Spectroscopy of hydrothermal reactions. 7. Kinetics of aqueous [NH3OH]NO3 at 463-523 K and 27.5 MPa by infrared spectroscopy

Article abstract of DOI:10.1021/jp9723877

The first kinetic measurements for aqueous [NH₃OH]NO₃ that are based on species concentrations at hydrothermal conditions are discussed. The decomposition rate of 0.1 and 0.2 m [NH₃OH]NO₃ was determined in real time by flow-cell IR spectroscopy at 463-523 K and 27.5 MPa. The loss of NH₃OH⁺ (2730 cm^-1 bending combination) and formation of N₂O (v₃ at 2230 cm^-1) occur at the same rate under these conditions and were fit to a first-order rate expression. Using all of the data, E_a = 103 ± 21 kJ/mol and ln(A/s) = 21 ± 5. The decomposition rate resembles that of solid [NH₃OH]NO₃ in a similar temperature range, but the kinetics for dilute solutions and the solid state are too slow to apply to the combustion rate of a more concentrated solution of [NH₃OH]NO₃.

Full text of DOI:10.1021/jp9723877

J. Phys. Chem. A 1997, 101, 8593-8596
8593
Spectroscopy of Hydrothermal Reactions. 7. Kinetics of Aqueous [NH3OH]NO3 at
4
63-523 K and 27.5 MPa by Infrared Spectroscopy
J. W. Schoppelrei and T. B. Brill*
Department of Chemistry, UniVersity of Delaware, Newark, Delaware 19716
X
ReceiVed: July 23, 1997; In Final Form: September 17, 1997
The first kinetic measurements for aqueous [NH
hydrothermal conditions are discussed. The decomposition rate of 0.1 and 0.2 m [NH
in real time by flow-cell IR spectroscopy at 463-523 K and 27.5 MPa. The loss of NH
3
OH]NO
3
that are based on species concentrations at
OH]NO was determined
OH (2730 cm
at 2230 cm ) occur at the same rate under these conditions
and were fit to a first-order rate expression. Using all of the data, E ) 103 ( 21 kJ/mol and ln(A/s) ) 21
5. The decomposition rate resembles that of solid [NH OH]NO in a similar temperature range, but the
kinetics for dilute solutions and the solid state are too slow to apply to the combustion rate of a more
concentrated solution of [NH OH]NO
3
3
+
-1
3
-
1
bending combination) and formation of N
2
O (ν
3
a
(
3
3
3
3
.
Introduction
primarily due to the stress of temperature cycling and required
occasional disassembly and resealing. The reaction zone of the
cell had an internal volume of 0.0604 mL as determined by the
cylindrical channel machined through the cell body and the flat
duct between the sapphire windows. A transport model of the
cell indicated laminar flow and that the solution reaches 85%
of the cell temperature well within the first 15% of the entrance
tube. The cell/reaction temperature was determined by a
K-type thermocouple located in the cell body near the flat duct
between the windows.
A highly concentrated aqueous solution of hydroxylammo-
nium nitrate, [NH3OH]NO3, mixed with alkylammonium nitrate
salts burns very rapidly, which makes it useful as a liquid
propellant.1
-3
In fact, even approximately 40 wt % of [NH3-
OH]NO3 in water (5 M) will burn steadily provided that the
pressure is greater than 5 MPa. A phenomenological model
describing this deflagration process has been developed. The
controlling event is believed to be the rate of decomposition of
4
9
5
[NH3OH]NO3 in the condensed phase. Therefore, the kinetics
Experiments were conducted on two concentrations of [NH3-
OH]NO3 at 27.5 MPa and with flow rates of 1.79-0.03 mL/
min, which resulted in residence times of 2-110 s. A 0.20 m
and pathway of decomposition in water at high temperature and
pressure are of interest. The main problem is, however, that
the reaction is highly exothermic, which makes it difficult to
control during the investigation. Thus, although [NH3OH]NO3
dissolves up to 95 wt % in water, the use of a dilute solution
has been necessary to measure the reaction rate.
