The mechanisms of molecular halogen production from frozen saline surfaces
remain incompletely understood, limiting our ability to predict atmospheric
oxidation and composition in polar regions. In this laboratory study,
condensed-phase hydroxyl radicals (OH) were photochemically generated in
frozen saltwater solutions that mimicked the ionic composition of ocean
water. These hydroxyl radicals were found to oxidize Cl-, Br-, and
I-, leading to the release of Cl2, Br2, I2, and IBr. At
moderately acidic pH (buffered between 4.5 and 4.8), irradiation of ice
containing OH precursors (either of hydrogen peroxide or nitrite ion)
produced elevated amounts of I2. Subsequent addition of O3
produced additional I2, as well as small amounts of Br2. At lower
pH (1.7–2.2) and in the presence of an OH precursor, rapid dark conversion
of I- to I2 occurred from reactions with hydrogen peroxide or
nitrite, followed by substantial photochemical production of Br2 upon
irradiation. Exposure to O3 under these low pH conditions also
increased production of Br2 and I2; this likely results from
direct O3 reactions with halides, as well as the production of
gas-phase HOBr and HOI that subsequently diffuse to frozen solution to react
with Br- and I-. Photochemical production of Cl2 was only
observed when the irradiated sample was composed of high-purity NaCl and
hydrogen peroxide (acting as the OH precursor) at pH = 1.8. Though
condensed-phase OH was shown to produce Cl2 in this study, kinetics
calculations suggest that heterogeneous recycling chemistry may be equally
or more important for Cl2 production in the Arctic atmosphere. The
condensed-phase OH-mediated halogen production mechanisms demonstrated here
are consistent with those proposed from recent Arctic field observations of
molecular halogen production from snowpacks. These reactions, even if slow,
may be important for providing seed halogens to the Arctic atmosphere. Our
results suggest the observed molecular halogen products are dependent on the
relative concentrations of halides at the ice surface, as we only observe
what diffuses to the air–surface interface.
Introduction
It is now well established that gas-phase halogen species influence
atmospheric composition through reactions with ozone (O3), volatile
organic compounds (VOCs), and gaseous elemental mercury (Hg0)
(Barrie
and Platt, 1997; Carpenter et al., 2013; Platt and Hönninger, 2003;
Saiz-Lopez and von Glasow, 2012; Simpson et al., 2007, 2015; Steffen et al.,
2008, 2014, and references therein). In polar regions, it is believed that
halogens build up to effective concentrations through a heterogeneous
reaction sequence known as the “halogen explosion” (Reactions R1–R4, where
X represents Cl, Br, or I)
(Garland and Curtis, 1981;
Tang and McConnell, 1996; Vogt et al., 1996; Wennberg, 1999).
R1X2+hν→2XR2X+O3→XO+O2R3XO+HO2→HOX+O2R4HOX+X-+H+→X2+H2O
In this sequence, a molecular halogen (X2) is photolyzed to produce two
reactive halogen radicals. These radicals can react with O3 to produce
halogen oxides (XO). The XO produced in Reaction (R2) rapidly photolyzes (or
reacts with NO) to regenerate O3 and X2 in a null cycle. To
irreversibly remove ambient O3, XO must react with another halogen
oxide or Hg0. Alternatively, XO can react with HO2 to form HOX
(Reaction R3) or NO2 to form XONO2. Gas-phase HOX can
heterogeneously react with salt-laden surfaces, including sea-salt aerosol
particles (McConnell et al., 1992) and the
“disordered interface” (often referred to as a quasi-liquid or quasi-brine
layer) that exists on frozen saline surfaces (Bartels-Rausch
et al., 2014; Cho et al., 2002) to produce X2, effectively returning
two halogen radicals to the gas phase. Additionally, this mechanism is
enhanced under acidic conditions, confirmed by laboratory studies of aqueous
(Fickert et al., 1999) and frozen solutions (e.g.,
Abbatt et al., 2010; Sjostedt and Abbatt, 2008; Wren et al., 2013) and by field observations (Pratt et al., 2013).
While much has been learned about the atmospheric chemistry of reactive
halogen species in the Arctic, knowledge gaps remain in the chemical
mechanisms by which molecular halogens are produced from frozen surfaces
(Liao et
al., 2014; Pratt et al., 2013). Recently, in situ, light-induced production
of Cl2 (Custard et al., 2016), Br2
(Pratt et al., 2013; Raso
et al., 2017), and I2 (Raso et al., 2017) within
snowpack interstitial air has been reported and was further demonstrated to
be enhanced following the addition of O3. The Br2-producing
snowpacks studied by Pratt et al. (2013) were
characterized as having a larger surface area, lower pH (≤6.3), greater
[Br-]/[Cl-] molar ratios (≥1/148), and lower salinity
relative to other frozen samples collected near Utqiaġvik, Alaska. The
proposed mechanism for this chemistry is based on laboratory studies of
condensed-phase, hydroxyl radical (OH)-mediated halogen oxidation (Reactions R5–R12), which is followed by partitioning of the molecular halogen to the
gas phase
(Abbatt
et al., 2010; Knipping et al., 2000; Oum et al., 1998b).
R5H2O2+hν→2OHR6NO2-+hν→NO+O-R7O-+H+→OHR8OH+X-↔HOX-R9HOX-+H+→X+H2O
R10X+X-↔X2-R112X2-→X3-+X-R12X3-↔X-+X2
Direct, light-induced halogen production from frozen surfaces in the
presence of OH has been previously demonstrated in the laboratory for
Br2 and possibly for I2 (Abbatt et al., 2010), but analogous
chemistry for Cl2 has yet to be observed. Additionally, photochemical
production of I2 has been directly observed in the absence of OH
(Kim et al., 2016). Employing cavity
ring-down spectroscopy, Kim et al. (2016) reported photochemical production
of I2 from a frozen solution by known aqueous-phase chemistry (Reactions R13–R17).
This proposed photochemical mechanism involves an (I-–O2)
charge-transfer complex (Levanon and Navon, 1969).
R13O2(aq)+4H++6I-→2I3-+2H2OR14I-+O2→I-O2→hνI+O2-R15I+I-↔I2-R162I2-→I3-+I-R17I3-↔I-+I2
Kim et al. (2016) also report enhanced photochemical I3-
production (determined spectrophotometrically) from sunlit frozen iodide
solutions placed on Antarctic snowpack, as well as from refrozen field snow
and glacier samples doped with iodide. A question is thus raised regarding
the necessity of OH for I2 production under environmentally relevant
conditions.
