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Atmospheric Environment 75 (2013) 188e195

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Atmospheric Environment

journal homepage: www.elsevier .com/locate/atmosenv

Measuring a 10,000-fold enhancement of singlet molecular oxygen(1O2

*) concentration on illuminated ice relative to the correspondingliquid solution

Jonathan P. Bower, Cort Anastasio*

Department of Land, Air, and Water Resources, University of California, Davis, One Shields Avenue, Davis, CA 95616, USA

h i g h l i g h t s

� We report the first results for singlet molecular oxygen (1O2* ) on illuminated ice.

� 1O2* concentrations are up to 11,000 times higher on ice compared to liquid.

� Freezing increases 1O2* sources e without changing sinks e in liquid-like regions.

� We find abundant 1O2* formed in illuminated natural snow.

� Our results suggest 1O2* could be an important sink for certain pollutants in snow.

a r t i c l e i n f o

Article history:Received 25 January 2013Received in revised form16 April 2013Accepted 17 April 2013

Keywords:SnowPhotochemistryFreeze-concentrationLiquid-like regionsOxidants

* Corresponding author. Tel.: þ1 530 754 6095; faxE-mail address: canastasio@ucdavis.edu (C. Anasta

1352-2310/$ e see front matter � 2013 Elsevier Ltd.http://dx.doi.org/10.1016/j.atmosenv.2013.04.054

a b s t r a c t

Much attention has focused on the highly reactive hydroxyl radical in the oxidation of trace organiccompounds on snow and ice (and subsequent release of volatile organics to the atmospheric boundarylayer) but other oxidants are likely also important in this processing. Here we examine the ice chemistryof singlet molecular oxygen (1O2

*), which can be significant in atmospheric water drops but has not beenexamined in ice or snow. To examine 1O2

* on ice we illuminate laboratory ices containing Rose Bengal(RB) as the source of 1O2

* , furfuryl alcohol (FFA) as the probe, and Na2SO4 to control the total soluteconcentration. We find that the 1O2

*-mediated loss of FFA (and, thus, the 1O2* concentration) is up to

11,000 times greater on ice than in the equivalent liquid sample at the same photon flux. We attributethis large increase in the 1O2

* steady-state concentration to the freeze-concentration of solutes intoliquid-like regions (LLRs) in/on ice: compared to the initial solution, in the LLRs of ice the sources for 1O2

*

are highly concentrated, while the concentration of the dominant sink for 1O2* (i.e., water) remains largely

unchanged. Similar to results expected in liquid solution, rates of FFA loss in ice depend on both theinitial sensitizer concentration and temperature, providing evidence that these reactions occur in LLRs.However, we find that the enhancement in 1O2

* concentrations on ice does not follow predictions fromfreezing-point depression, likely because experiments were conducted below the eutectic temperaturefor sodium sulfate, where all of the salt should have precipitated. We also explore a method for sepa-rating 1O2

* and �OH contributions to FFA oxidation in laboratory ices and show its application to twonatural snow samples. We find that 1O2

* concentrations in these snows are approximately 100 timeshigher than observed in polluted, mid-latitude fog waters, showing that the enhancement of 1O2

* on ice isenvironmentally relevant and that 1O2

* could be a significant sink for electron-rich organic compounds insnow.

� 2013 Elsevier Ltd. All rights reserved.

1. Introduction

Polar and seasonal snowpacks provide a vast, photochemicallyactive reservoir for the cycling of chemicals between the landsurface and atmosphere, as well as to ground and surface waters if

: þ1 530 752 1552.sio).

All rights reserved.

the snowpack melts (Domine and Shepson, 2002; Grannas et al.,2007a; Hoffmann, 1996). When considering ice and snow photo-chemistry the hydroxyl radical (�OH) is usually considered thedominant oxidant because it reacts rapidly with most compoundsand it is ubiquitous in illuminated snow (Anastasio et al., 2007;Anastasio and Jordan, 2004; Beyersdorf et al., 2007; France et al.,2007). Snowgrain �OH is primarily formed by the photolysis ofhydrogen peroxide and, to a lesser extent, by photoreactions of

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195 189

nitrate and nitrite (Anastasio et al., 2013; Chu and Anastasio, 2003,2005; Dubowski et al., 2001; Honrath et al., 2000). Organic mattercan also be a source of �OH on illuminated snow and ice, as it is insurface waters (Grannas et al., 2006, 2004; Vione et al., 2006).

Chromophoric organic matter in surface and atmospheric wa-ters is also a source of singlet molecular oxygen (1O2

* ), which isformed by the transfer of energy from a photo-excited triplet stateof an organic compound to ground state dissolved molecular oxy-gen (Anastasio and McGregor, 2001; Faust and Allen, 1992; Haagand Hoigne, 1986; Haag et al., 1984; Zepp et al., 1977). While 1O2

*

is also probably formed from sunlit organic compounds in snowand ice, these reactions have not been examined. Although it ismore selective than �OH, 1O2

* reacts with a variety of electron-richorganic species, in some cases forming peroxides that can initiatesecondary free radical oxidation (Zepp et al., 1977). Singlet molec-ular oxygen can be a significant sink for electron-rich organics inboth surface and atmospheric waters (Anastasio and McGregor,2001; Zepp et al., 1981), although these reactions are in competi-tion with the physical deactivation of 1O2

* by water, which is thedominant sink in solution.

