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OCR for page 48
l
6
Heterogeneous Chemical Processes
in Ozone Depletion
MARIO J. MOLINA
Jet Prop vision Laboratory
National Aeronautics and Space Administration
This discussion of heterogeneous chemistry focuses first on the
reaction between hydrogen chloride (MCI) and chlorine nitrate
CONGO. By the term Heterogeneous chemistry" is meant a reac-
tion that does not occur only in the gas phase but instead requires a
condensed phase in order to proceed. As the previous speakers have
indicated, free chlorine atoms are believed to be the principal agent
for destroying ozone. Hydrogen chloride and chlorine nitrate species
act as sinks, or reservoirs, of chlorine in the atmosphere. These
species are not directly reactive with ozone. Hydrogen chloride is a
well-known species with well-known properties. Chlorine nitrate is
more esoteric, requiring a description of some of its properties.
Chlorine nitrate has two peculiar properties in the context of
atmospheric chemistry. One is that in contrast to most other species
containing chlorine and composed of many atoms, it absorbs light
relatively inefficiently; hence it is an efficient sink for storing chIo-
rine. Related species with fewer oxygen atoms photolyze much more
readily and do not serve as efficient reservoirs of chlorine. The other
interesting property of chlorine nitrate is that it is a very difficult
species to synthesize in the laboratory. On the surfaces of reac-
tion vessels, it is chemically unstable and decomposes rapidly. Thus
it came as a surprise to many atmospheric chemists that it was a
comparatively stable chemical in the stratosphere. If it reacts with
48
OCR for page 49
HETEROGENEOUS OH~ICAL PROCESSES
.
49
hydrochloric acid on a suitable surface, then the molecular chlorine
gas that is produced is quickly broken down into chlorine atoms by
the action of absorbed light.
In the context of antarctic chemistry, when this reaction occurs
on ice surfaces, the other product of the reaction, nitric acid, re-
mains bound to the ice. The atomic chlorine, on the other hand,
reacts with ozone to produce chlorine monoxide and molecular oxy-
gen, destroying two ozone molecules in the process. The chlorine
monoxide combines with nitrogen dioxide to reform chlorine nitrate,
thereby partially replenishing the chlorine reservoir and removing
any remaining nitrogen dioxide from the air. The net effect of the
cycle is to release chlorine from the hydrogen chloride (normally very
stable), to scavenge nitrogen from the atmosphere, and to generate
chlorine monoxide while simultaneously converting ozone to molecu-
lar oxygen.
Before the discovery of the antarctic ozone hole, the reaction
between hydrogen chloride and chlorine nitrate was considered to be
improbable for several reasons. The principal surfaces believed to
be effective in the stratosphere in promoting the reaction were those
of sulfuric acid, but they appeared to be only weakly favorable for
making the reaction proceed. In adclition, a simultaneous collmion
of two gaseous molecules on a surface appeared to be necessary in
order for the reaction to occur. Such a simultaneous collmion has a
very low probability of occurring.
In 1986, we began to study reactions that could be taking place in
the polar night stratosphere over Antarctica. We and others surmised
that stratospheric ice clouds relight be important in causing reactions,
since such stratospheric clouds are found almost exclusively in the
Antarctic in winter. When it does get cold enough for ice clouds to
form, the crystals are not just pure ice. The conclusions of other
researchers were that nitric acid would be bound to the solid ice
crystal, even at temperatures a few degrees above freezing, through
the formation of various hydrates of nitric acid, but that hydrochloric
acid would not condense and hence not be bound.
S. Pickering (1893) published a study on the formation of pure
Ice and of hydrogen chloride hydrate crystals at low temperatures
by freezing aqueous solutions. The result of his study was that the
concentration by weight of hydrogen chloride compared to water must
be greater than about 24 percent for any of it to be incorporated into
the frozen crystalline form. This would imply that hydrogen chloride
at usual atmospheric concentrations would have no affinity for ice.
OCR for page 50
l
50
C)
o
-
MARIO J. MOLINA
o
-20
-40
-60
-80
-100
~1 1 1 ~
_ \\ \ LIQUID ~
-\\ \ err
~ \ /HCQ-3H2O
- \\ \' /
SOLID) \ /
_ (ICE) l ~ / SOLID
~\ / (HYDRATES) _
~ V
1
1 1 1 1 1
0 20 40 60
% HC]
FIGURE 6-1 Equilibrium phase diagram for the hydrogen chloride (HC1) water
system as a function of temperature and HC1 concentration.