The role of O3 in halogen production on frozen surfaces is also
unclear. Previous laboratory studies have demonstrated that halide-doped
frozen surfaces exposed to O3 can lead to Br2 production
(independent of radiation; Reactions R18–R19 and R4)
(Oldridge and Abbatt, 2011; Oum et al., 1998a; Wren et al., 2013).
R18O3+Br-↔BrO-+O2R19BrO-+H+↔HOBr
It has recently been shown that this process proceeds at the surface,
through a water-stabilized ozonide, Br⚫OOO-, as shown in
Reactions (R20)–(R22). Artiglia et al. (2017) observed this Br⚫OOO- intermediate via liquid-injection X-ray photoelectron
spectroscopy.
R20Br-+O3→Br⚫OOO-R21Br⚫OOO-+H+→HOBr+O2R22Br⚫OOO-+H2O→HOBr+O2+OH-
Wren et al. (2013) found that Cl2
was produced primarily via heterogeneous recycling of HOCl, resulting from
BrCl photolysis, on halide-rich artificial snow. However, the observation
that O3 induces halogen production from natural frozen surfaces has yet
to be confirmed by field observations of snowpack chemistry, in which
exposure to only O3 in the absence of light has not been shown to
produce molecular halogens (Custard
et al., 2017; Pratt et al., 2013; Raso et al., 2017). This raises a question
of whether O3 is more important for initial halogen release or in a
gas-phase propagation/recycling capacity (i.e., per the halogen explosion).
In this study, we utilized a custom ice-coated-wall flow reactor in tandem
with chemical ionization mass spectrometry to study Br2, Cl2,
and I2 production from frozen surfaces with compositions mimicking
sea ice. The effects of photochemically generated OH radicals, O3
addition, and pH are tested as they relate to the production of these
halogens. Surface pH was controlled through the use of buffers.
MethodsMaterials
Sample solutions were composed to mimic the halide composition of seawater.
This was achieved using either dissolved Instant Ocean (Spectrum Brands) or
commercially available halide salts at a composition that mimics Instant
Ocean (for consistency) in solutions referred to hereafter as
“saltwater”. The halide
concentrations in these solutions were made to a final concentration of
0.56 M Cl-, 7.2 × 10-4 M Br-, and
1.9 × 10-6 M I-. Except for Instant Ocean, all chemicals were purchased
from Sigma Aldrich. Halide salts include solid NaCl (puriss. p.a. grade, ≥99.5 % purity), NaBr (puriss. grade, > 99 % purity), and
KI (puriss. p.a. grade, ≥ 99.5 % purity). We note that these halide
concentrations are comparable to those in actual seawater (Herring and Liss,
1974; Luther et al., 1988; Tsunogai and Sase, 1969), which typically contains
Cl-, Br-, and I- at ratios of
1:1/660:1/200000. Solutes were dissolved in ultrapure water
(Birck Nanotechnology Center). Dissolved organic carbon for Instant Ocean and
halide salt solutions was analyzed using a Shimadzu TOC-VCSH
Total Organic Carbon Analyzer and determined at approximately
70 mg L-1 for Instant Ocean solutions and less than 5 mg L-1
for saltwater solutions. No further characterization of carbon-containing
compounds was performed.
While previous investigators have adjusted the pre-freezing pH of their
samples, it is very difficult to know the pH in the surface brine (or
disordered interface) of frozen samples (Bartels-Rausch
et al., 2014), though there is evidence from laboratory studies suggesting
that the pH of salt solutions remains largely unchanged after freezing
(Wren and Donaldson, 2012b). To obviate this problem,
the aqueous solutions used in this study were buffered so that the same pH
should exist in the surface brine layer. All solutions were buffered by
either a 20 mM acetic acid (ACS reagent grade, ≥99.7 %
purity)/acetate (puriss. p.a. grade) buffer (pH ≈ 4.5–4.7) or a 20 mM bisulfate (ReagentPlus grade, 99 % purity)/sulfate (ReagentPlus grade,
≥99.0 % purity) buffer (pH ≈ 1.7–2.2). These buffer
concentrations were chosen as a compromise between using as little buffer as
possible, yet enough buffer to ensure adequate buffering ability, as buffer
capacity rapidly decreases as constituent species concentrations approach
the acid Ka value. pH values of sample solutions were determined
before and after experiments with no significant changes observed,
suggesting the buffer composition/buffering capacity does not appreciably
change over the course of an experiment (discussed further in the
Supplement). 100 µM of either hydrogen peroxide (trace
analysis grade, ≥30 % purity) or sodium nitrite (ReagentPlus grade,
≥99.0 % purity) was included as photochemical hydroxyl radical
precursors, via Reactions (R5)–(R7).
Flow tube
Experiments were performed in a custom-built 150 cm long, 2.5 cm ID
frozen-walled Pyrex flow tube contained within a temperature-controlled
cooling jacket. In each experiment, 80.0 mL of sample solution was poured
into the tube in the presence of room air, which was subsequently sealed
with vinyl caps (McMaster-Carr). The flow tube was then rotated on motorized
rollers within a 170 cm × 50 cm × 50 cm insulated wooden cooling chamber.
Crushed dry ice was placed along the bottom of the chamber, and fans were
used to circulate the air throughout the chamber such that the flow tube was
evenly cooled. After ∼30 min, the sample was evenly
frozen (ice thickness of 0.9 mm). The flow tube was subsequently transferred
to an enclosed 156 cm × 50 cm × 50 cm wooden Mylar-lined experiment
chamber and connected to a recycling chiller set to 258 K (i.e., above the
NaCl⚫2H2O eutectic point. At this temperature, the relevant chemical
reactions are expected to occur with/in a brine on the ice surface; Cho et al., 2002;
Oldridge and Abbatt, 2011). This conjecture is based on the work of
Oldridge and Abbatt (2011), who reported from a series of similar
experiments that when O3 is flowed over frozen NaCl/NaBr solutions
above the NaCl eutectic temperature, reaction kinetics were strongly
consistent with chemistry occurring in a liquid brine. The cooling liquid
used for the chiller was a mixture of 60 % ethylene glycol and 40 %
distilled water. Six UVA-340 solar simulator lamps (Q-Lab, 295–400 nm
with maximum wattage at 340 nm, irradiance spectrum in Fig. S1 in the Supplement) were
installed in the experiment box (two on each side except bottom). Each side
was lined with reflective Mylar sheets to evenly irradiate the flow tube
when the lamps were powered.
