Rate constants for the reactions of XO3-(H2O)n (X = C, HC, and N) and NO3-(HNO3)n with H2SO4: Implications for atmospheric detection of H2</

Article abstract of DOI:10.1021/jp971768h

Relative rate constants for the reactions of CO₃^-(H₂O)_n, HCO₃^-(H₂O)_n, NO₃^-(H₂O)_n, and NO₃^-(HNO_3<

Full text of DOI:10.1021/jp971768h

J. Phys. Chem. A 1997, 101, 8275-8278
8275
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Rate Constants for the Reactions of XO3 (H2O)n (X ) C, HC, and N) and NO3 (HNO3)n
with H2SO4: Implications for Atmospheric Detection of H2SO4
A. A. Viggiano,* John V. Seeley,† Paul L. Mundis,† John S. Williamson,† and Robert A. Morris
Phillips Laboratory, Geophysics Directorate (GPSC), 29 Randolph Road,
Hanscom AFB, Massachusetts 01731-3010
ReceiVed: June 2, 1997; In Final Form: August 13, 1997X
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Relative rate constants for the reactions of CO3 (H2O)n, HCO3 (H2O)n, NO3 (H2O)n, and NO3 (HNO3)n with
H2SO4 have been measured for n ) 0-2, 0-2, 0-3, and 0-3, respectively. All rate constants have been
found to occur in the ratio expected for the collisional values. This strongly indicates that the reactions
proceed at the collision rate, and we assign absolute rate constants on the basis of this assumption. The
present results show that all ions previously used for chemical ionization detection of H2SO4 concentrations
in the atmosphere react rapidly, as has been previously assumed.
Introduction
for n ) 0-2, and
Sulfuric acid is extremely important in the formation of
atmospheric aerosols in several different environments.1 Even
when present in low concentrations, sulfuric acid condensation
processes are often the dominant route to formation of new
aerosols. Examples include the stratospheric aerosol layer1
found at about 20 km and jet engine exhaust.2,3 It is extremely
difficult to detect H2SO4 in the low concentrations responsible
for the aerosol formation. The only presently available tech-
nique to detect the gas-phase H2SO4 in the atmosphere is
chemical ionization mass spectrometry (CIMS).4-11 H2SO4 is
detected by monitoring HSO4--containing product ions from
the reactions of CO3- and NO3- core ions with H2SO4. Usually,
the ions are solvated by HNO3 or hydrated. For instance, in
measuring jet exhaust emissions the following reactions have
been used to derive H2SO4 concentrations,10,12
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NO3 (HNO2) + H2SO4 f HSO4 (HNO3) + HNO2 (5)
These include all the important primary ions used in CIMS
detection of H2SO4.
Experimental Section
The experiments were performed in a variable temperature
selected ion flow tube. The apparatus has been described in
detail elsewhere,15,16 and only details pertinent to the present
experiments will be discussed. The ions were produced in a
supersonic expansion ion source.15 We could not make the
CO3-(H2O)n cluster directly. Instead, O-(H2O)n ions were made
by expanding a mixture of N2O, Ar, and H2O, held at about 4
atm total pressure, through a 25 mm orifice and then ionizing
the gas just downstream of the expansion with an electron
filament (ThO2/Ir) biased at about 50 V with respect to the
expansion orifice. Other configurations were tried but with less
success. Both O-(H2O)n and OH-(H2O)n ions were produced,
at least in part, by reaction of hydrated electrons with N2O.17
The O-(H2O)n and OH-(H2O)n were sampled by a blunt
skimmer and passed into a quadrupole mass filter. The
quadrupole was operated in the “dc off” mode so that more
than one ion species entered the flow tube, which made the
determination of relative rates easier and more accurate. In the
upstream inlet of the flow tube CO2 was added in sufficient
quantity to convert (completely) the O-(H2O)n ions into
CO3-(H2O)n and the OH-(H2O)n into HCO3-(H2O)n upstream
of the reactant inlet. The ions are as written and not CO2
clustered to the original ions. The evidence for this comes from
the fact that water ligands are displaced in this process and
would not be if CO2 simply clustered. The added CO2 did not
disrupt the kinetics measurements since all the ions were
uncoupled from each other due to the lack of H2O in the flow
tube. We attempted to make the CO3-(H2O)n ions in the source
region with very limited success; addition of CO2 and H2O to
the expansion mainly resulted in HCO3-(H2O)n ions. For
CO3-(H2O)2 and HCO3-(H2O)2, C18O2 was used to avoid mass
coincidences with HSO4- and SO4-. The latter ion came from
a small amount of SO3 entering the flow tube from decomposi-
tion of H2SO4. We estimate that the [SO3] was at most 10% of
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CO3 (H2O)n + H2SO4 f
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HSO4 (H2O)m + HCO3 + (n - m)H2O (1)
and
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NO3 (H2O)n + H2SO4 f
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HSO4 (H2O)m + HNO3 + (n - m)H2O (2)
These reactions have been used although the rate constant for
only reaction 2 with n ) 0 had been previously measured,13,14
and no calibration was used. Instead, the assumption that the
reactions occurred at the collision rate independent of solvation
was made. In the present study we have measured the rate
constants for reaction 1, n ) 0-2, and reaction 2, n ) 0-3,
relative to that for reaction 2, n ) 0. In addition, we have
measured
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NO3 (HNO3)n + H2SO4 f HSO4 (HNO3)n + HNO3 (3)
for n ) 1-3,
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HCO3 (H2O)n + H2SO4 f
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HSO4 (H2O)m + H2CO3 + (n - m)H2O (4)
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the [H2SO4] as evidenced by the relative intensities of the SO4
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† Under contract to Wentworth Institute, Boston, MA.