Other studies conducted at temperatures close to freezing came up
with the same result.
We carried out experiments to check these results. We measured
the hydrogen chloride concentration directly in the solid and in the
liquid phase at temperatures appropriate to achieve equilibrium. For
initially liquid solutions with a concentration of less than 24 percent,
the resultant hydrogen chloride concentration in the solid phase was
between about one-third and one-fourth of the corresponding liquid-
phase value (Figure ~1~. This result, in contradiction to those of the
earlier studies, implies that polar stratospheric ice clouds will absorb
significant amounts of hydrogen chloride vapor.
We concluded that ice is actually very efficient in scavenging
gas-phase hydrogen chloride. We therefore expect that most of the
hydrogen chloride in the antarctic stratosphere is in the condensed
phase and bound to ice crystals when they are present. The nature of
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HETEROGENEOUS CHEMICAL PROCESSES
51
the nitric acid hydrates that are likely to be present in the antarctic
ice crystals needs further study, as the chemistry involved appears to
be more complicated than was originally thought. Nevertheless, the
end result is simply that both hydrogen chloride and nitric acid are
efficiently scavenged by stratospheric ice particles.
We thus realized that if the hydrogen chloride is already in
the ice, a simultaneous collision of two species on a surface is not
required to cause the hydrogen chioride-chIorine nitrate reaction to
occur. Instead, the collision of a chlorine nitrate gas molecule with
an ice particle containing hydrogen chloride, an event with a much
higher probability of occurring, might be sufficient.
We first carried out experiments to measure diffusion rates for
hydrogen chloride in ice at temperatures around 200 K. We used
optical absorption techniques to determine the rate of penetration of
hydrogen chloride into ice. We were surprised to find that hydrogen
chloride moves extremely rapidly within ice. This, we believe, is a re-
sult of the relative ~openness" of the ice crystalline lattice, combined
with the fact that hydrogen chloride molecules are comparatively
small, allowing their rapid movement through the lattice. We con-
cluded that hydrogen chloride diffuses almost as quickly in solid ice
at 200 K as it does in liquid water, in contrast to the conventional
view that diffusion in solids is always much slower. We measured
movements of several millimeters on a time scale of just minutes.
The next experiment we did was to look at the infrared spectra
of ice samples. We compared the spectrum of pure ice with ice
that contained a fraction of a percent of hydrogen chloride: it can
be used as a diagnostic to determine if there is hydrogen chloride,
even in very small amounts, in ice crystals (Figure 6-2~. We did
the same experiment with nitric acid. When it is first deposited on
ice, the spectrum is very similar to that for condensed nitric acid.
After about an hour, the spectrum changes, to show a structure
around wave number 700. This spectrum change indicates that the
nitric acid is being absorbed by the ice. Finally, the spectrum for
ice containing hydrogen chloride and chlorine nitrate is essentially
the same as that for ice containing hydrogen chloride and nitric acid
(Figure ~3~. This implies that when chlorine nitrate is deposited on
an ice crystal that was previously treated with hydrogen chloride,
nitric acid is produced that remains in the solid phase within the ice.
We also carried out experiments with ice-coated tubes in order
to determine the probability of interaction with ice of several gases.
For hydrogen chloride, we determined that at least one of every five
OCR for page 52
l
52
I1J
At:
CY
lo
MARIO J. MOLINA
ICE
ECU+ ICE
I ~
4500 3500 2500
WAVENUMBERS
1500 500
FIGURE 6-2 Infrared spectra of pure ice compared with that for ice containing
approximately 1 percent HCI. (Reprinted, by permission, from Molina et al.,
1988. Copyright @1988 by The American Association for the Advancement of
Science.)
collisions resulted in the hydrogen chloride being scavenged by the
ice. For chlorine nitrate, we also found that the "sticking coefficient"
is very large if the ice on the surface of the tube is doped with
hydrogen chloride. We also measured molecular chlorine as a product
of the collisions. The measured reaction rate is sufficient to explain
the observations in the antarctic stratosphere.
The conclusion of all these experiments is that the chlorine
nitrate-hydrogen chloride reaction is very efficient in the presence
of ice and produces molecular chlorine, thus converting chlorine from
an inactive reservoir form to a form that is readily affected by ultravi-
olet (UV) radiation (Figure 6-4~. There are two other reactions that
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HETEROGENEOUS CHEMICAL PROCESSES
53
HNO3 + HCQ + ICES
Lo
Ad
c
CO
on
`_.