Given that 1O2* can be an important sink for some organic

compounds in surface waters and atmospheric drops, might it alsobe significant in/on illuminated ice and snow? When a solution isfrozen, the ordering of water molecules into ice crystals generallyexcludes solutes to highly concentrated liquid-like regions (LLRs) atthe aireice interface, at crystal grain boundaries, and as inclusionswithin the ice (Cho et al., 2002; Sadtchenko and Ewing, 2002; Weiet al., 2001; Wettlaufer, 1999). Because of this, solute concentra-tions within LLRs can be orders of magnitude higher than in thepre-frozen solution, which can significantly enhance the reactionkinetics for some photochemical reactions in the ice compared to inthe original solution (Grannas et al., 2007b; Kahan et al., 2010).

In this work we examine, for the first time, the dynamics of 1O2*

in/on illuminated ice samples, with two specific goals. The first is toinvestigate 1O2

* kinetics in illuminated laboratory samples todetermine whether 1O2

* concentrations are enhanced on icecompared to in the same sample studied as a liquid. Our secondgoal is to apply a method that uses selective sinks to distinguish 1O2

*

from other oxidants (e.g., �OH) in order to assess the importance of1O2

* as a sink for organic species in natural snow and ice.

2. Experimental methods

2.1. Materials

Furfuryl alcohol (FFA, 99%), Rose Bengal sodium salt (RB, �85%),2-nitrobenzaldehyde (2NB, 98%), sodium sulfate (SigmaeUltra),and sodium chloride (99.999%) were obtained from SigmaeAldrich.Acetonitrile (ACN, HPLC grade) and hydrogen peroxide (HOOH,certified ACS) were from Fisher, while deuterium oxide (D2O,99.95%) and glycerol (GlyOH, reagent ACS) were from Acros, andmethionine (Met, 99.5%) was from Fluka. All reagents were used asreceived. Solutions were made with purified water (Milli-Q) from aMilli-Q Plus system (�18.2 MU cm) with an upstream Barnstead B-Pure cartridge to remove organics.

2.2. Sample preparation and illumination for 1O2* kinetic

measurements (Goal #1)

The first goal of this work is to compare 1O2* kinetics in illumi-

nated ice and liquid samples. For this portion of the work we pre-pared air-saturated solutions with 5e50 nM of Rose Bengal (RB) (asthe source of 1O2

* ), 0.10e2.0 mM of furfuryl alcohol (FFA) (as theprobe for 1O2

* ), and 19e208.3 mM Na2SO4 (to control the concen-tration of total solutes, TS). RB is a sensitizer with a high,

wavelength-independent, quantum yield for 1O2* (F ¼ 0.75 (Allen

et al., 1991)). FFA is a commonly used probe for 1O2* in aqueous

solutions (Anastasio and McGregor, 2001; Faust and Allen, 1992;Haag and Hoigne, 1986; Haag et al., 1984) that reacts rapidly with1O2

* but is not expected to react with the RB triplet state (3RB*) sinceFFA is not known to react with triplet excited states (Halladja et al.,2007). Finally, we assumed complete dissociation of Na2SO4, whichyields total solute concentrations from 57 to 625 mM for our saltconcentrations.

Ice pellets were made by aliquoting 750 mL of solution into pre-cooled, custom made, PTFE ice pellet molds and allowing them tosit in a freezing chamber for 1 h at �10 �C. After freezing, pelletswere transferred to a temperature-controlled, illumination cham-ber at�10 �C; the chamber contains a quartz window (purged withdry air to eliminate frosting) to transmit light from the illuminationsystem. For liquid samples, solutions were transferred to 2-cm far-UV quartz cuvettes (Spectrocell), placed in the illumination cham-ber at 5 �C, and continuously stirred with a magnetic stir bar. Priorto illumination, all samples were allowed a minimum of 10 min toequilibrate to the chamber temperature.

Liquid water and ice samples were illuminated with light at549 nm (the peak absorption wavelength of RB) from a mono-chromatic illumination system (1000 W Hg/Xe lamp, Spectral En-ergy power source, lamp housing and monochromator). Comparedto the lifetime for RB in this system, which was w80 min (Bower,2012), the longest illumination time for the ice samples was short(�10 min). The majority illuminations of liquid RB solutions werecompleted in less than 15min. To achieve desired reaction times forice tests, six wire screens were placed in the light path to slowreactions to reasonable rates. For liquid tests no screens were used.For a given ice test, individual ice pellets of the same compositionwere illuminated for sequentially increasing amounts of time,melted in the dark at room temperature, and analyzed. For liquidtests, the sample was illuminated and small sub-samples wereconsecutively removed for analysis over the course of the experi-ment. After illumination, liquid samples were immediatelyanalyzed for FFA concentration.