and HSO4 product signals. In most experiments [SO3] was
X Abstract published in AdVance ACS Abstracts, September 15, 1997.
considerably less than 10% of the [H2SO4].
S1089-5639(97)01768-4 CCC: $14.00 © 1997 American Chemical Society

8276 J. Phys. Chem. A, Vol. 101, No. 44, 1997
Viggiano et al.
For NO3-(H2O), two techniques were used. The first was
simply to substitute NO2 in the flow tube for CO2. This allowed
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for rapid switching between CO3 and NO3 core ions.
Alternatively, N2O was left out of the expansion region and
HNO3 was added downstream of the expansion. At low flows
of HNO3 the resulting ions were mainly NO3-(H2O)n. At higher
flows of HNO3, the resulting ions were mainly NO3-(HNO3)n.
A small amount of NO3-(HNO2) was made as well.
The cluster ions were unstable with respect to thermal
decomposition at higher temperatures so that the larger the
cluster, the colder the temperature at which the data were taken.
In this manner small clusters were studied at or near room
temperature, and larger ones could only be studied at temper-
atures of ∼160 K.
The necessity of working at cold temperatures created
problems in introducing H2SO4 into the gas phase. Previous
experiments with H2SO4 had been made above room temper-
ature by passing a dry gas over glass wool coated with
H2SO4.13,14,18 For the present experiments we settled on a
1
Figure 1. Ion signals as a function of He flow through the heated
H2SO4 reservoir. The lines are crude fits to the data and are intended
only as guides in viewing the data.
modified version of a previous H2SO4 inlet design. A /4 in.
glass tube was blown so that a small bulbed section resided
just inside the inlet flange of the flow tube. The bulb region
was stuffed with glass wool. Approximately 1 mL of H2SO4
was added. The glass was wrapped with insulated nichrome
wire in two zones: one around the bulb containing H2SO4 and
one on the straight piece of glass tubing downstream from the
bulb. An RTD monitored the temperature in each of the two
zones. In order not to heat the buffer gas significantly, the inlet
was wrapped with thermal insulation. The insulation was
covered with stainless steel foil so that no electrically insulating
surface was exposed to the ions. Insulating surfaces charge in
the unipolar environment of a selected ion flow tube and cause
disruption of the ion signal. Both zones of the inlet were
typically run at ∼373 K.
are clearly secondary and tertiary products, respectively. No
evidence was found for direct formation of any of these clusters.
The possible conclusions are that (1) no neutral H2SO4 clusters
are formed, (2) any H2SO4 clusters react just like unclustered
H2SO4, or (3) any H2SO4 clusters are completely unreactive.
We interpret our data to be due to either case 1 or case 3, since
it seems likely that (H2SO4)n clusters would directly form some
HSO4-(H2SO4)n ions. The total density of H2SO4 in the flow
tube is between 1011 and 1012 cm-3
.