Experimental schematic. Purple bars represent powered solar
simulator bulbs. The green shading around the flow tube (flow reactor)
represents cooling liquid (60 % ethylene glycol, 40 % water) circulated
through the chiller. The flow reactor region itself has an inner diameter of
2.5 cm.
A flow schematic representing typical experiments is shown in Fig. 1. The
carrier gas (Air, Ultra Zero grade, Praxair) was scrubbed of volatile
organic compounds using activated charcoal and water by traveling through
coiled stainless-steel tubing surrounded by crushed dry ice (replaced
throughout the course of an experiment). This gas was measured to contain
≤300–400 pmol mol-1 NO (experimentally determined limits of
detection) using the Total REactive Nitrogen Instrument (TRENI) (Lockwood
et al., 2010; Xiong et al., 2015). Though NO2 was not measured, it
should have been removed by the charcoal trap. Before entering the
coated-wall flow tube, the carrier gas flowed through a commercial O3
generator (2B Technologies model 306). Carrier gas air entered the tube near
room temperature (20 ∘C). At the start of experiments, the O3
generator was set to 0 nmol mol-1. Carrier gas then entered the flow
tube in the dark experiment chamber. In most experiments, the carrier gas
was regulated to a volumetric flow rate of 4.0 L min-1, which yields a
residence time in the flow tube of ∼12 s. On exiting the
flow tube, sample air was characterized using a Thermo Environmental 49i
O3 monitor (flow rate of ∼1.5 L min-1) and a
chemical ionization mass spectrometer (CIMS; sampling flow rate of
∼1.7 L min-1; described below in Sect. 2.3). Excess flow
air was vented away. At set times in an experiment, the solar simulator
bulbs were activated, and O3 was added to the system by powering the
O3 generator. At the end of each experiment, the ice was melted and the
water collected for pH measurements. To clean the flow tube, its interior
was washed three times with ultrapure water before a final rinse with wash
acetone. The flow tube was then connected to a compressed nitrogen gas
cylinder (Praxiar, > 99.99 % purity) to dry for at least 2 h. Once dry, the flow tube was disconnected and capped until the next
experiment.
CIMS
Halogen species were detected using a chemical ionization mass spectrometer
(CIMS), described previously by Liao et al. (2011) and Pratt et al. (2013).
Chemical ionization is achieved by ion-molecule reactions that occur between
iodide–water reagent clusters, I(H2O)n- in N2, and the
gas-phase analytes in zero air. The iodide–water clusters are formed when
gas-phase iodide ions, generated by flowing 5 ppm methyl iodide through a
210Po ionizer (NRD), combine with water in the humidified ion-molecule
region of the CIMS. Ion were filtered using a quadrupole mass filter. The
ice-coated flow tube was connected to the CIMS via approximately 50 cm of
i.d. 1/2 in. PFA Teflon tubing.
List of relevant species monitored by chemical ionization mass
spectrometry (I(H2O)n- as reagent ion) with corresponding
m/z values.
A typical CIMS sampling cycle consisted of an 8.35 s duty cycle. Dwell times
for all monitored species were 250 ms except for the reagent ion (detected
as m/z 147, I(H218O)-), which was set to a dwell time of 100 ms.
The 18 ions analyzed in this study are listed in Table 1, but we focus
herein on results concerning masses related to Br2 (m/z 285 and 287:
I79Br79Br- and I81Br79Br-, respectively),
Cl2 (m/z 197, 199, and 201: I35Cl35Cl-,
I37Cl35Cl-, and I37Cl37Cl-), and I2
(m/z 381: I3-). In addition, IBr (m/z 333 and 335: I79IBr-,
I81IBr-) was unambiguously detected in some experiments. The
presence of Br2, Cl2, and IBr was confirmed by measuring the
ratios between the two isotope signals for each mass, compared to the
natural abundances (i.e., 1.95 for m/z287:285; 1.54 for m/z197:199; and 1.03 for
m/z333:335, respectively). Data outside ±25 % the expected isotope
ratio were excluded from analysis. The signals for BrCl (m/z 241 and 243:
I79Br35Cl-, I81Br35Cl-,
I79Br37Cl-) masses were never observed at the correct ratios
(1.3 for m/z243:241), and so those data were not reported here. As the
introduction of ∼60 nmol mol-1O3 to the
experimental system significantly increased the baseline signal of m/z 197, but
not m/z 199 or 201, the presence of Cl2 could not be confirmed under
elevated O3 conditions. In addition, background-subtracted, relative
signals for m/z 271 (IHOI-) and m/z 225 (IHO81Br-) are discussed
(signals are relative to that of the ionization gas (m/z 147,
I(H218O)-). According to isotope ratios, IHOBr- was not
unambiguously observed, however, due to an interference at m/z 223
(IHO79Br-), and our results here should be considered for only
qualitative purposes as we only discuss relative changes in the signal.
CIMS calibrations were performed using I2, Br2, and Cl2
permeation devices (VICI) at the start and conclusion of each experiment.
Br2 and Cl2 permeation outputs were quantified using the
spectrophotometric method described by Liao et al. (2012).
The I2 permeation output was quantified by flowing the I2 through
an impinger containing a NaHCO3 (30 mM)/NaHSO3 (5 mM) reducing
solution. This solution quantitatively reduces I2 to I-, which was
then determined using a Dionex DX500 ion chromatography system. Permeation
rates were calculated for each experiment and found to average (1.9±0.1) × 10-11, (5.5±0.1) × 10-10, and (8.6±0.1) × 10-10 mol min-1
of I2, Br2, and Cl2, respectively
(uncertainties representing standard error of the mean). CIMS calibration
factors were calculated for individual experiments. These factors are based
on the average of the signal sensitivities, determined from the permeation
sources, calculated at the start and completion of each experiment.
Corresponding uncertainties for these calibration factors thus represent the
1σ standard deviation of the mean sensitivity. An approximate
I79IBr- calibration factor was assumed to be the average of the
sensitivities for m/z 287 (IBr2-) and 381 (I3-).