[
NH3OH]NO3 solution was studied at four temperatures between
63 and 493 K and 0.10 m [NH3OH]NO3 was studied at seven
temperatures between 463 and 523 K. IR spectra (32 summed
4
6
-1
interferograms, 4 cm resolution) were collected with a Nicolet
0 SX FTIR spectrometer using 17 different flow rates at each
Recent development of spectroscopy/flow cells suitable for
real-time IR and Raman studies at hydrothermal conditions (T
6
temperature. After each spectral set was collected and the flow
rate was adjusted, the system was stabilized and flushed with
at least 10 reactor volumes of sample at the new conditions
prior to collection of the next spectral set. Collection of data
at each set of flow rates was repeated three times before moving
to the next temperature. The cell required resealing after the
data sets at the lowest four temperatures were collected on 0.10
m [NH3OH]NO3. A replicate data set was collected at 493 K
with the refurbished cell for comparison with the data collected
earlier.
)
450-620 K, P ) 27.5 MPa) make direct spectrokinetic
7-9
measurements now feasible.
Described herein is a real-time
investigation of the decomposition rate of 0.1 and 0.2 m [NH3-
OH]NO3 (m ) mol/kg of H2O) at 463-523 K and 27.5 MPa.
Experimental Section
Within 2 h of use, solutions having 1:1 NH2OH:HNO3 were
prepared from a stock solution of approximately 50 wt % (16.5
M) hydroxylamine (Howard Hall Division of R.W. Greef and
Co.), 70 wt % (15.8 M) nitric acid (Fisher Scientific, ACS),
and HPLC-grade water and were stored in a glass container.
Results
7
Reduction of the spectra to kinetic data required several steps.
First, each reaction spectrum was ratioed against the spectrum
of pure H2O in the same cell at the same pressure and
temperature to remove as much of the absorption by water as
possible. Second, the reactant and product bands were resolved,
and areas were calculated by using curve-fitting software
The flow reactor used permits excellent control of the pressure
((0.1 MPa), temperature ((1.5 K), and flow ((0.01 mL/min,
pulseless). The data were collected with a PC that was
programmed to control the flow reactor automatically. Such
automation permitted multiple data sets to be collected which
revealed their overall consistency.
(Peakfit, Jandel). The absorbances of interest were fitted with
The reaction takes place inside the heated spectroscopy cell
constructed of 316 stainless steel (SS) and fitted with sapphire
a four-parameter Voigt function. This method permitted ac-
curate integration of the weaker absorptions with the rather poor
baselines frequently encountered in IR spectra of water solutions.
Third, the absorbances were converted into concentrations at
each condition so that the reaction rate could be obtained.
Two distinct absorptions of [NH3OH]NO3 were observed in
7
windows. The cell was sealed by compression against a set
of gold gaskets whose thickness also controlled the path length
(0.0030-0.0045 mm). These seals have limited durability
*
Corresponding author.
Abstract published in AdVance ACS Abstracts, October 15, 1997.
X
-1
the 1800-2900 cm window available with sapphire and water.
S1089-5639(97)02387-6 CCC: $14.00 © 1997 American Chemical Society

8594 J. Phys. Chem. A, Vol. 101, No. 46, 1997
Schoppelrei and Brill
Figure 3. Example of first-order rate plots for the conversion of [NH
OH]NO under 27.5 MPa from the N O data at the temperatures shown
see text). Error bars have been omitted for clarity.
3
-
3
2
Figure 1. Example infrared spectra of 0.1 m [NH
and 27.5 MPa showing the conversion of the NH
bending combination) into N O(aq) as a function of the residence time.