2.3. Sample preparation and illumination for tests to distinguish1O2

* from other oxidants (Goal #2)

The second goal of our work is to develop a method to identifythe contribution of 1O2

* in samples (e.g., natural snows) where theoxidants for FFA are not necessarily known. To test this approach,we made a series of laboratory samples containing either RB (as asource of 1O2

* ), HOOH (as a source of �OH), both RB and HOOH, orneither (to test direct photodegradation of FFA). For samples with1O2

* we used RB concentrations of 25 nM for liquid samples and0.10 nM for ice samples. For hydroxyl radical (�OH) we used 100 mMHOOH in both liquid and ice. For liquid and ice samples we used50 nM of FFA as a probe for 1O2

* (Section 2.2) and �OH. We alsoperformed these tests on two samples of natural snow, where onthe day of each experiment, the snow was melted at room tem-perature, manipulated (e.g., by adding FFA), and refrozen in PTFE icepellet molds, just as with laboratory solutions. To distinguish thecontributions of 1O2

* and �OH to FFA loss for each sample type (e.g.,with RB), we perform four separate experiments (described inSection 3.3) in liquid and in ice.

Liquid and ice samples were prepared and illuminated asdescribed in Section 2.2, except that samples used in this portion ofthe work were illuminated with simulated sunlight (1000 W Xelamp, Spectral Energy housing, Newport power source, and opticalfilters described by Faust (1993)). The lifetime for RB in the SolarSimulator was approximately 20 min so we kept the total illumi-nation time to less than 10 min for both liquid and ice experiments.

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195190

2.4. Photon flux normalization

We used 2-nitrobenzaldehyde (2NB) actinometry (Galbavyet al., 2010) to account for differences in the photon flux betweenice and liquid experiments and to adjust for day-to-day fluctuationsin the output of the illumination systems. On each experimentalday we illuminated 10 mM 2NB samples in the same manner as theFFA experiments (e.g., for ice experiments with FFA, 2NB experi-ments were performed in similar ice pellets). The one exception isfor the monochromatic illumination system, where we illuminated2NB samples with 313 nm light to keep actinometry from beingprohibitively long. This method is appropriate since the photon fluxat 549 nm is proportional to that at 313 nm and since a relativephoton count is all that is needed for normalization (Bower, 2012).

2.5. Kinetic analysis of FFA loss

FFA concentrations were measured in melted (or unfrozen)samples using an isocratic highperformance liquid chromatography(HPLC) system consisting of a: Shimadzu SPD-10A UVeVis detectorand LC-10AT pump; 250 � 3 mm, 5 mm bead, BetaBasic-18 column(Thermo Hypersil-Keystone); eluent of 10% CH3CN/90% Milli-Q;flow rate of 0.6 mL min�1; detection wavelength of 230 nm; andinjection loop volume of 400 mL. 2NB was quantified on the sameHPLCwith amobile phase of 60%CH3CN/40%Milli-Qwith aflowrateof 0.6 mL min�1 and a detection wavelength of 258 nm.

Apparent-first order rate constants for loss of FFA (k0FFA) weredetermined from the slope of plots of ln([FFA]/[FFA]0) versus illu-mination time, where [FFA] and [FFA]0 are themolar concentrationsof FFA at time t and t¼ 0, respectively. First-order rate constants forloss of 2NB (j2NB) were determined in the same manner. Each FFAdecay rate constant was normalized to the daily value of j2NB:

k*FFA ¼ k0FFAj2NB

(1a)

where k*FFA is the normalized apparent-first order rate constant forFFA loss. In samples containing only RB (i.e., where 1O2

* dominatesthe observed FFA loss) the measured, apparent-first-order-rateconstant is proportional to the steady-state concentration of 1O2

*

by the second-order rate constant ðkFFAþ1O*

2Þ:

k0FFAzkFFAþ1O*2

h1O

*

2

i(1b)

where the equality is an approximation because we have notaccounted for the very small contribution of direct photo-degradation to FFA loss when illuminated with 549 nm light. Sta-tistical errors and error bars reported for results of Goal #1 and allactinometry experiments are the propagated standard error. Errorsand error bars for results of Goal #2 represent the 90% confidenceinterval calculated from the average slope of triplicate measure-ments, unless noted otherwise in the text.

Fig. 1. (A) First-order plots for loss of furfuryl alcohol in a liquid sample (1.0 mM FFA,5 �C) and ice sample (0.10 mM FFA, e10 �C). The inset shows details of FFA decay in theice sample. Both samples contained 250 mM total solutes (TS) as Na2SO4 and 10 nMRose Bengal (RB) and were illuminated with 549 nm light. (B) Calculated first-orderrate constants for loss of FFA (normalized to photon flux, which results in unitsof s�1/s�1 (Equation (1a))) calculated from the data in panel (A). The rate constant forFFA loss on ice is 11,300 (2200) times higher than the value in the liquid sample.

3. Results & discussion

3.1. 1O2* in ice and liquid samples

Our goal in this section is to examine how 1O2* concentrations in/

on illuminated ice compare to values in the same solution studied asa liquid. As described in Section 2.2, we estimate 1O2

* concentrationsbymonitoring the loss of FFA in illuminated samples containingRoseBengal (RB) as the source of 1O2

* : themeasured rate constant for FFAloss, k0FFA, is proportional to the singlet oxygen steady-state con-centration (Equation (1b)). An example of these results is shown in

Fig. 1A for samples containing 10 nM RB that were illuminated with549 nm light. The figure shows that loss of FFA in the ice sample isdramatically enhanced compared to in solution: the loss occurringon ice (�10 �C) in only 45 s is greater than that in the equivalentliquid (5 �C) sample after w11 h of illumination. Converting thesekinetic plots to photon-normalized apparent first-order rate con-stants for FFA loss (i.e., k*FFA; Section 2.5) shows that FFA loss on ice is11,300 times faster than in solution (Fig. 1B). To complement our549 nm experiments, we also conducted several tests where weilluminated samples with simulated sunlight. The loss of 50 nM FFAin/on ice containing 0.10 nM RB is 35 times faster than a liquidsample containing 25 nMRB (Supplemental Information, Figure S1),which corresponds to an ice-phase enhancement of the 1O2

* con-centration bya factor of 8600 after correcting for the difference in RBconcentration between liquid and ice.