Background measurements were performed before and after the experiment
(minimum of 5 min) by passing the carrier gas through the experimental flow
tube (without O3, in the dark) and subsequently through a glass wool
scrubber, previously shown to remove molecular halogens with greater than
95 % efficiency (Liao
et al., 2012; Neuman et al., 2010). Temporal variations in bromine-species
signals while using the low pH sulfate/bisulfate buffer were observed in
some experiments (Fig. S2) and are discussed in the Supplement.
Analysis of experimental data was based on 1 min averages, with
uncertainties representing the standard deviation of these averages.
Subsequently, signals were converted to concentrations using the
sensitivities calculated above, propagating the sensitivity uncertainty into
the measurement uncertainty. Average limits of detection (3σ) across
all experiments for the molecular halogens during background periods were
1.8±0.4, 1.2±0.3, and 9±2 pmol mol-1 for
Br2, Cl2, and I2 respectively (uncertainties representing
standard error of the mean). Additionally, reported uncertainties for
integrated amounts of formed halogens are calculated as integrated halogen
concentrations multiplied by the relative uncertainty in the CIMS signal
sensitivity.
Results and discussion
The experiments described here address the extent to which condensed-phase
OH radicals in an ice surface brine (Cho et al., 2002) can produce I2,
Br2, and Cl2 through condensed-phase reactions within frozen
saline surfaces, as hypothesized by recent field (Custard
et al., 2017; Pratt et al., 2013; Raso et al., 2017) and laboratory
experiments (Abbatt et al.,
2010). In addition, we test the pH dependence of this chemistry and whether
gas-phase O3 enhances this production. We find the relative and
absolute amounts of halogens produced from ice are a complex function
of the relative concentrations of the precursor halide ions, pH, presence of
oxidants, radiation, and O3.
Results for all experiments performed. The first line in an
experiment represents the integrated totals of molecular halogen production
after 1 h of irradiation (t=0 through t=1 h). The results on
italicized lines are 1 h integrated production amounts beginning once
additional ozone was introduced to the flow tube. Average LODs across
experiments were 1.8±0.4, 1.2±0.3, and 9±2 pmol mol-1 for Br2, Cl2, and I2
respectively. “IO#” represents samples composed of Instant Ocean, and
“SW#” represents saltwater samples, composed of reagent salts. “CL1”
here represents the experiment performed using 0.56 M high-purity NaCl.
* The Br2 and
I2 values presented for IO2 are discussed further in the Supplement.
The ice-coated flow tube experiments started under dark conditions and
without addition of O3 (Sect. 3.1). Once signals stabilized, lights
were activated (Sect. 3.2). After 1–2 h, ∼60 nmol mol-1 of O3 was introduced into the carrier gas (Sect. 3.3).
Integrated amounts of produced molecular halogens are presented in Table 2
for all experiments. Unless otherwise specified, integrated amounts of
produced halogens represent amounts produced over the course of 1 h of
exposure to light (Sect. 3.2) and/or ozone (Sect. 3.3). Saline ices tested
include frozen Instant Ocean (IO) solutions, saltwater (SW) solutions
composed of dissolved reagent grade salts mimicking seawater composition,
and 0.56 M high-purity NaCl (CL1). OH-radical precursors used include
hydrogen peroxide (H2O2) or nitrite (NO2-), which have
been estimated to account for 96 % of snowpack photochemical OH formation
at Utqiaġvik, Alaska (France et al., 2012). Many of
the salient features of our results are demonstrated by example experiments
shown in Fig. 2, including the impact of irradiation in the presence of ice-phase OH-radical precursors, varied pH, and the presence of O3. Below
we discuss the results and interpretations of our experiments, organized by
the mechanism of halogen production and halogen products themselves.
Representative experiments of OH-mediated production of X2
and subsequent production of X2 from O3 addition. (a) Saltwater
experiment (SW2) at pH = 4.5. (b) Instant Ocean experiment (IO4) at pH = 1.8.
Time-varying Br2 and IBr signals before t=0 are shown in Fig. S2 in the Supplement. Inset more clearly shows the increase of I2 signal after
irradiation. (c) NaCl experiment (CL1) at pH = 1.8. Timescale represents
hours from the activation of the lights, and the yellow shading represents
presence of radiation from solar simulator bulbs. Gaps in data represent
periods when the isotopic ratios showed an interference.
Dark reaction production of I2
After the initial connection of the flow tube to the CIMS (i.e., before
irradiation or addition of O3), large I2 signals (measured as
I3-, m/z 381) were observed in several cases in which OH-radical
precursors were utilized, especially when pH ≤ 2 (e.g., Figs. 2b and
S2). Integrated calibrated sums of this dark I2 production are
estimated in Table S1 in the Supplement and span from the time when the flow tube was
connected to the CIMS until lights were activated. When pH ≤ 2, dark
production of I2 sometimes caused significant depletion of reservoir
I-. Experiment IO4 and SW5 (both using H2O2 as an OH precursor) only had, at most, ∼36 % of the
initial 152 nmol of I- by the time lights were turned on (remaining I- was
estimated by subtracting twice the observed I2 (i.e., two I- for
every I2) from the initial 152 nmol of I- in the IO or SW
solutions). Considerably less dark I2 production occurred using
NO2- as an OH precursor (depleting I- by an average of
4.5 %; Table S1). However, the amounts in Table S1 represent lower limits
of the dark-produced I2; it is impossible to accurately determine the
extent of dark I2 production since some was lost from the flow tube
during its connection to the CIMS after freezing (Fig. 1). At pH ≈ 4.7, this production was relatively modest. Only Experiment IO2 was
noticeably affected, in which only ∼0.5 % of initial
I- was removed by dark mechanisms (Table S1). Under both pH regimes
(i.e., ∼4.7 and < 2), this signal subsequently decayed
as I2 flushed out of the system until a low steady-state
concentration was reached. No corresponding dark production of Br2 or Cl2 was
observed for any experiments at any pH.
As previously reported, both H2O2 and NO2- can directly
convert I- to I2 under dark acidic conditions. The oxidation of
I- by H2O2 occurs through the condensed-phase
reactions,
Reactions (R23)
and (R24) (Küpper et al.,
1998):
R23I-+H2O2↔HOI+OH-R24HOI+I-+H+→I2+H2O.