3 3
OH]NO at 493 K
(
+
+
3
3
OH ion (-NH
2
OH+
3
TABLE 1: First-Order Rate Constants for NH
Decomposition and N
2
O Formation from Aqueous
[NH
3
OH]NO
3
k, s- (×10 )
1
3
expta
T, K
NH
+
bend combination
3
3 2
ν (N O)
1
2
1
2
1
2
1
2
3
3
3
3
463
2.5 ( 0.2
3.6 ( 0.5
4.3 ( 0.3
4.0 ( 0.5
11.2 ( 0.5
13.3 ( 1.6
11.2 ( 1.3
18.5 ( 1.5
13.3 ( 2.7
24.6 ( 2.3
26.5 ( 3.3
58 ( 43
1.6 ( 0.4
1.7 ( 0.2
3.1 ( 0.1
4.4 ( 0.3
6.1 ( 0.4
5.8 ( 1.1
7.7 ( 0.3
13.7 ( 0.6
15.5 ( 0.8
17.4 ( 3.5
32.0 ( 5.1
43.5 ( 4.5
473
483
493
503
513
523
a
1
: [[NH
3
OH]NO
3
]
0
) 0.20 m; initial seal of cell. 2: [[NH
3 3 0
OH]NO ]
OH]NO ) 0.10 m; second
)
0.10 m; initial seal of cell. 3: [[NH
seal of cell.
3
3 0
]
Figure 2. Concentrations of NH
residence time from 0.1 m [NH OH]NO
3
OH+ and N
at 493 K under 27.5 MPa.
2
O as a function of
3
3
Points represent the average of three data sets, and error bars were
determined from the sample standard deviation.
The collection of multiple data sets permitted determination
of the uncertainties in the concentration-time data. Attention
was given to the propagation of these uncertainties throughout
the kinetic analysis. Each group of three flow rate sets at a
particular temperature essentially represented a separate kinetic
experiment. Thus, the average concentration ( xj ) at each flow
rate was determined from the three individual measurements,
and the uncertainty was determined by the sample standard
deviation (s). The averaged concentration data and resulting
uncertainties were then plotted against the residence time which
was calculated by dividing the reactor volume given above by
the flow rate. The reaction profiles result, such as are shown
in Figure 2. The flow rate was corrected for the difference in
the density of water at each temperature.
-
1
The weak, broad absorption at about 2730 cm resulted from
+
the -NH3 bending combination mode of the hydroxylammo-
1
0
-1
nium ion reactant. A stronger absorption at about 2230 cm
is the ν3 stretch of the aqueous N2O product (Figure 1). The
concentrations represented by the integrated band areas were
calibrated internally at each temperature.11 Over most of the
temperature range used, the spectra at the shortest residence
times indicated no significant reaction and provided an absor-
+
bance value for the initial concentration of NH3OH . The
absorbance intensity was converted to concentration by assuming
Beer’s law. In the highest two temperature sets, some decom-
position was evident even at the shortest residence times and
was taken into account in the analysis. The N2O concentration
The N2O concentration was converted to conversion of [NH3-
OH]NO3 by [[NH3OH]NO3]0 - ([N2O]t/0.8), and the kinetic
analysis was performed separately on the reactant and the
product data using both zeroth- and first-order rate equations.
Zeroth-order plots of concentration vs time and first-order plots
of ln(concentration) vs time were constructed up to ap-
proximately 60% conversion. The first-order plots are shown
in Figure 3. The rate parameters were determined from the slope
of weighted-least-squares regression (WLSQR) of the averaged
concentration data weighted by their uncertainties. In the first-
order kinetic analysis, uncertainties were translated into log
-
1
was calibrated similarly from the 2230 cm band by the fact
that complete reaction was achieved at most temperatures in
the longest residence time spectra as indicated by the unchanging
area of the N2O band and the absence of the reactant band.
Complete reaction was not achieved at the lowest three
temperatures, so the area of the N2O band at completion at a
higher temperature was adjusted to account for the density
difference of water and then was set equal to the band area at
lower temperature. At completion the N2O concentration is 0.8
of the initial [NH3OH]NO3 concentration based on the stoichi-
ometry determined by Raman spectroscopy at hydrothermal
conditions.8
1
2
space as s/ xj . The values of k and the resulting uncertainties
are listed in Table 1.