Why is the concentration of 1O2* enhanced in/on illuminated ice

containing Rose Bengal? To explore this question, we first considerthe freeze-concentration effect. Based on freezing point depression(FPD), in a frozen solution the fraction of water molecules that arepresent in liquid-like regions in/on the ice (f(T)) depends inverselyon temperature and directly on the total solute concentration of thepre-frozen solution ([TS]LIQ) (Cho et al., 2002):

f�T�z

mH2ORTf1000DH0

f

T

T � Tf

!½TS�LIQ (2)

where mH2O is the molecular weight of water, R is the ideal gasconstant, Tf is the freezing point of purewater, T is the experimentaltemperature, andDH0

f is the enthalpy of fusion of water. The inverseof the fraction of initial water molecules present in LLRs is thefreeze-concentration factor (F), i.e., the ratio by which solute con-centrations in LLRs are enhanced relative to in the unfrozen solu-tion (assuming that all of the solutes are segregated to LLRs in thefrozen sample):

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195 191

F ¼ ½TS�LLR½TS� z

1fðTÞ (3)

Fig. 2. Dependence of furfuryl alcohol (FFA) decay on (A) concentration of 1O2* sensi-

tizer (Rose Bengal, RB), (B) temperature, and (C) total solute (TS) concentration. Unlessnoted otherwise, ice samples were illuminated at �10 �C with 549 nm light and weremade from solutions containing 2.0 mM FFA, 250 mM total solutes (as Na2SO4), and10 nM RB. In panel (A) values of k*FFA (in units of s�1/s�1) are 0.00024 (No RB), 0.13(5 nM RB), 0.22 (10 nM RB), and 0.80 (50 nM RB). In panel (B) values of k*FFA (s�1/s�1)are 0.24 (5 �C), 0.22 (�10 �C), and 0.14 (�20 �C). In panel (C) all samples converged toan apparent rate constant of k*FFA ¼ 0.19 s�1/s�1.

LIQ

Freezing-point depression (Equation (2)) predicts a concentra-tion factor (Equation (3)) of approximately 21,000 for ice at �10 �Cmade from a solution containing 250 mM TS, which is in goodagreement with our observation in the 549 nm experiments (i.e.,the 11,300-fold enhancement in k*FFA on ice compared to in solu-tion; Fig. 1B). The enhancement of k*FFA on ice compared to solutionin the solar simulator experiments (Fig. S1) is similarly high (8600),though 2.4 times lower than F predicted from freezing-pointdepression. Although natural snows are not formed in the samemanner as our ice pellets, there is good evidence that many solutesare concentrated in liquid-like regions in/on natural snow (e.g.,Domine et al. (2008), Mulvaney et al. (1988), etc.).

Our measured enhancements of 1O2* on ice are likely lower

bounds because the concentration of FFA in the liquid-like regionsof our ice samples is sometimes high enough that FFA might be animportant sink for 1O2

* ; that is, the fraction of 1O2* that reacts with

FFA (f FFA; Supplemental Information Section S1) is sufficiently highto depress the steady-state concentration of 1O2

* . For example, usingfreezing-point depressionwe estimate that [1O2

* ] on ice at �10 �C issuppressed by a factor of 1.4 for samples made from solutionscontaining 0.1 mM FFA and 250 mM TS (i.e., the conditions in Fig. 1A(inset)) and by a factor of 2.4 for samples made from solutionscontaining 50 nM FFA and 30 mM TS (i.e., conditions in Fig. S1).When we consider this suppression of 1O2

* by FFA, the measuredenhancement of k*FFA from Fig. 1B and S1 are very close (<25%) to Fcalculated with freezing-point depression.

To further explore the kinetics of 1O2* on ice, we examined k*FFA

in ice samples as a function of sensitizer concentration, tempera-ture, and the total solute concentration of the pre-frozen solution.Fig. 2A shows that the rate constant for FFA loss in/on ice increaseswith sensitizer concentration and that there is no loss of FFA in 549-nm-illuminated ice that does not contain RB. Similarly, FFA loss wasnot observed in the absence of RB in liquid samples illuminated at549 nm. While increasing RB increases the FFA loss rate constant(i.e., the 1O2

* concentration), k*FFA in/on ice is not proportional to theRB concentration (Fig. S2). For example, increasing the RB con-centration by a factor of 10 (from 5 to 50 nM) only increases k*FFA bya factor of 5.9 0.6 (1 SE). One possibility for the observation thatFFA loss is not as sensitive to sensitizer concentrations as we expectcould be that at higher concentrations of RB, some of the RBmay beprecipitating or aggregating in/on ice. However, the increasing lossof FFAwith increasing RB indicates that RB is collocated with FFA inthe ice pellets, most likely in liquid-like regions.