Nitrite ions react with hydronium ions to form the nitroacidium ion,
H2ONO+, which has been previously shown to produce I2
(Reactions R25–R27) (Hellebust
et al., 2007; O'Driscoll et al., 2006, 2008; O'Sullivan and Sodeau, 2010):
R25NO2-+H3O+↔HONO+H2OR26HONO+H3O+↔H2ONO++H2OR272H2ONO++2I-↔2NO+I2+2H2O.
Therefore, it is likely the I2 observed on connection of the flow tube
to CIMS originated from the above reactions, Reactions (R23)–(R27), as the pH ≤ 2
experiments in this work (IO3-5, SW3-5) favor these forward reactions.
At pH ≈ 4.7, frozen solutions without OH-radical precursors
produced no (IO6, SW6–SW7) or little (IO7, 0.10±0.06 nmol of
I2) amounts of molecular halogens above their respective limits of detection (LODs) after
activation of lights (Table 2). The small amount of I2 produced in IO7
possibly originates from the light and O2-mediated production mechanism
proposed by Kim et al. (2016) as summarized within Reactions (R13)–(R17). However, as
shown below, this mechanism of I2 production is of relatively minor
importance at this pH.
In the presence of H2O2 at pH ≈ 4.7, I2 mole
fractions increased rapidly upon irradiation, as shown in Fig. 2a. Of the
four experiments performed in these conditions (IO1, IO2, SW1, SW2),
three experiments (IO1, SW1, SW2) produced statistically similar amounts of
I2 (mean: 8±2 nmol) after 1 h of irradiation (Table 2).
The I2 signal behavior in Experiment IO2 qualitatively shared the same
features as Experiment IO1, SW1, and SW2 (Fig. S3) but provided an
apparently statistically different amount of I2 (0.6 (±0.4) nmol) based on the objectively chosen integration limits. This experiment is
discussed further in the Supplement.
Regarding other molecular halogens, IBr was observed above the estimated
limits of detection (3 pmol mol-1) upon irradiation during Experiment SW2 (Fig. 2a), starting
approximately 20 min before the addition of O3. No photochemically
produced (OH-induced) Br2 was unambiguously observed at this pH (note
that the apparent IO2 Br2 production of 0.034±0.003 nmol is
likely overestimated and is discussed in more detail in the Supplement). Cl2 mole fractions remained below limits of detection in
all cases with OH precursors at this pH.
pH ≤ 2
In cases without OH precursors at pH ≤ 2, photochemical I2
production was observed (integrated production of 14±10 nmol for
IO8, and 6±2 nmol for SW8) (Table 2), contrasting with experiments
performed at pH ≈ 4.7 in which very little was produced. This
production likely stems from the mechanisms outlined by Kim et al. (2016)
(Reactions R13–R17), which require only light and oxygen to form a charge-transfer
complex that results in I2 production (discussed in Sect. 1). Molecular
Br2 and Cl2 concentrations remained below limits of detection,
consistent with Abbatt et al. (2010), in which no Br2 or
Cl2 was observed without an OH precursor.
As discussed in Sect. 3.1, inclusion of H2O2 or NO2- can
result in direct oxidation of I- and reduce the available [I-] for
photochemical OH oxidation when pH ≤ 2. Photochemical production of
I2 across experiments yielded ≤0.82 nmol (IO4, IO5, and SW5) when
H2O2 was used as an OH precursor. However, when instead
NO2- was used (as in IO3, SW3, and SW4), initial observations of
I2 on connection of the flow tube to CIMS were as much as 90 % less than when
H2O2 was used (Table S1), thereby leaving more I- available
for reaction. For Experiment IO3 (using NO2-), the reduced pH led
to an observed photochemical I2 production amount of 39±1 nmol,
approximately 4 times larger than the largest amount observed at pH ≈ 4.7 (9±3 nmol; Table 2). That production would be
enhanced at lower pH was expected based on the halogen activation reactions, Reactions (R4)–(R22). The corresponding saltwater experiments using
NO2-
were not as conclusive; Experiment SW3 only yielded 4.0±0.1 nmol of
photochemical I2 (Fig. S5). Experiment SW4 (a repeat of SW3) did not
produce any photochemical I2 and qualitatively resembles the
H2O2 experiments performed at this pH. It is possible that, for
SW3 and SW4, more I2 was produced by dark reactions and flushed out of
the tube during connection with the CIMS and therefore would not have been
measured.
Photochemical production of Br2 does not appear until I2
production decreases. The results shown in Fig. 2a and b demonstrate
that when [I-]/[Br-] approximates the initial conditions of
Instant Ocean (∼2.6×10-3), OH-mediated I2
production precedes Br2 and IBr production (as in the pH ≈ 4.7 experiments and IO3, in which significant dark I2
production was not observed). After [I-]/[Br-] has sufficiently
decreased, Br2 eventually becomes the dominant photochemical product.
As demonstrated by Experiment IO4 (Fig. 2b and inset), there is a delay in
Br2 production until I- was removed as I2, then as IBr. For
experiments that used H2O2, photochemical Br2 yields
averaged 4.5±0.5 nmol between IO4 and IO5 and 6.0±0.7 nmol
from SW5. Experiment SW4 (using NO2-) produced a comparable amount
of Br2 (5.4±0.7 nmol). Given the initial depletion of I-
from dark I2 production (Sect. 3.1), we can estimate
[I-]/[Br-] at pH ≤ 2 in ice with H2O2 just before
irradiation based on the remaining moles of I- in solution (Table S1)
and the total moles of Br- in the solution. Averaging values from
Experiment IO4-5 and SW5, [I-]/[Br-] was calculated as (1.6±0.7)×10-4
(compared to the initial ratio of 2.6×10-3)
and was sufficiently low to result in photochemical production of Br2.
Photochemical Cl2 production was only observed from a frozen solution
of “pure” 0.56 M NaCl and H2O2 at pH = 1.8 (CL1), as shown in
Fig. 2c. The initial Br- impurity of this CL1 solution was determined
to be (4.5±0.3)×10-6 M via ion chromatography, while any
I- impurity concentration could not be detected above the 3σ
LOD of 90 nM. When the lights were turned on, slight increases in I2
and IBr were observed in concert with a rapid rise in Br2. After about
1 h of apparent equilibrium, I2 concentrations began decreasing,
while Br2, IBr, and Cl2 continued rising. Over 1 h of
illumination, 93±3 pmol of Cl2, 100±10 pmol of
Br2, and 100±10 pmol of I2 were observed.