Spectroscopy of Hydrothermal Reactions. 7
J. Phys. Chem. A, Vol. 101, No. 46, 1997 8595
+
two variables, the reactant NH3OH and product N2O, are
-
1
observable in the 1800-2900 cm spectral window available
(
Figure 1). On the other hand, a 1:1 ratio of NH2OH:HNO3
1
-3
has the widest practical application.
A pathway of the decomposition of [NH3OH]NO3 in water
1
7
leading to (1) is given by reactions 5-8.
k1
+
-
NH OH + NO { } NH OH + HNO
(5)
3
3
2
3
k-1
k2
NH OH + HNO
9
8 HNO + H O + HNO
(6)
(7)
(8)
2
3
2
2
k3
2
HNO
98 N O + H O
2 2
k4
2
(NH OH + HNO
98 N O + 2H O)
2 2
2
2
Figure 4. An Arrhenius plot of all of the rate constants in Table 1.
Steady-state treatment of HNO, HNO2, and NH2OH yields the
rate equation (9).
The Arrhenius plot of the first-order rate parameters is shown
in Figure 4. The WLSQR of the rates of disappearance of NH3-
+
d[NH OH]
2k1k
2
OH and the appearance of N2O were both calculated at the
5% confidence interval. The error bars assigned to each rate
3
-
)
[NH OH][NO ]
(9)
3
3
9
dt
k-1 + 2k2
constant on the Arrhenius plot represent the precision of the
individual rate measurements and do not represent the uncer-
tainty in their accuracy. The uncertainty of the rate at each
temperature is better represented by the uncertainties of the slope
and intercept of the WLSQR of all the data shown in the
Arrhenius plot.
The decreasing bulk dielectric constant of H O with increasing
2
2
7
temperature favors the right side of eq 5, i.e., k increases
1
relative to k-1 as the temperature increases. If k-1 , k ,then
2
eq 9 reduces to k [NH OH][NO ]. For the special case of [NH -
1
3
3
2
OH] ) [HNO ], the rate is expected to have apparent first-
3
Reaction Pathway. Extensive work has been devoted to
learning about the mechanism of the reactions of [NH3OH]NO3
order behavior in [NH OH]NO . The same expression applies
3
3
to d[N O]/dt. In Figure 2 the profiles for depletion of NH -
2
3
3,8,10,13-26
+
in aqueous solution.
OH and formation of N O derived from the IR absorbances
2
closely track one another. The incompleteness of the mecha-
nism (5)-(8) at other conditions is evident, however, when
excess NH2OH or HNO3 is added. For NH2OH:HNO3 of 1.0:
4
[NH OH]NO f 3N O + 7H O + 2HNO
(1)
3
3
2
2
3
The limiting stoichiometry of (1)generally describes a 1:1 ratio
of NH2OH and HNO3.
pressure, the Raman scattering intensities of N2O and NO3 in
.87-1.74 m [NH3OH]NO3 are consistent with the slightly
different stoichiometry of (2). It is known, however, that both
the limiting stoichiometry and the reaction details depend on
the conditions of the experiment.14
0
1
.5, the rate indeed decreased by about 0.5, as expected. A
:2 ratio, however, increased the rate by much more than 2 (i.e.,
13-15
At 420-470 K under 27.5 MPa
-
about 10 ( 3). These results are qualitatively consistent with
0
the role of a catalytic species, such as HNO2, whose concentra-
8
19
tion increases as the HNO3 concentration increases.