Table 1 lists the relative rate constants for the reactions
studied here at five different temperatures. At 300 K, the only
cluster ion that could be studied was HCO3-(H2O), and the rate
constants are measured relative to that for bare CO3-. At 273
K, the rate constants are relative to that for NO3-, which in
turn is assumed to be the same relative rate constant as found
at 300 K. At 233 K the rate constants are relative to that of
bare CO3-. The 200 K rate constants are relative to that of
NO3-(H2O), which was assumed to be the same as found at
higher temperatures. Finally, at 158 K the rate constants are
relative to that of NO3-(HNO3), which is assumed to be the
average of the relative rate constants found at higher temper-
Decay curves were taken by varying the flow of helium over
the hot H2SO4. The decays of two or more ions in the flow
tube were monitored so that relative rates could be measured
simultaneously. The relative rate constants were determined
simply from the relative slopes in this case. When changing
core ions, the NO2 and CO2 in the flow tube were switched
rapidly, and decays were measured minutes apart. The gases
were then switched back to check for any drift in the H2SO4
source.
No attempt was made to mass select the ions upstream since
we were only attempting to measure relative rates. All product
ions had HSO4- cores. No attempt was made to determine the
distribution of ligands on the product ions. All data are the
average of several runs, and the variation in the relative rates
from run to run was within 10% and often only a few percent.
atures. In all cases the relative rate constants are close to unity,
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ranging from 0.69 to 1.1 of that for CO3
.
For the most part the relative rate constants were found to
be independent of temperature. However, for NO3-(HNO3)2
and NO3-(HNO3)3 the rate constants are larger at warmer
temperatures. We believe this is due to the hot helium/H2SO4
mixture issuing from the H2SO4 bulb thermally decomposing
some of these ions at the highest temperatures at which they
were studied. The anomalous rate constants were measured at
the highest temperature where the cluster of interest was
thermally stable with respect to surviving the length of the flow
tube. Thermal decomposition of a portion of the reactant ions
down the length of the flow tube was certainly possible and
indeed probable for the highest flow tube temperatures used.
Therefore, a slight increase in buffer gas temperature (even
locally) with increasing addition of the hot helium/H2SO4
mixture from the H2SO4 bulb could have had an appreciable
effect on the ion signals and therefore the rate constants.
Results and Discussion
Figure 1 shows a typical decay plot for the reaction of
CO18O2-(H2O)2 with H2SO4 at 208 K. It shows the primary
decay and four HSO4--containing products. Other primaries,
namely CO18O2-(H2O) and HCO18O2-(H2O)2, were present in
small quantities and are not included in the figure since our
computer program can only monitor five masses at one time,
and the decays were not recorded for this run. Other runs
showed that all the primaries decayed at essentially the same
rate. The decay of the primary is reasonably linear on this
semilogarithmic plot. One concern in taking these data was
that a considerable amount of the reaction was due to (H2SO4)n
clusters. The data in Figure 1 address this issue. The data
Table 1 contains a column where the relative rate constants
for each temperature are combined to yield relative rate constants
for each ion with respect to CO3-. For all reactions that did
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clearly show that the primary product ions are about 70% HSO4
and 30% HSO4-(H2O). The HSO4-(H2SO4) and HSO4-(H2SO4)2

Reactions of XO3-(H2O)n and NO3-(HNO3)n with H2SO4
J. Phys. Chem. A, Vol. 101, No. 44, 1997 8277
TABLE 1: Relative Rate Constants of Chemical Ionization Primary Ions with H2SO4
temperature (K)
relative
a
d
ion
300
(1)e
273
233
200
158
k
kcb
(1)
k/kcc
k300
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CO3
(1)
0.89
(1)
(1)
(2.39)
2.15
2.39
2.63
1.98
2.37
2.32
2.06
1.89
1.89
1.86
1.72
1.79
1.86
CO3-(H2O)
CO3-(H2O)2
0.90
1.0
0.90
1.0
1.1
0.92
0.87
0.99
0.92
0.86
0.99
0.92
0.87
0.82
0.76
0.72
0.85
0.84
0.98
1.15
1.11
0.9
1.15
0.98
0.94
0.91
0.96
1.03
1.00
0.87
0.93
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HCO3
1.1
0.83
HCO3-(H2O)
0.82
0.83
0.99
0.97
0.86
0.79
0.79
0.78
0.72
0.75
0.78
HCO3-(H2O)2
0.99
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NO3
0.99
(0.99)
0.86
0.95
0.85
NO3-(H2O)
(0.86)
0.79
NO3-(H2O)2
NO3-(H2O)3
NO3-(HNO3)
NO3-(HNO3)2
NO3-(HNO3)3
NO3-(HNO2)
0.79
(0.78)
0.69
0.75
0.80
0.74
{1.1}
0.78
{0.87}f
0.75
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a Average relative rate constants. b Calculated collision rates ratioed to the calculated collision rate for CO3
.
c Ratio of previous two columns.