However, as shown in Fig. 2c, the greatest rate of increase in Cl2
signal occurred just after this time. Integrating instead from t=0 until
t=2 h, the amount of Cl2 produced was 190±10 pmol, while
the amount of Br2 increased to 310±20 pmol. Utilizing the
starting halide concentrations of Br- and Cl- for CL1, our results
show Cl2 production was observed at [Br-]/[Cl-] of
8.1×10-6 (1/124000), compared to the Instant Ocean
[Br-]/[Cl-] of ∼1/800. Unfortunately,
BrCl could not be observed due to an unknown interference at m/z 241 and 243.
The observations in this study indicate competition for the OH radical in
which the most oxidizable halide is oxidized, and the corresponding
molecular halogens are produced until that halide ion is depleted in the ice
surface brine reaction environment. The trends in molecular halogen
production confirm acid-enhanced mechanisms in which the dominant products
are largely dependent on relative halide ratios. Here, Br2 and IBr were
not observed until I2 production sufficiently decreased the
[I-]/[Br-] ratio, and Cl2 was not observed unless the
[Br-]/[Cl-] ratio was sufficiently low ([Br-]/[Cl-]=8.1×10-6,
as discussed above). This observation is consistent with
Sjostedt and Abbatt (2008), who exposed
frozen salt solutions to gas-phase OH and found peak BrCl production
occurred as Br- decreased from an initial [Br-]/[Cl-] of
7.3×10-5. Additionally, Abbatt et al. (2010) generated condensed-phase OH
on frozen surfaces via the photolysis of nitrate and similarly found lower
Br2 and IBr integrated amounts at lower [Br-]/[Cl-] when
temperatures were higher than the eutectic point of sodium chloride. These
halide ratios are also consistent with in situ snowpack observations of
Br2, BrCl, and Cl2 formation (Custard et al., 2017;
Pratt et al., 2013).
Relative reactivities of OH-induced halogen production
I2, Br2, and Cl2 have been previously observed at mole
fractions within less than 2 orders of magnitude of each other in snowpack
interstitial air at Utqiaġvik, Alaska (Custard et al., 2017; Raso et al.,
2017). Custard et al. (2017) observed gas-phase [Br2]/[Cl2] values
for artificially irradiated, acidic snowpacks ranging from 2 to 95 for
corresponding snowpack [Br-]/[Cl-] ratios of (6±1)×10-4. Under similar conditions, Raso et al. (2017) observed
[I2]/[Br2] ranging from ∼0.4 to 0.8 from corresponding
snowpack [I-]/[Br-] amounts of (2.6±0.6)×10-3.
Despite the large differences in relative halide abundance (i.e.,
[I-] ≪ [Br-] ≪ [Cl-]), it appears
that halogen activation reaction kinetics favor the larger halide ions,
effectively leveling the relative molecular halogen production rates. The
observations herein provide an opportunity to explore the relative
reactivities of OH-mediated halogen production.
If we assume that the observed X2 flux out of the ice is proportional
to the production rate (i.e., X2 desorbs as it is produced, within the
residence time of the flow tube) and that halogen production is limited by
halide reaction with OH radicals, effective relative reactivities,
kX-/kY-, (where X and Y represent Br, Cl, or I) can be calculated
using Eq. (1).
FluxX2FluxY2=kX-[X-][OH][H+]kY-[Y-][OH][H+]
The initial molecular halogen flux is calculated as the integrated sum of
X2 (in moles) divided by both integration time (t=0–3 min,
starting from the beginning of irradiation) and the surface area of ice
coverage in the flow tube. Because the surface area, as well as the [OH] and
[H+] in the ice surface reaction environment, are identical within
individual experiments and cancel in these calculations, the relative fluxes
are simply equivalent to the relative outflow concentrations of halogens.
The pre-freezing halide ion concentrations (defined in Sect. 2) thus allow
us to solve for the effective relative reactivity, kX-/kY-, by
assuming the ratios of the halide ice concentrations are the same after
freezing.
At pH = 1.8, kBr-/kCl- was estimated to be (2.4±0.2)×105 from Experiment CL1; in other words, production of
Br2 is 240 000 times more efficient than production of Cl2 via (OH + halide) in the surface layer. Across the six experiments
performed at pH ≤ 2 (average of 1.85) using Instant Ocean (IO3, IO4,
IO5) and saltwater (SW3, SW4, SW5), kI-/kBr- was
calculated to average (9±4)×103 (reported uncertainty is the
standard error of the mean and thus only represents the experiment
repeatability). These relative reactivities are substantially larger than
the corresponding relative aqueous OH + halide rate constants
(kI-=kBr-=1.1×1010 M-1 s-1;
Buxton et al., 1988; Zehavi and Rabani,
1972;
kCl-=3.0×109 M-1 s-1; Grigor'ev et al., 1987), which are different by less
than a factor of 4. However, these rate constants refer to the specific
fundamental reaction of OH with X- to produce HOX-, as in Reaction (R8).
Ultimately, X2 production would occur via Reactions (R8)–(R12), and this
condensed-phase chemistry is much more complex when also considering
interhalogen reactions, such as Reaction (R28), that involve combinations of the three
molecular halogens, halides, and mixed molecular halogens (XY, where Y =Cl, Br, or I).
XY+X-→X2+Y-
Thus, it must be the case that there exist competing reactions that make the
production of the larger X2 more efficient. For example,
Cl+I-→ClI- may be faster than
Cl+Cl-→Cl2-.
Alternatively, the relative rates of the
disproportionation reaction, Reaction (R11), are likely different, favoring the larger
molecular halogens. We can thus only state from these observations that the
apparent relative reactivities calculated are consistent with the overall
reactivity of the larger ions compensating for their lower abundances. This
may lead to comparable production rates in our laboratory experiments and
comparable snowpack gas-phase concentrations.