Arrhenius Kinetics. The activation energy, Ea, for reaction
3
was reported to be 107 kJ/mol from the temperature rise during
1
4
14
reaction in the 289-308 K range. Ea for (4) was estimated
4
[NH OH]NO f 3.2N O + 7.2H O + 1.6HNO +
3 3 2 2 3
to be 65 kJ/mol at 273-298 K from the rate of change of the
-
0
.4O (2)
UV absorbance of NO2 . More recently, the time-to-reaction
2
method based on thermal explosion theory was applied to
The details of the NH2OH-HNO3 reaction have mainly been
aqueous solutions of [NH OH]NO at 420-470 K under 27.5
3
3
8
probed with a low concentration of NH2OH in the presence of
MPa. The results indicated E ) 129 ( 29 kJ/mol for 0.76-
a
a large excess of HNO3.1
4,16-21
This condition enables the
1.52 m solutions and E ) 66 ( 8 kJ/mol for 1.58-1.74 m
a
reaction to be investigated at 273-308 K and minimizes the
amount of heat released. It is generally agreed that both HNO3
and HNO2 competitively oxidize NH2OH by reactions 3 and 4.
solutions. The initial stage of gas evolution from solid [NH -
3
2
8
OH]NO is described by eq 10, where E is in kJ/mol.
3
a
-
1
4.6 -69/RT
k (s ) ) 10 e
(10)
NH OH + 2HNO f 3HNO + H O
(3)
(4)
2
3
2
2
Isothermal measurements of species concentrations are dif-
ficult to conduct with [NH3OH]NO3 because of the rapid
formation of heat during the reaction. To succeed in the current
work, it was necessary to use E0.2 m solutions because the
heat generated was sufficiently small so as to be controlled by
the cell. The concentration-time profiles for which Figure 2
is representative provide rates of conversion in the 463-523 K
range. These data fit the first-order rate expression of eq 9 as
is shown in Figure 3. It was found, however, that a zeroth-
order rate plot gave a statistically equivalent fit. Zeroth-order
behavior could be rationalized by assuming that the decomposi-
tion of [NH3OH]NO3 is mainly wall-catalyzed. However, a
previous decomposition study of [NH3OH]NO3 in the hydro-
NH OH + HNO f N O + 2H O
2
2
2
2
Reaction 3 is most important at small NH2OH:HNO3 ratios,
1
7,19
whereas (4) is increasingly prevalent as this ratio increases.
The decomposition of [NH3OH]NO3 is accelerated by the
formation of HNO2.1
4,17,18,23
Nonlinear behavior in the form
of two stationary reaction states also occurs at small NH2OH:
2
0
HNO3 ratios.
Additional mechanism details have been
provided by 15N and O isotopic labeling studies.
18
23-26
As a result of maintaining the NH2OH:HNO3 ratio at one in
the current work, a mechanism analysis in the sense of most
previous work was not conducted. Also, the behavior of only

8596 J. Phys. Chem. A, Vol. 101, No. 46, 1997
Schoppelrei and Brill
thermal regime revealed only a small difference in the rate in
a 316 SS cell compared to a Ti cell. This result suggests that
dominate, and heat will not be generated at a fast enough rate
to sustain burning. In the absence of H2O, it is possible that
the transport properties and concentrations of the key catalytic
species are too low to accelerate the rate in the stage of the
reaction studied. Only when the [NH3OH]NO3 concentration
is 5 M or higher do all of the necessary properties for fast
reaction become sufficiently optimized to sustain combustion.
As a result, the global Arrhenius parameters yielding the
appropriate rates to model the combustion process of concen-
trated [NH3OH]NO3 solutions are most likely the result of a
different rate-controlling process than exists below 475 K. In
any case the Arrhenius parameters for the combustion regime
are probably a mixture of transport rates and reaction rates in
the heated surface reaction zone.
8
wall catalysis is not a major factor, which makes zeroth-order
kinetics a less appropriate description for the hydrothermal
regime. Alternatively, the complexity of the rate constant in
eq 9 may be responsible for the absence of a simple order
behavior. We have no method to probe the individual steps of
(5)-(8) in the hydrothermal regime.