We next examined how temperature affects 1O2* kinetics in/on

ice. Based on freezing-point depression, the value of F in/on iceshould increase with decreasing temperature: for example, coolingfrom �5 to �10 �C should double F (Equations (2) and (3) and thusdouble [1O2

* ] and k*FFA. However, we find the opposite temperaturedependence, though with a much weaker response: for example,moving from �5 to �10 �C decreases k*FFA, but only by 7% (Fig. 2B).While this behavior is not what we would expect from freezing-point depression, it is consistent with the temperature depen-dence of the second-order rate constant for the solution reaction ofFFA with 1O2

* ði:e:; kFFAþ1O*

2Þ. As shown in Fig. S3, using our data

from Fig. 2 the activation energy for FFA loss on illuminated icecontaining RB is 19 4 kJ mol�1 K�1, quite similar to the value of Eadetermined for kFFAþ1O*

2in aqueous solution (22.7 kJ mol�1 K�1;

Gottfried and Kimel (1991)).We also quantified FFA loss as a function of total solute (TS)

concentration in the pre-frozen solution. As described in Equation(3), if freezing-point depression controls the freeze-concentrationfactor of liquid-like regions in our Na2SO4eH2O ices then

decreasing total solute concentration from 625 to 57 mM shouldincrease F by a factor of 11. However, as shown in Fig. 2C, the rateconstant for FFA loss is insensitive to changes in Na2SO4 concen-tration in the tested range. While this effect may in part be due tosuppression of 1O2

* by FFA (as discussed above), we also attributethis result to the fact that our experiments were conducted attemperatures below the eutectic temperature for aqueous Na2SO4(�1.3 �C; Hall and Hamilton (2008)) and thus the majority ofNa2SO4 in our samples should have precipitated in/on the ice.However, the dependence of k*FFA on Rose Bengal concentrationand on temperature (Fig. 2A and B) suggests that the 1O2

* reactionsare occurring in liquid-like regions of the ice, even though theseLLRs are not behaving as predicted from freezing-point depression.This is consistent with previous studies that have found evidence ofliquid-like regions in ices below the eutectic temperature (Choet al., 2002; Grannas et al., 2007b).

3.2. Mechanism for the enhancement of 1O2* concentrations on ice

Our results above show that 1O2* concentrations are greatly

enhanced on ice, by approximately the freeze-concentration factorF predicted by freezing-point depression. Here we explore the

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195192

mechanism for this enhancement. We start by considering a short-lived oxidant (Ox) whose steady-state concentration is determinedby the balance of formation and loss:

½Ox� ¼ Rf ;OxPkOxþm½m� ¼

Rf ;Oxk0d;Ox

(4)

Here Rf ,Ox is the rate of formation of the oxidant, kOxþm is thesecond-order rate constant for reaction of Ox with sink “m”, andk0d,Ox is the overall first-order rate constant for loss of Ox (i.e., thesum of the apparent first-order rate constants from each losspathway of Ox). In the case of singlet oxygen in illuminated solutionwe can write this equation as

h1O

*

2

iLIQ

¼ jRB;LIQ ½RB�LIQk1O*

2þH2O;LIQ½H2O�LIQ þP k1O*

2þm;LIQ ½m�LIQ

¼ jRB;LIQ ½RB�LIQk0H2O;LIQ

(5)

where jRB is the rate constant for formation of 1O2* during illumi-

nation of RB, k1O*

2þH2Ois the second-order rate constant for deacti-

vation of 1O2* by water, and the concentrations of RB, H2O, and sinks

m (including FFA) are explicitly denoted in liquid phase. The productk1O*

2þH2O[H2O]LIQ is the apparent first-order deactivation rate con-

stant for 1O2* by water ði:e:; k0H2O

Þ. Since quenching by water is the

dominant sink for 1O2* in solution ði:e:; k0H2O

>>P

k1O*

2þm½m�LIQ Þ,the denominator simplifies to k0H2O

.Then if the concentration of FFA (or any other reactive solute) is

sufficiently low to remain a minor sink for 1O2* in an ice sample, the

steady-state concentration of 1O2* in the LLRs is:

�1O2�LLR ¼ jRB;LLR½RB�LLR

k1O2þH2O;LLR½H2O�LLR

¼ jRB;LLR½RB�LLRk0H2O;LLR

(6)

The case where FFA is an important sink for 1O2* is discussed in

the Supplemental Information (Section S1). RB in the ice should befound predominantly in liquid-like regions, where its concentra-tion is enhanced by a factor of approximately F (the freeze-concentration factor) relative to the liquid value. In contrast, theconcentration of liquid (or liquid-like) water, which is a dominantsink for 1O2

* , remains essentially unchanged, based on ourmodelingof ice conditions using the E-AIM model (Wexler and Clegg, 2002);thus k0H2O; LLR

z k0H2O; LIQ. Making these substitutions into Equation

(6) yields:

h1O

*

2

iLLR

zjRB;LIQ ½RB�LIQF

k0H2O;LLRz

Rf ;1O*

2;LIQF

k0H2O;LIQ¼h1O

*

2

iLIQ

F (7)

where Rf ; 1O*

2; LIQis the rate of 1O2

* formation in solution. Equation(7) assumes that the temperature dependence for a rate constant inLLRs is similar to the value in liquid, consistent with Fig. S3, andholds when jRB,LLR[RB]LLR z jRB,LIQ[RB]LIQF. Thus, we expect con-centrations of 1O2

* in LLRs to be higher than those in the equivalentliquid solutions by a factor of approximately F, which is roughlywhat we see in our results (Fig. 1 and S1). A similar treatmentshowing why we do not expect the �OH concentration to beenhanced in LLRs is provided in the Supplemental Information(Section S2): because both the sources and sinks of �OH areenhanced by a factor of F in LLRs compared to in solution, its steady-state concentration should be very similar in solution and ice.