The above relative reactivity calculations are considered upper limits since
the halide ratios used represent those in the pre-freezing solution. In
other words, it is assumed that the ions are excluded to the ice surface
reaction environment–air interface in amounts proportional to their
pre-freezing concentration. Malley et al. (2018) recently demonstrated that
brine can be distributed throughout ice in channels, suggesting that only
the solutes at the liquid–air interface (a fraction of the total
pre-freezing solution) participate in heterogeneous chemistry. Indeed, we
find evidence here suggesting not all ions are available for reaction at the
ice brine surface, particularly for experiments for which little I- was
lost from dark I2 production mechanisms (i.e., pH = 4.7 with OH
precursors: IO1, IO2, SW1, SW2). Considering Experiment IO2 as an example
(Fig. S5; pH = 4.7), integration of the I2 signal during
∼15 h of exposure to both light and O3 shows that
54 % (82 nmol) of the original 152 nmol of I- remained unreacted in
the frozen solution despite the signal apparently stabilizing at its
baseline. It is therefore probable that a significant number of the ions, as
well as H2O2, exist within brine channels within the ice (Bartels-Rausch
et al., 2014; Malley et al., 2018). Oxidation chemistry would then be
occurring throughout the ice, but release of molecular halogens to the flow
tube air would be determined by diffusion rates. The diffusion rates of the
product molecular halogens through bulk ice are likely slow, such that only
production occurring in the brine that is in the near-liquid–air interface
is observed here (Abbatt et
al., 2012). Of the halogens produced from frozen solutions here, it is
expected that I2 is observed most readily given the high polarizability
and surface affinity of I- in aqueous solutions
(Gladich et al., 2011) and the relative ease of
oxidation of I-. That is, surface concentrations will be relatively
enhanced with larger, more polarizable anions (I- > Br- >Cl-) (Gladich et al., 2011), which
favors production of I2 over Br2, and Br2 over Cl2. As
the larger and more reactive ions are depleted through oxidation, the next
largest ion then becomes more favorably oxidized. Thus, in addition to the
impact of differential reactivities and competing reactions for Reactions (R9)–(R12), what
we observe in the laboratory and in the field can also be influenced by the
relative surface enhancements of the anions, especially with respect to
O3 impacts as discussed below.
Effects of O3 on halogen production
In experiments without an OH source (IO6–IO8, SW6–SW8), I2 production
was greatest when O3 was introduced to the irradiated tube for both
pH regimes (Table 2). The amount of I2 produced over 60 min in
these experiments was large, ranging from 26±9 to 80±1 nmol at pH = 4.7 and from 2.6±1.7 to 38±12 nmol at
pH < 2. This production likely results from a combination of
heterogeneous recycling and the surface and aqueous reactions between
O3 and I- (k=2.0×10-12 cm3 molecules-1 s-1; Liu et al., 2001). While the I2
produced when pH < 2 appears to be lower, I2 had already been
produced in the presence of light prior to addition of O3 (Sect. 3.2.2), yielding a lower [I-]/[Br-] ratio when O3 was
eventually added. Br2 production amounts ranged from 0.012±0.001 to 0.16 ±0.01 nmol at pH = 4.7 and took up to 6 h
to rise above detection limits after O3 was added. At pH ≤ 2,
Br2 production amounts ranged from 0.14±0.02 to 0.93±0.05 nmol. While O3-mediated halogen production has been observed
directly from frozen surfaces in the absence of light in previous laboratory
studies (Artiglia
et al., 2017; Oldridge and Abbatt, 2011; Oum et al., 1998a; Wren et al.,
2013), Br2 production has not been directly observed from the Arctic
snowpack without irradiation (Pratt et al.,
2013). This raises a question of the role of O3 in initial halogen
release in the Arctic spring.
Normalized, background-subtracted HOX signals from Experiment IO2,
pH = 4.7. (a) Comparison of Br2 mole fractions to HOBr (m/z 225). Note that
the HOBr signal should be considered only qualitatively as its identity
could not be confirmed using isotopic ratios with m/z 223 due to its relatively
large background signal. (b) Effect of O3 on I2 and HOI.
Normalized, background-subtracted HOX signals from Experiment SW5,
pH = 1.8. (a) Comparison of Br2 mole fractions to HOBr (m/z 225). Note that
the HOBr signal should be considered only qualitatively as its identity
could not be confirmed using isotopic ratios with m/z 223. (b) Effect of O3
on I2 and HOI.
When OH precursors were present, the addition of O3 to the zero-air
flow over the irradiated frozen sample caused additional production of
I2 and Br2, as shown in Fig. 2a and b,
under both pH regimes (Table 2). In experiments at pH ≈ 4.7, in
which [I-]/[Br-] remained sufficiently large due to minimal dark
production of I2 (i.e., IO1-2, SW1-2), exposure to O3 caused a
sharp increase in I2 (as in Fig. 2a). I2 production amounts for
frozen Instant Ocean at pH ≈ 4.7 (IO1, IO2) averaged 22±10 nmol, about 2 times less than for frozen saltwater experiments
SW1 and SW2 (average production amount of 51±25 nmol). As the
I2 signal decayed, the corresponding Br2 signals gradually
increased above detection limits, approximately 3 h after the introduction of
O3 (Fig. 2a). The average integrated amounts of Br2 produced from
these pH ≈ 4.7 experiments were very similar (0.05±0.01 nmol for IO experiments and 0.03±0.01 nmol for SW experiments).
When pH < 2, the effects of O3 addition varied according to
the remaining availability of I-. When the surface I- reservoir
had been reduced from dark reactions with H2O2 or
NO2-
(Reactions R17–R21; Sect. 3.1), exposure to O3 did not increase I2 above the
LOD except in Experiment IO5, which exhibited a small spike before decaying
below the LOD (0.11±0.06 nmol in IO5). However, O3 did cause
additional Br2 production after 1 h (average of 10±2 nmol for IO4 (Fig. 2b) and IO5 (Fig. S4) and 14±2 nmol
for SW4 and SW5). In contrast, for SW3 (using NO2- as an OH
source), there was relatively little initial consumption of I- by dark
reaction; therefore, when O3 was added, 1.1±0.1 nmol additional
I2 was observed, comparable to what was observed with the higher
pH experiments (Fig. S5). The amount of Br2 produced (0.46±0.01 nmol) was also significantly less than observed when I- was initially
depleted, demonstrating the importance of the halide ratios.