Table 1 lists the first-order rate constants treated statistically
as discussed above. The precision in the rate constants is
generally satisfactory, but it is apparent that an uncertainty of
a factor of 2 or less exists in the accuracy on the basis of the
differences incurred when the cell was dismantled and reas-
sembled. This degree of uncertainty in the accuracy of rate
constants is in line with our past experience in reaction rate
7,8,29
Acknowledgment. We are grateful for support of this work
by the Army Research Office on the URI Project DAAL03-
measurements at hydrothermal conditions.
A possible
explanation is that flow-rate changes occur from an occasional
phase separation (bubble) somewhere in the flow path. A phase
separation did not take place in the observation region because
no gaseous N2O was detected.
9
2-G-0174. The sample of hydroxylamine in H2O was gener-
ously supplied by Hugo Galletta of Howard Hall Division, R.W.
Greeff & Co.
+
Figure 4 shows the Arrhenius plot for NH3OH destruction
References and Notes
and N2O formation. The rates of these processes were statisti-
cally indistinguishable and are, therefore, included together on
the same plot. The resulting values of Ea ) 103 ( 21 kJ/mol
and ln(A/s) ) 21 ( 5 are obtained. This value of Ea lies close
(
1) Klein, N.; Freedman, E. Proc. JANNAF Propuls. Mtg. 1984, 2,
287.
(2) Klein, N. Prog. Aeronaut. Astronaut. 1988, 109, Chapter 14.
(3) Liquid Propellant XM46 Handbook; Dowler, W. L., Ferraro, N.
W., Eds.; ARDEC: Picatinny Arsenal, NJ, July 1994.
1
4
to that previous estimated for reaction 3 at 292-308 K, but
no preexponential factor was previously available. It is also in
the range of 129 ( 29 kJ/mol given previously based on the
thermal explosion method for <1.5 m solutions of [NH3OH]-
NO3.8 These results suggest that the components of [NH3OH]-
NO3 react by (3), itself a complex process, and that (3) controls
the overall rate. At higher concentrations of [NH3OH]NO3, the
autocatalysis, possibly by the HNO2 thus formed, accelerates
the rate of reaction and release of heat, giving a lower apparent
activation energy, i.e., 66 ( 8 kJ/mol.8
(4) Vosen, S. R. Combust. Sci. Technol. 1989, 68, 85.
(
5) Shaw, B. D.; Williams, F. A. Proceedings of the 24th Symposium
(
International) on Combustion; The Combustion Institute: Pittsburgh, PA,
1992; p 1923.
(6) Klein, N.; Freedman, E. 19th JANNAF Combust. Mtg. 1982, 1.
7) Kieke, M. L.; Schoppelrei, J. W.; Brill, T. B. J. Phys. Chem. 1996,
00, 7455.
(
1
(
8) Schoppelrei, J. W.; Kieke, M. L.; Brill, T. B. J. Phys. Chem. 1996,
1
00, 7463.
(
9) Schoppelrei, J. W.; Kieke, M. L.; Wang, X.; Klein, M. T.; Brill, T.
B. J. Phys. Chem. 1996, 100, 14343.
10) Cronin, J. T.; Brill, T. B. J. Phys. Chem. 1986, 90, 178.
(11) Schoppelrei, J. W.; Brill, T. B. Unpublished results.
12) Cvetanovic, R. J.; Singleton, D. L. Int. J. Chem. Kinet. 1977, 9,
81.
Relation to Combustion. Some of the mechanistic and
(
kinetic details discussed above appear to apply to very concen-
10,22,23,28,30
trated solutions and to solid [NH3OH]NO3.
In
(
particular, accelerating rate observed is consistent with auto-
catalysis. A type of nonlinear behavior is also noted in the
formation rate of the decomposition gases from pyrolysis of
4
(13) McKibben, J. M.; Bercaw, J. E. USAEC, DP-1248, 1971.