,~_
4:O2 + HCQ + ICE
~4,
~I ~
at,
\
i ~
~1~
/
~ ~ '
1 1 1 1 1
1 1
4500 4000 3500 3000 2500 2000 1500 1000 500
WAVENUMBERS
FIGURE 6-3 Comparison of infrared spectra for ice exposed to small amounts
of the following compounds: HCl (approximately 1 percent), HCl together
with nitric acid (HNO3), and HCl together with chlorine nitrate (ClONO2~.
(Reprinted, by permission, from Molina et al., 1988. Copyright (~1988 by The
American Association for the Advancement of Science.)
can also result in the "activation" of chlorine. One is the reaction of
nitrogen pentoxide (N205) with the hydrogen chloride in the ice to
produce nitric acid and nitry} chloride, which is relatively unstable in
the presence of UV. This reaction has been studied by the SRI group
and also by a group at the Jet Propulsion Laboratory and has been
shown to be an important one for chlorine conversion. The second
reaction is that of hypochIorous acid (HOCI) with hydrogen chloride
to produce molecular chlorine and water. This reaction also proceeds
fairly quickly. Thus, the presence of other species besides chlorine
nitrate can result in liberation of the chlorine from the hydrogen
chloride in the ice.
Clearly, the presence of ice crystals in the atmosphere is very fa-
vorable for the release of chlorine that can destroy ozone. A question
that remains is the relative efficiency of ice that contains nitric acid,
OCR for page 54
54
MARIO J. MOLINA
HC1 + CloNO2-HNO3 (S) + C12
C12 + he 2C1
2tC1+O3-C10+02]
C10 ~ NO2-ClONO2
NET:
HC1 + 2O3 + NO2-HNO3(s) + 2O2 ~ C10
[HC1; > ~ [NO2]
FIGURE 6-4 Series of reactions that result in the destruction of ozone (03)
molecules by chlorine compounds in the presence of polar stratospheric ice
clouds. The net result of these reactions is also shown.
compared to pure ice, in scavenging hydrogen chloride and promot-
ing a reaction with chlorine nitrate. We have done some prelirn~nary
studies to answer this question. When we doped the ice with large
amounts of nitric acid, the reaction no longer occurred. With pro-
gressively smaller amounts of nitric acid, the reaction rate increased.
A very similar situation probably obtains with sulfuric acid. The
critical point is the amount of water available in the condensed phase
for the reaction to occur. Here, we can make a thermodynamic ar-
gument: if the temperature is very close to the freezing point of
water, even if not to the point where pure ice can be crystallized,
the existence of a condensed phase will be sufficient to promote these
reactions. Therefore, it seems that any of these acid solutions that is
sufficiently close to the frost point of water will promote the reactions
in the same way as we have shown.
(In answer to a question on the possible role of large volcanic
eruptions in producing sulfuric acid-water ice): The extent to which
sulfuric acid droplets would be effective in promoting these types
of reactions would depend on the available water that the droplets
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HETEROGENEOUS CHEMICAL PROCESSES
55
contain. Concentrated sulfuric acid (95 percent acid) has the small
amount of water present so tightly bound that the water is not avail-
able for reaction. Fairly dilute solutions may be effective in promot-
ing reactions; laboratory experiments will be needed to determine
this. A plausible argument is that if sulfuric acid droplets exist at
temperatures close to the frost point of water in the surrounding
atmosphere, then they will become very dilute. What happens will
depend critically on the temperature, and experiments should be
conducted with that in mind.
(In answer to a question about the effects of methane increase
and greenhouse gas cooling on stratospheric water-vapor content):
Methane and carbon dioxide are usually considered to be the "good
guys" in terms of countering the loss of ozone due to chlorine. But
the effect of these gases is the opposite for the kind of heterogeneous
chemistry that ~ have discussed. Warming of the surface and lower at-
mosphere implies cooling of the stratosphere. Both the temperature
decrease and the additional water vapor, formed from methane, that
reaches the stratosphere will favor greater formation of stratospheric
ice clouds.
REFERENCES
Molina, M.J., T.-L. Tso, L.T. Molina, and F.C.-Y. Wang. 1988. Antarctic
stratospheric chemistry of chlorine nitrate, hydrogen chloride, and ice:
release of active chlorine. Science 238:1253-1257.
Pickering, S. 1893. Die hydrate der chlorwasserstoffsaure. Ber. Dtsch. Chem.
Ges. 26:277-289.
Representative terms from entire chapter:
chlorine nitrate