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d Calculated rate constants at 300 K based on the relative rates and assumption that the CO3 rate is collisional. Units are cm3 s-1 × 10-9
.
e ( )
indicates the ion whose rate constant the other rate constants are relative to. f { } indicates the rate constant is not used in the average; see text for
explanation.
not suffer from thermal decomposition, no temperature depen-
dence was found, justifying taking relative rate constants at
different temperatures and combining them. Averages were
taken when an ion was studied at more than one temperature,
with the anomalously high rate constants for NO3-(HNO3)2 and
NO3-(HNO3)3 omitted from the averages for these ions. Also
included is a column labeled kc, for collision rate constants
normalized to that for bare CO3-. Another column shows the
ratio of the relative rate constants to the relative collisional rate
constants. These numbers vary from 0.87 to 1.15. The two
highest numbers are for the reactions of CO3-(H2O)2 and
HCO3-(H2O)2 at the highest temperatures at which they were
studied. This probably indicates that they also suffer from the
thermal decomposition problem discussed above. This appears
to affect the measured rate constant by 15%.
The present findings of all negative ions reacting rapidly with
H2SO4 are consistent with the previous findings of Viggiano et
al.13,14 that showed similar behavior for the reactions of H2SO4
with a number of negative ions. In that study it was found that
the rate constant for NO3-(HNO3)2 was only 0.6 the collision
rate. The present results show no such trend and indicate that
n ) 3 is also fast. The previous measurements were taken at
343 K, and the warmest temperature at which NO3-(HNO3)2
was reliably studied in the present set of experiments was 200
K. Normally, ion-molecule reactions become faster at lower
temperatures so this may explain the discrepancy.23 Alterna-
tively, the previous data were taken in a flowing afterglow in
which large amounts of HNO3 were added. In similar com-
parisons of selected ion flow tube data to flowing afterglow
data involving clusters, the flowing afterglow data were
shown to be in error due to equilibrium processes occurring.24,25
Tanner and Eisele22 have concluded that the reactions of
NO3-(HNO3)(H2O)n ions also did not depend on the level of
hydration, consistent with the present results.
The fact that the relative rate constants are in the same ratio
as that expected for the collision rates indicates that all the
reactions occur at the collision rate within about 10%. With
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this in mind we set the rate constant for CO3 equal to its
collisional value and scaled all values to this. Collision rate
constants are calculated by the method of Su and Chesna-
vich,19,20 and the molecular parameters are the same as those
used by Viggano et al.14 This is shown for 300 K in the last
column of Table 1.
The present results show that all ions previously used for
CIMS detection of H2SO4 concentrations in the atmosphere react
rapidly with H2SO4, producing HSO4 core ions. Therefore,
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the assumption of calculating the collision rate constant for the
ion of interest with H2SO4 has led to little error in previous
measurements of H2SO4 concentrations. This is important since
we have recently shown that for CO3-(H2O)n reactions with
SO2 the rate constants do depend on the hydration number and
that either in-flight calibration or a knowledge of the hydrate
distribution is needed for the CIMS measurements of SO2 with
this series of core ions.26
One should note, however, that the collisional value depends
relatively strongly on temperature so that the appropriate rate
constant for a stratospheric balloon mission is about 15% higher
than the room-temperature value. The assumption that the
reactions occur at the collisional value is supported by the work
of Eisele and co-workers.21,22 They titrate OH into H2SO4 and
then use the reaction of NO3-(HNO3) with H2SO4 to detect the
OH. Use of the collisional value for the H2SO4 reaction gives
the same answer as a calibrated OH source within 10%.
It is difficult to put error limits on the data. Although the
assumptions used to derive the rate constants warrant relatively
large uncertainty, a number of factors indicate that the measure-
ments are better than one might expect. Two indications that
the relative rate constant measurements are probably precise to
within 10-15% are the fact that there is such good reproduc-
ibility and the tight scatter about the ratio of rates to expected
collisional values. The reactions where thermal dissociation
may be occurring are obvious exceptions. The assumption that
the rate constants are collisional is corroborated by the good
agreement in the OH calibration mentioned in the previous
paragraph.
Acknowledgment. The research was supported by the Air
Force Office of Scientific Research under task 2303EP4, the
NASA Atmospheric Effects of Aviation Program, and the
Strategic Environmental Research and Development Program.
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