(14) Pembridge, J. R.; Stedman, G. J. Chem. Soc., Dalton Trans. 1979,
>
12 M solutions of [NH3OH]NO3 upon flash heating under
1657.
(15) Klein, N. ARBRL-TR-02471, Ballistic Research Laboratory, Ab-
erdeen Proving Ground, MD, Feb 1983.
16) Morgan, T. D. B.; Stedman, G.; Hughes, M. N. J. Chem. Soc. B
968, 344.
.5 MPa of Ar.10 At high concentrations, however, the gaseous
3
10,22,23,30
products frequently include N2, NO, and NO2,
latter two products can form by various secondary reactions in
but the
(
1
31
mixtures containing HNO2 and HNO3. The activation energy
for solid [NH3OH]NO3 (eq 10)28 resembles that for eq 4, which
is consistent with the catalytic role of a species such as HNO2.
The rate constant of eq 10 and the kinetics determined in this
article are the same at 408 K, which is reasonable on the basis
of the roughly similar temperature range in which they were
measured.
(17) Gowland, R. J.; Stedman, G. J. Inorg. Nucl. Chem. 1981, 43, 2859.
14
(18) Bennett, M. B.; Brown, G. B.; Maya, L.; Posey, F. A. Inorg. Chem.
1982, 21, 2461.
19) Gowland, R. J.; Stedman, G. J. Chem. Soc., Chem. Commun. 1983,
038.
(
1
9
1
(
91.
(
61.
20) Horv a´ th, A.; P o´ ta, G.; Stedman, G. Int. J. Chem. Kinet. 1994, 26,
21) Bourke, G. C.; Stedman, G. J. Chem. Soc., Perkin Trans. 1992, 2,
From the practical point of view, however, it is useful to know
whether the kinetics of reaction 2 for a dilute solution or eq 10
for solid [NH3OH]NO3 apply to the combustion of concentrated
(22) Lee, H. S.; Thynell, S. T. 33rd AIAA/ASME/SAE/ASEE Joint
Propulsion Conference, AIAA97-3232, Seattle, WA, July 1997.
(
(
(
23) Oxley, J. C.; Brower, K. R. SPIE 1988, 872, 63.
24) Bothner-By, A. A.; Friedman, L. J. Chem. Phys. 1952, 20, 459.
25) Clusius, K.; Effenberger, E. HelV. Chim. Acta 1955, 38, 1834.
[NH3OH]NO3 solutions. With an estimated surface tempera-
5
ture of 600 K for a burning solution of 9.1 M [NH3OH]NO3,
-
2
-1
the kinetics of neither eq 10 (k ) 4 × 10 s ) nor Figure 4
(26) Hussain, M. A.; Stedman, G.; Hughes, M. N. J. Chem. Soc. B 1968,
597.
(27) Uematsu, M.; Franck, E. U. J. Phys. Chem. Ref. Data 1980, 9,
291.
28) Ratveev, V. A.; Rubszov, U. E. IzV. Akad. Nauk Ser. Khim. 1993,
1, 1897.
-
1
(
k ) 1 s ) are fast enough. The model rate constants required
5
to fit the burning rate of 5-9 M [NH3OH]NO3 solutions are
1
4
5 -1
in the range of k ) 2 × 10 -2.2 × 10 s . Thus, we concluded
that the rate of decomposition of [NH3OH]NO3 is suppressed
either by the large excess of H2O in 0.1-0.2 m [NH3OH]NO3
or by the virtual absence of H2O as in the solid [NH3OH]NO3.
When the [NH3OH]NO3 concentration is low, reaction 4 cannot
(
1
(29) Maiella, P. G.; Brill, T. B. J. Phys. Chem. 1996, 100, 14352.
(30) van Dijk, C. A.; Priest, R. G. Combust. Flame 1984, 57, 15.
(31) Kaiser, E. W.; Wu, C. H. J. Phys. Chem. 1977, 81, 187.