from 0.423 to 0.844
M) and [Cl-] (from 0.462 to 0.928 M), consistent with
(fast) (16)
+
2kBrCl[N2H5 ]T
Relative Reactivities for the Oxidation of N2H5+. Second-
order rate constants for the oxidation of N2H5 by halogens
kobsd
)
(10)
+
(1 + KBrCl2[Cl-])(1 + KP2[H+])
and interhalogens are summarized in Table 3. The relative
reactivities for the oxidation of N2H5+ are in the order of BrCl
> Cl2 > Br2 > ICl . IBr . I2. The reactions occur by X+ (X
) I, Br, Cl) transfer rather than electron transfer. A linear free
energy relationship was reported12 for log kIX (X ) I, Br, Cl)
vs log KC, where KC values are the complex formation constants
of IX with pyridine and 3-methylpyridine. Halogens and
interhalogens can be classified as Lewis acids with decreasing
acid strength as ICl . BrCl > IBr . I2 > Br2 . Cl2 by the
comparison of the free energies of trihalide formation.22 This
is clearly not the dependence we observe. The reactivities of
halogens and interhalogens with nucleophiles depend on their
electrophilicity. In the present case the energetics of N-X bond
formation and Y- loss are both important. The contribution from
Lewis acidity is less important, although it may account for the
greater reactivity of BrCl compared to Cl2.
The decrease of the first-order rate constants with increasing
2+
-
[H+] and [Cl-] is due to the formation of N2H6 and BrCl2
(eq 8). The terms (1 + KP2[H+]) and (1 + KBrCl2[Cl-]) in eq 10
are used for that correction.
General Mechanism for the Oxidation of N2H5 by XY.
+
From the kinetic studies, a general mechanism for the oxidation
of N2H5+ by XY is proposed in eqs 11-16, where XY ) Br2,
+
Cl2, and BrCl. The N2H5 ion acts as a weak nucleophile in
the reaction with XY electrophiles to form a reactive intermedi-
ate, XN2H4+, in the rate-determining step (eq 13). The transition
state for reaction 13 is shown in Scheme 1, where the reaction
proceeds by concerted X-Y bond cleavage and X-N bond
formation with X+ transfer as a net result. An electron-transfer
mechanism has been ruled out for the oxidation of hydrazine
by I2, where the calculated rate constants are 9 orders of
magnitude smaller than the measured ones.11 In the oxidation
From the kinetic studies of halogen and interhalogen addition
to olefinic compounds in acetic acid solutions, White and
Robertson23 estimated the following relative reactivities, I2 (1),
IBr (3 × 103), ICl (105), Br2 (104), and BrCl (4 × 106), which
has some similarities to our results except for the order of Br2
and ICl and the smaller range of reactivities. Our values of the
of N2H5 by ICl and IBr,12 the first step to form IN2H4 is
reversible. However, for Br2, Cl2, and BrCl the reverse reaction
in eq 13 is not appreciable, and hence, no additional H+, Br-,
or Cl- suppressions are observed. The reactive intermediates
+
+
+
+
+
BrN2H4 and ClN2H4 are more acidic than IN2H4 because
bromine and chlorine are more electronegative than iodine,
which tends to make eq 13 irreversible. In the subsequent rapid
+
second-order rate constants (kXY) for the XY/N2H5 reactions
range over 9 orders of magnitude and involve Y- release in
aqueous solutions. It is possible to establish an electrophilicity
scale for halogens and interhalogens (EXY) according to the
+
+
reactions, XN2H4 can eliminate X- and H+ to form N2H3
(eq 14), which is in rapid equilibrium with diazine, N2H2 (eq
+
15). The dissociation constant of N2H3 to form N2H2 (eq 15)
(21) McKee, M. L. J. Phys. Chem. 1993, 97, 13608-13614.
(22) Scott, R. L. J. Am. Chem. Soc. 1953, 75, 1550-1552.
(23) White, E. P.; Robertson, P. W. J. Chem. Soc. 1939, 1509-1515.
was calculated to be 32 M.21 We propose diazine as an
intermediate, but it is not detected. Finally, diazine is oxidized

1978 Inorganic Chemistry, Vol. 39, No. 9, 2000
Swain-Scott free energy relationship:24
Jia et al.