This additional O3-induced halogen production could result from a
combination of mechanisms. First, as discussed above, O3 can react with
halides on frozen saline surfaces to produce Br2 or I2 per
Reactions (R18)–(R19) and then Reaction (R4) (Artiglia
et al., 2017; Carpenter et al., 2013; Gladich et al., 2015; Hayase et al.,
2010; Oum et al., 1998a; Shaw and Carpenter, 2013; Wren et al., 2013). It is
possible that Br2 (as well as other halogens) may have been produced
via this mechanism at levels below the LOD in previous Arctic snowpack
studies (Custard
et al., 2017; Pratt et al., 2013; Raso et al., 2017).
The presence of O3 also yielded HOX compounds (Figs. 3–4), likely formed
in the flow tube in part by O3 reactions with halides (Reactions R18–R19).
Additionally, given a flow tube residence time of 12 s, gas-phase
production of HOX is possible via Reactions (R1)–(R3) and could act as an additional
X2 production source (via Reaction R4), given a timescale for molecular
diffusion of 6.5 s for HOBr from the center of the tube to the ice
surface. At this flow rate, there is enough time for one–two heterogeneous
reaction cycles. Figure 3 shows HOX for IO2 (pH = 4.7 with
H2O2
present, analogous to IO1, SW1, SW2). For each experiment in this series,
increases in I2, HOI, and Br2 were readily observed when the
O3 was introduced at hour 2 (Figs. 3 and S4). However, corresponding
HOBr production was not observed, perhaps either due to a high LOD or the
relatively low abundance of Br2 that would limit production of HOBr.
Conversely, in pH ≤ 2 cases when substantial portions of I- had
already reacted prior to irradiation (IO4, IO5, SW4, SW5), the addition of
O3 produced negligible amounts of I2 and HOI (Fig. 4). But, in
these cases, following the addition of O3, HOBr (m/z 225
IHO81Br-), was observed together with Br2 (Figs. 4 and S4). We note in this case that m/z 223, representative of
IHO79Br-, does not appear to show an enhancement when O3 is added to the system.
There was a much higher background signal for m/z 223 compared with m/z 225
(IHO81Br-) resulting from an unknown interference.
Summary and conclusions
It was shown in this ice-coated-wall flow tube laboratory study that the
hydroxyl radical can act as an effective condensed-phase halide oxidant, leading to I2, IBr, Br2, and Cl2 production under acidic
conditions. Rates of molecular halogen production and release were dictated
by both pH and relative halide concentrations. The identities of the
molecular halogens produced appear to be highly influenced by which ions
are enhanced at the ice surface, with I2 production occurring prior to
Br2 production, which commenced as the [I-]/[Br-] was
reduced. An opportunity exists to further explore this chemistry via
surface-sensitive methods, for which recent developments have been shown to
effectively enable characterization of the surface composition of frozen
solutions of sodium chloride under near-atmospherically relevant conditions
(Artiglia et al., 2017; Orlando et
al., 2016). It would be useful to confirm the dominant ions involved in this
surface-based chemistry over time. Further investigations into the effects
of halide ratios on halogen production are also suggested, including
measurements of how the ratios vary for different frozen Arctic surfaces, as
well as how they vary spatially. While condensed-phase OH produces Br2
and I2 most rapidly in this study, it appears that other mechanisms,
such as heterogeneous recycling of HOCl or ClONO2, could be a more
dominant mechanism for in situ production of gas-phase Cl2 (Wang and Pratt, 2017). We find the addition of gas-phase O3
produces additional Br2 and I2, likely through aqueous reactions
with halides and/or gas-phase production of HOX or possibly XONO2
(Deiber et al., 2004) and subsequent
halogen explosion chemistry. These results lend support for the
photochemical, condensed-phase molecular halogen production mechanisms
proposed by the recent in situ snowpack experiments (Custard
et al., 2017; Pratt et al., 2013; Raso et al., 2017).
Understanding the environmental pH dependence of halogen activation
necessitates study of the pH on relevant Arctic frozen surfaces. Pratt et
al. (2013) found that the frozen surfaces
most conducive to in situ photochemical Br2 production had acidic pH
after melting, while no production was observed from those with a
well-buffered alkaline ice brine. Similarly, we find herein that
condensed-phase OH-induced halogen production is enhanced at lower pH. Wren
and Donaldson (2012a, b) found
in laboratory studies that pH of acidic and basic solutions remains
essentially unchanged after freezing and that saline solutions with buffers
(i.e., seawater) maintain their buffering capacity following trace gas
deposition, supporting the lack of observed Br2 production from the sea
ice surface (Pratt et al., 2013). Therefore, it would be useful to test
in situ production of halogens from Arctic frozen surfaces in tandem with
measurement of the pH of said surfaces to determine the atmospherically
relevant surface pH range required for halogen production.
Data availability
Data used for this work are available for download from the NSF Arctic Data Center under the dataset
“Laboratory experiments of the pH-dependent production of molecular chlorine, bromine, and iodine from frozen saline
surfaces”
(10.18739/A22804Z17; Halfacre et al., 2018).
The supplement related to this article is available online at: https://doi.org/10.5194/acp-19-4917-2019-supplement.
Author contributions
JWH and PBS designed the research, and JWH performed the experiments and
data analysis. All three authors contributed to the discussion and
interpretation of the results and writing of the paper.
Competing interests
The authors declare that they have no conflict of interest.
Acknowledgements
We thank the National Science Foundation for their funding (PLR-1417668 and
PLR-1417906, OPP-1417668). We also express thanks to Jonathan H. Slade, L. Gregory Huey, David J. Tanner, Fulizi Xiong, Angela R. W. Raso, and Kyle D. Custard for their
assistance with CIMS operation and maintenance. Additionally, we thank the
Purdue Chemistry Shop for helping build both the cooling and photolysis
boxes, as well as the Jonathan Amy Facility for Chemical Instrumentation for
their support in the fabrication of the experimental flow tube and setup of
our experimental boxes. Thanks are also extended to Megan Haas and Marianne Bischoff
for performing total organic carbon analysis of our samples and Angela R. W. Raso for confirmation of the iodide concentrations in our Instant Ocean
samples. Finally, we thank Timothy Miller and the Purdue Birck Nanotechnology
Center for the provision of the nanograde water used for our samples.
Review statement
This paper was edited by James Roberts and reviewed by two anonymous referees.
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