[Advances in Chemistry] Aquatic Humic Substances Volume 219 (Influence on Fate and Treatment of Pollutants) || Aquatic Humic Substances as Sources and Sinks of Photochemically Produced Transient Reactants

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  • 23 Aquatic Humic Substances as Sources and Sinks of Photochemically Produced Transient Reactants

    Jrg Hoign, Bruce C. Faust1, Werner R. Haag2, Frank E. Scully, Jr.3, and Richard G. Zepp4

    Swiss Federal Institute of Water Resources and Water Pollution Control (EAWAG), 8600 Dbendorf, Switzerland

    In sunlit surface waters aquatic humic substances and nitrate act as sensitizers or precursors for the production of photoreactants such as singlet oxygen, humic-derived peroxy radicals, hydrogen peroxide, solvated electrons, and OH radicals. Lifetimes of the various reac-tants are controlled by their reactions with aquatic humic substances (OH radicals), by solvent quenching (singlet oxygen), by reactions with molecular oxygen (solvated electron), or by other processes (per-oxy radicals). The steady-state concentration of each transient formed during solar irradiation was determined from the apparent first-order disappearance rate of added organic probe compounds. The probe compounds used had selective reactivities with the individual tran-sient species of interest. Effects of the photoreactants on the elimi-nation of micropollutants and on chemical transformations of DOM are discussed.

    1Current address: School of Forestry and Environmental Sciences, Duke University, Durham, NC 27706

    2Current address: SRI International, Menlo Park, CA 94025 3Current address: Department of Chemical Sciences, Old Dominion University, Norfolk, VA 23508-8503

    4Current address: Environmental Research Laboratory, U.S. Environmental Protection Agency, Athens, GA 30613

    0065-2393/89/0219-0363$07.00/0 1989 American Chemical Society

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    In Aquatic Humic Substances; Suffet, I., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1988.

  • 364 AQUATIC HUMIC SUBSTANCES

    DISSOLVED ORGANIC MATERIAL IN SURFACE WATERS HAS A ROLE in pro

    ducing or consuming different types of photoreactants. The conclusions stated in this chapter are based on data from a series of our recent publications. Extensive literature reviews are given in these publications and are not repeated here. This chapter wil l focus on the reactions for which humic materials act as sources or sinks. We make no attempt to include all other possible photochemical processes. For example, no discussion of heterogeneous processes is included, although there is evidence that they are important (e.g., in the redox cycling of metals). Moreover, photochemical processes mediated by superoxide and hydrogen peroxide are discussed in Chapter 22 by Cooper et al.

    During a cloudless summer noon hour, surface waters receive approximately 1 k W / m 2 of sunlight, or about 20 einsteins/m 2 (20 mol of pho-tons/m 2) (Figure 1). Within 1 year about 1300 times this dose is accumulated (2). A large portion of these photons is absorbed by dissolved organic material (DOM) present in natural water. In addition, a rather small fraction of short-wavelength light is absorbed by nitrate (Figure 2).

    From a chemist's viewpoint, the resulting rate of interactions between photons and absorbers is very high. Assuming that most of the photons are absorbed in a well-mixed 1-m water column, we estimate that about 20 mmol/(L*h) of interactions occur between photons and absorbing sub-

    Figure 1. Sofor radiation, (a) Mean dose intensity in a mixed 1-m water column in which all light is absorbed, (b) Monthly solar flux (280 < < 2800 nm) in

    Dubendorf(47.5 N\ 1985.

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    In Aquatic Humic Substances; Suffet, I., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1988.

  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 365

    (nm)

    Figure 2. Decadic molar absorptivities of Greifensee DOM and N03~ anion. The values for N03~ have been multiplied by a factor of 10. Typical DOC and [N03~] values for Greifensee are 4 mg/L (300 of carbon units) and 100 , respectively. Solar irradiation data are for sea level, after ref 1.

    strates (Figure la). Assuming an average chromophore unit weight of 120 for D O M in water containing 4 mg of dissolved organic carbon (DOC) per liter, we arrive at a chromophore concentration of 0.033 mM. Thus, each chromophore is excited at a high rate of 600 times per hour.

    Some of these interactions lead to direct photochemical transformations of D O M and aqueous micropollutants to secondary products. But in addition, aquatic humic materials act as sensitizers or precursors for the production of reactive intermediates (so-called "photoreactants") such as singlet oxygen (l02) (1, 3, 4), DOM-derived peroxy radicals ( R O C ) (5-7), hydrogen peroxide (8, 9), solvated electron (eaq") (10-12), superoxide anion (02~) (13, 14) and humic structures excited to triplet states (IS).

    In addition, U V light absorbed by nitrate and nitrite produces OH* radicals (16, 17). Light-absorbing redox-active metal species may also be important sources of photoreactants, such as metals in lower valence states (18, 19).

    Of these photoreactants only H 2 O z , because of its relative inertness, accumulates and decomposes during extended illumination periods (hours).

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    In Aquatic Humic Substances; Suffet, I., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1988.

  • 366 AQUATIC HUMIC SUBSTANCES

    A l l other species are highly reactive and short-lived; they are present only at very low concentrations and only during illumination. The role of D O M as a source and a sink of photoreactants is of interest because these photoreactants can chemically transform pollutants and, in many cases, the D O M itself.

    In principle the role of D O M as source and sink for photoreactants can be discussed without detailed knowledge of particular kinetic models (see Conclusions). However, a reaction kinetic approach is required for designing experiments that yield generalizable results. The main ideas of the model are summarized here (for details, see ref. 20).