3.3. Distinguishing 1O2* from other oxidants in/on ice

The enhancement of 1O2* concentrations in frozen samples

suggests that singlet oxygen could be a significant oxidant forelectron-rich organic compounds in natural snow and ice, whereother oxidants, like �OH, will also be present. As shown in Section3.1, FFA is a useful probe for 1O2

* on ice, but it also reacts with �OH,which needs to be taken into account in order to determine the 1O2

*

concentration. In liquid samples D2O dilution is typically used todistinguish the 1O2

* contribution to FFA loss (Anastasio andMcGregor, 2001; Bilski et al., 1997; Faust and Allen, 1992). Butthis approach is problematic for snow samples because it requires alarge dilution of the original sample (typically by a factor of two),which both reduces the concentration of 1O2

* sensitizers in thesnow sample and potentially introduces contaminants.

Our second goal of this paper is to try an approach that usesselective sinks to differentiate the reactions of 1O2

* from �OH in/onice. First, we note that the measured rate of FFA loss (Rd,FFA) in asample is the sum of the rates of loss from all pathways, includingreaction with 1O2

* ðRFFAþ1O*

2Þ, with �OH (RFFAþ�OH), and via direct

photodegradation (RFFAþhv):

Rd;FFA ¼ RFFAþ1O*2þ RFFAþOH þ RFFAþhv (8a)

Factoring out the concentration of FFA from each term we candescribe the system in terms of photon-flux normalized, apparent-first order rate constants for FFA loss:

k*FFA ¼ k*FFAþ1O*

2þ k*FFAþOH þ k*FFAþhv (8b)

In order to determine the contribution of 1O2* to FFA loss in a

sample we need to identify the contributions from the otherpathways. In order to do this, we perform four separate experi-ments on the same laboratory sample, which contains either asource of 1O2

* (RB), a source of �OH (HOOH), both, or neither. In thefirst experiment, termed “Base”, we measure k*FFA in a laboratorysolution prepared to simulate solute levels in a remote continentalsnow, using 8 mM glycerol as a proxy for natural �OH scavengersand 7.3 mM Na2SO4 to set the total solute concentration at 30 mM.In the second experiment, “Ions”, we add 200 mM TS (i.e.,66.67 mM Na2SO4) to the Base sample to test the effect of addedsolutes on FFA loss, which serves as a total solute control forsubsequent experiments. The third experiment is “GlyOH” inwhich we add 200 mM glycerol to the Base sample; this shouldsuppress �OH (which reacts rapidly with glycerol) but not affect1O2

* (which does not react with glycerol (Müller-Breitkreutz et al.,1995). Thus, glycerol addition should yield a value of k*FFA that isapproximately the sum of FFA loss due to 1O2

* and direct photo-degradation (k*FFA z k*

FFAþ1O*

2

þ k*FFAþhv). The fourth and final

experiment is “Met” in which we add 200 mM methionine to theBase sample to suppress both 1O2

* and �OH; methionine reactsrapidly with both 1O2

* and �OH (Supplemental Table 1). Thus in theMet experiment we expect that FFA loss is due only to directphotodegradation (i.e., k*FFA z k*FFAþhv).

An application of this method is shown in Fig. 3; the resultsare summarized here, while a full discussion is in theSupplemental Information (Section S3). Results in liquid samplesare shown in Row A of Fig. 3. FFA loss in illuminated solutions notcontaining RB or HOOH is low (see “hv-only” in Column I),especially compared to solutions with oxidants (Columns II e IV).In solutions with RB or HOOH the addition of glycerol ormethionine qualitatively acts as expected: glycerol does notaffect 1O2

* (Fig. 3A.II) but reduces the contribution of �OH to FFAloss (Fig. 3A.III), while methionine suppresses both 1O2

* and �OH.

Fig. 3. Photon-flux-normalized apparent first-order rate constants for FFA loss in laboratory solutions studied as liquid (top row, A) and ice (bottom row, B) during illumination withsimulated sunlight. Column I shows FFA loss in samples with illumination only (i.e., no added oxidants); column II shows FFA loss in samples containing Rose Bengal, a photo-chemical source of 1O2

* ; column III shows k*FFA in samples containing 100 mM HOOH as a photochemical source of �OH; and column IV shows samples with both 1O2* and �OH. In

systems containing RB, liquid samples contained 25 nM RB while the ice samples contained 0.10 nM. The “Base” sample contains 50 nM FFA, 8 mM glycerol, 7.33 mMNa2SO4 (to makethe total solute concentration of 30 mM), in addition to RB and/or HOOH as noted. Experiments labeled “Ions” contain an additional 200 mM of ions to the Base sample (as 66.67 mMNa2SO4). “GlyOH” experiments include the addition 200 mM of glycerol, an �OH sink, to the Base sample, while “Met” includes the addition of 200 mM methionine, a sink for 1O2

* and�OH, to the Base sample. Note that the y-axis scale for the ice samples is 50 times greater than that for the liquid samples.