Scheme 2. Proposed Mechanism for the Base-Catalyzed
Reaction of HNCl2 and NCl3
log(kXY/kI ) ) s′EXY
(17)
2
+
We have selected the second-order rate constant of the I2/N2H5
reaction (kI2) as a reference and have arbitrarily assigned a value
of unity for the sensitivity factor, s′. The resulting electrophi-
licities of halogens and interhalogens are listed in Table 3.
Legon25 established a limiting gas-phase electrophilicity scale
for halogens and interhalogens based on the intermolecular
stretching force constants (kσ) determined from the rotational
spectra of prereactive complexes [B‚‚‚XY]. An empirical
equation kσ ) cNE was used for calculation, where c ) 0.25 N
m-1, E (electrophilicity) ) 10 for HF, and N (nucleophilicity)
) 10 for H2O.26 Legon’s gas-phase electrophilicity scale of XY
is a measure of their capacity to interact with a nucleophile B
without atom transfer. However, for our reactions in aqueous
is OH-, the third-order rate constant is 108 M-2 s-1 for the rate
expression28
d[N2]
) kB[B][HNCl2][NCl3]
(18)
+
solution, the halogen cation X+ transfers to N2H5 as Y- and
dt
H+ are released. Despite these differences in the nature of the
reactions, the gas-phase and aqueous-phase electrophilicities are
similar for Br2 and BrCl but not for Cl2. The arbitrary choice
of kI2 as a reference and s′ as unity automatically gives an EXY
value of zero for I2, but this is for comparison purposes only
and does not mean that it has zero electrophilicity. The fact
The decomposition of the proposed intermediates tetrachloro-
hydrazine and dichlorodiazine are rapid and occur after the rate-
determining step, in which N-N bond formation takes place.
A question that arises is why the HNCl2/NCl3 reaction pair
carries the bulk of the N2 generation process. Our present work
shows that monochlorohydrazine (eq 14) reacts very rapidly to
give diazine, which in turn is rapidly oxidized to N2. Thus, the
replacement of one H by Cl in hydrazine would be sufficient
to rapidly generate N2. In water treatment processes, there are
many other reactant pairs such as NH2Cl/NH2Cl, NHCl2/NHCl2,
NH3/NCl3, NH2Cl/NCl3, etc. that might form N-N bonds, where
the resulting chlorohydrazine species would rapidly generate
N2. The fact that eq 18 is the preferred path indicates the
importance of generating a strong nucleophile (NCl2-) that can
react with the electrophile (NCl3) to form an N-N bond and
displace Cl-. The present results also suggest that some of the
other reactant pairs could contribute to the breakpoint chlorina-
tion process as conditions are altered.
+
that N2H5 is a very weak nucleophile permits comparison of
the relative rates of the XY electrophiles. A strong nucleophile
such as N2H4 would react at diffusion-controlled rates with all
the XY electrophiles except I2, where its rate constant is 2.4 ×
107 times larger than with N2H5+. Our electrophilicity scale of
halogens and interhalogens could be used for prediction of the
rate constants in their reactions with other weak nucleophiles.
Halohydrazines and Breakpoint Chlorination. In water
treatment plants throughout the world, a process known as
breakpoint chlorination27 is used to destroy chloramines and
generate dinitrogen. The base-catalyzed reaction of dichloramine
with trichloramine is a key pathway for this process, as shown
by the proposed mechanism in Scheme 2.28 When the base (B)
Acknowledgment. This work was supported by National
Science Foundation Grant CHE-96-22683 and by a grant from
the Purdue Research Foundation.
(24) Swain, C. G.; Scott, C. B. J. Am. Chem. Soc. 1953, 75, 141-147.
(25) Legon, A. C. Chem. Commun. 1998, 2585-2586.
(26) Legon, A. C.; Millen, D. J. J. Am. Chem. Soc. 1987, 109, 356-358.
(27) Johnson, D. L. Water Chlorination; Jolley, R. L., Ed.; Ann Arbor
Science: Ann Arbor, MI, 1978; Vol. 1, p 44.
Supporting Information Available: Listings of kinetic data. This
material is available free of charge via the Internet at http://pubs.acs.org.
(28) Yinn, B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 2135-2141.
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