Fig. 4. Effect of D2O on FFA reactivity in liquid and ice Base samples (50 nM FFA, 8 mMglycerol, and 7.33 mM Na2SO4). Panel (A) shows FFA loss in samples with no addedoxidant (k*FFAþhv), while panel (B) shows FFA loss ðk*

FFAþ1O*2

Þ in samples with 25 nM RB(for liquid) or 0.10 nM RB (for ice). Samples were prepared in H2O, or a 50:50 mixtureof D2O and H2O, and illuminated with simulated sunlight at 5 �C for liquid and �10 �Cfor ice. Error bars are 1 SE propagated from the uncertainties in kFFA and j2NB; errorbars for liquid results in H2O in both panels are too small to discern on this scale.

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195 193

Similarly, in solutions with both RB and HOOH (Fig. 3A.IV), theaddition of glycerol appears to remove the �OH contribution toFFA oxidation, while not affecting 1O2

* , while methionine sup-presses both 1O2

* and �OH. The addition of salt in the “Ions” testshas no effect in liquid samples, as expected.

Row B of Fig. 3 shows the application of this method to icesamples. Unfortunately, our analysis on ice is confounded by therapid loss of FFA in the “hn only” experiments with simulatedsunlight (Fig. 3B.I): this loss is nearly as fast as in samples con-taining oxidants and is 600 times faster than in the “hn only” liquidexperiments. As discussed in Supplemental Section S3, our labo-ratory samples apparently contained a low concentration ofcontaminant that forms 1O2

* . While the contaminant also appears tobe in the liquid samples (Section S3), the freeze-concentration ofsolutes into LLRs in/on ice intensifies the effect of the contaminanton ice. We verified the presence of 1O2

* in the “hn only” experimentsby measuring FFA loss in liquid and ice prepared with a 50:50mixture of D2O and H2O (Fig. 4A). Because D2O deactivates 1O2

* lessrapidly than does H2O (Supplemental Table S1), if 1O2

* is thedominant oxidant for FFA k*FFA should be 1.9 times faster in theD2O/H2O mixture relative to in pure H2O (Supplemental SectionS4). Consistent with this, we observe increases of 1.9 0.63(1 SE) in liquid and 2.0 0.44 (1 SE) on ice (Fig. 4A), suggesting 1O2

*

is the dominant sink for FFA in these samples, even though nosource for 1O2

* was added. Similarly, in a D2O/H2O mixture to whichwe add a source of 1O2

* (i.e., RB), FFA loss is enhanced by 2.1 0.51(1 SE) in liquid and 2.5 0.63 (1 SE) in/on ice, both consistent with1O2

* as a dominant sink for FFA.While our experiments in this section confirm that 1O2

* isenhanced on ice compared to in solution, because of this tracecontaminant, we cannot quantitatively use our technique todistinguish the 1O2

* contribution to FFA loss in laboratory ice sam-ples. However, the expected reduction in k*FFA as a result of addingsinks to the mixed oxidant system in/on ice is comparable tothe liquid results (Fig. 3, Column IV), showing that the addition ofglycerol and methionine at least qualitatively suppress 1O2

* (and�OH) in/on ice as they do in solution. Thus, we can use the differencebetween results of the GlyOH and Met experiments to estimate aconservative value of k*

FFAþ1O*

2

and thereby the concentration of 1O2*

in ice samples.

3.4. Application to natural snow samples

In our final experiments we applied ourmethod for determining1O2

* to two snow samples to examine whether appreciable 1O2* is

formed. We tested one surface snow sample from the Central SierraSnow Laboratory in Soda Springs, California and one from the cleanair sector at Summit, Greenland. As shown in Fig. 5, in these naturalsamples (with no added oxidants), the loss of FFA in the Baseexperiment is comparable to the laboratory ice samples containing

Fig. 5. Photon-flux-normalized apparent first-order rate constants for FFA loss in icesamples made from snow collected in the Sierra Nevada Mountains of California and atSummit, Greenland. For a given sample, the snow was melted, divided into three ali-quots, 50 nM furfuryl alcohol (and, in some cases, glycerol or methionine) was added,and then the samples were refrozen as ice pellets and illuminated with simulatedsunlight at �10 �C. Samples labeled (A) contain only the addition of FFA and areanalogous to the “Base” case in the laboratory samples. Samples labeled (B) contain FFAand 200 mM GlyOH, while samples labeled (C) contain FFA and 200 mM Met.

J.P. Bower, C. Anastasio / Atmospheric Environment 75 (2013) 188e195194

0.1 nM RB. FFA loss is suppressed in the GlyOH experiment, prob-ably because this decreases F (by increasing the total solute con-centration) and thus decreases the 1O2

* concentration in liquid-likeregions. We did not have enough sample volume to perform Ionsexperiments (i.e., with the addition of Na2SO4) on these samples toconfirm this. But it is unlikely that the GlyOH result represents the�OH contribution to FFA loss since 100 mM HOOH (which is at theupper end of snow concentrations) produces a negligible �OHcontribution on this scale (Fig. 3). The addition of methionine re-duces FFA loss more than GlyOH, suggesting that even when sup-pressed by a decrease in F, 1O2

* is still important. We use thisdifference between the GlyOH and Met experiments to conserva-tively estimate the singlet oxygen concentration (SupplementalSection S5).

Our estimates for [1O2* ] in these samples are (2 1) � 10�11 M

for the Sierra snow and (2 0.3) � 10�11 M for the Summit snow,with both results normalized to winter solstice sunlight in Davis,California (Supplemental Section S5). Even though these snowsamples are much cleaner than Davis fog waters (e.g., organic car-bon concentrations in surface snow at Summit are at least 100times lower (Anastasio and McGregor, 2001; Hagler et al., 2007)),our estimated 1O2

* concentrations in the snows are approximately100 times higher than in illuminated fog waters (1.1e6.1)� 10�13 M(Anastasio and McGregor, 2001). These preliminary 1O2

* concen-trations in snow suggest that singlet oxygen can be an importantoxidant on illuminated natural snow: with 10�11 M 1O2

* we calcu-late lifetimes on the order of minutes for tetrabromobisphenol A toa few hours for phenol based on aqueous rate constants (Wilkinsonet al., 1995). However, these calculations require that the pollutantsand precursors for 1O2

* are collocated in/on ice, which may not bethe case. In addition, it is not clear whether these field samplessuffer similar contamination as observed in laboratory samples, sothese 1O2

* concentrations should be regarded with caution. How-ever, our results of elevated 1O2

* in illuminated natural snows couldhelp explain the enhanced loss of pollutants in ice compared to inliquid. For example, Rowland et al. (2011) found that the loss ofaldrin (an unsaturated, organochlorine insecticide) was twice asfast in illuminated, refrozen solutions made from Arctic snowcompared to in the melted snow solution (Rowland et al., 2011).Since 1O2

* generally reacts rapidly with unsaturated compounds(Wilkinson et al., 1995), this observed enhancement in aldrinreactivity on ice might in part be due to oxidation by 1O2

* .

4. Conclusions

Reactions in the snowpack can release volatile compounds tothe overlying atmosphere, including NOx, HONO, Br2, and volatileorganic compounds (VOCs) such as formaldehyde and acetalde-hyde (Grannas et al., 2007a). While some of the VOC release is duesimply to thermal reactions, a portion appears to be linked to VOCformation on snowgrains via direct and indirect photooxidation oforganic compounds. The indirect degradation of organics in snow istypically attributed to reactions with hydroxyl radical, but our re-sults suggest that singlet molecular oxygen may also be important.

More specifically, we have found that the steady-state concen-tration of singlet molecular oxygen is approximately 104-timeshigher in illuminated ice compared to in the same sample studiedas liquid. This enhancement arises because the concentration of the1O2

* source (Rose Bengal in our laboratory ices) is enhanced inliquid-like regions in/on the ice but the concentration of the majorsink for 1O2

* (liquid H2O) is not. For oxidants where both the sourceand sink are enhanced in LLRs compared to in the precursor solu-tion (e.g., for �OH), we do not expect a concentration enhancementon ice. The enhancement of 1O2

* is insensitive to the total solutelevel when using Na2SO4 and is only weakly dependent on tem-perature in ice below the Na2SO4 eutectic temperature. Since totalsolute concentration and temperature both affect the freeze-concentration factor of LLRs in/on ice based on freezing-pointdepression, our results suggest that at temperatures below theeutectic there are still liquid-like regions in/on ice, but that most ofthe Na2SO4 has precipitated.

We also sought to develop a scavenger technique that coulddiscern the contributions of 1O2

* and �OH to the decay of a furfurylalcohol probe. While this method shows some promise, it iscomplicated by some unexplained behaviors as well as by highbackground 1O2

* levels in ice samples with no added sensitizer.Determining the importance of 1O2

* in/on ice to pollutant cyclingwill require a more robust technique. But using our current methodindicates that 1O2

* concentrations in natural snows are approxi-mately 100 times higher than in polluted, mid-latitude fog waters,which suggests that singlet oxygen is a significant sink for someelectron-rich organic species in/on illuminated ice.

Acknowledgments

This research has been supported by a graduate fellowship to J.B.from the U.S. Environmental Protection Agency’s Science to AchieveResults (STAR) program (grant # FP916999). Although the researchdescribed in the article has been funded wholly or in part by theU.S. EPA, it has not been subject to any EPA review and thereforedoes not necessarily reflect the views of the Agency, and no officialendorsement should be inferred. We are also grateful for fundingfrom the National Science Foundation (grants ANT-0636985 andCHE-1214121). We extend special thanks to Shauna McKellar forpreliminary laboratory work, and to Katarina Makmuri and ErinRausch VanHouten for laboratory assistance. We also thank HarryBeine and Bill Casey for comments that improved the manuscript.

Appendix A. Supplementary data

Supplementary data related to this article can be found at http://dx.doi.org/10.1016/j.atmosenv.2013.04.054.

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