    The steady-state concentration of relatively short-lived photoreactants ([X] s s) is given by the rate with which these reactants are produced (r*), relative to the pseudo-first-order rate constant with which they become consumed kx'):

    [XL = rx * '

    (1)

    The formation rate () is proportional to the rate of light absorption by the photochemical source substance (i.e., proportional to fcA[A]) and to the quantum efficiency (). As shown in equation 2, the rate of X consumption or quenching can be controlled by solvent quenching (kq), reaction with D O M acting as a scavenger (S) for the photoreactant (X) (fc x s [DOM]), reaction with oxygen (kXto2[02]), reaction with other scavengers, and possibly by bimolecular reactions with itself (&X,X[X]).

    DOM or

    NO3~ r = kA [ ]

    A:photon absorber X:photo^ react ant : probe molecule or

    micropollutant kA : specific light absorption

    rate-constant

    2 OH

    02" ROO*

    H 2 0 2

    solvent

    + DOM kx.S[D0M] + 02 k X , 0 2 0 2

    It

    *x,x

    ^transformed

    (2)

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 367

    Rate constants of reactions controlling the fate of transient reactants cover a wide range. Therefore, in most cases only one of these reactions dominantly controls the lifetime of a specific photoreactant in a given system.

    To quantify production rates and steady-state concentrations of the main photoreactants ( 1 0 2 , , R O O ' , and e a q"), rates of their selective reactions with added probe molecules (P) were determined (equation 2). Highly selective probe molecules were chosen to discriminate between different types of photoreactants. Whenever possible, probe compounds with structures similar to those of the micropollutants of interest were applied.

    To probe for , and R O O ' , experiments were performed in a way that produced a simple second-order rate law. The rate of transformation of was first order in concentration of both and X .

    - ^ = * P , X [ X L [P] (3)

    where k?x is the second-order rate constant for the reaction of X with P. For a closed parcel of water, and if the concentration of is low enough not to change [X] s s significantly, equation 3 integrates to

    - l n j S ^ = *p , x [X] s s X t (4) IT Jo

    Therefore, the logarithm of the relative residual concentration of declines linearly with time (t) with a slope of JfcP X[X] s s (apparent first-order kinetics). Then [X] s s can be calculated directly from the experimental elimination rate constants in cases where fcPX of a selected probe molecule is known. A l l kinetic experiments used in this overview for deducing rate constants yielded very good first-order plots.

    Transient production rate decreases with depth (z) in a surface water because of light screening by D O M (Figure 2). In waters that contain particles, light-scattering terms and absorption by particles must also be accounted for. Assuming complete mixing, measured surface rates (r ( 0 )) for production of photoreactants are normally converted to depth-averaged rates (r ( z )) by multiplying by the light-screening factor S[z ) (25).

    rw = X S f (5)

    where

    2.3 (6)

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  • 368 AQUATIC HUMIC SUBSTANCES

    Here is the deeadic absorption coefficient of the water at wavelength and is the average water depth. In principle, the choice of wavelength for the corrections must account for the overlap (product) between the action spectrum (quantum efficiency times molar absorptivity) of the reaction considered and the spectrum of the light. Application of the screening factor to waters of medium depth may be highly complicated. However, for large depths (such that x > 1), the light-screening factor approximates

    = (7) 2.3

    and the depth-averaged production rate becomes

    r(o)

    = 1 (8) 2.3

    Given the assumption of vertical mixing, equation 8 corresponds essentially to a dilution of the photochemical effect with increasing depth. For example, most U V light is absorbed within the top meter of even slightly eutrophic lakes (J). When most light is absorbed within the considered depth, z, r{z) is independent of the concentration of D O M (DOC) if both and r ( 0 ) increase proportionally to D O C . However, it decreases proportionally with the depth of the water body, because a higher absorbance at the surface is directly compensated by lower light penetration. If r ( 0 ) is independent of D O C (e.g., is proportional to N 0 3 " ) , then r ( z ) decreases with D O C and z.

    Finally, diurnal and seasonal variations in light intensity (Figure lb) must be taken into account in any generalization of results.

    Characteristics of Various Photooxidants Singlet Oxygen. It is possible to measure the steady-state concen

    tration of singlet oxygen by following the oxygenation of selective probe molecules such as dimethylfuran and furfuryl alcohol (I, 3). Furfuryl alcohol is less volatile and leads to products that are highly specific for *0 2 reactions. Dimethylfuran, however, needs lower exposure times.

    From the literature and our own studies (I) we conclude that the formation of * 0 2 from ground-state oxygen (30^ is sensitized by D O M , and that in natural waters its destruction rate is generally controlled by solvent (water) quenching. The sequence of reactions is given in equation 9.

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 369

    SINGLET OXYGEN FORMATION :

    atmos.

    : e.g. furfurylalcohol

    R}xid

    At D O M concentrations typical for surface waters (DOC < 20 mg/L), quenching of singlet oxygen by D O M can be neglected.

    Different types of aquatic D O M exhibit different quantum efficiencies for l02 production. Selected examples are presented in Table I. In summarizing phenomenological studies using a variety of different humic substances, we conclude that D O M with higher specific light absorption exhibits somewhat lower quantum efficiencies (I). No significant relationship between quantum efficiency and molecular weight fraction was found (I).

    As an example, the steady-state concentration of singlet oxygen at the surface of the somewhat eutrophic pre-alpine Lake Greifensee in Switzerland (DOC ~ 4 mg/L) during June, with summer midday sunshine (1 kW/m 2 ) is 8 10" 1 4 M . This number results from the observed furfuryl alcohol elimination rate of 3%/h.

    A light-screening factor can be estimated on the basis of action and absorption spectra typical for this lake water. Accounting for this, equations

    Table I. Half-Lives of Added Furfuryl Alcohol and Corresponding Steady-State Concentrations of l Q 2

    Water DOC tie FF A* [ ] . a ' b Source (mg/L) (ft) (M x JO

    14) (M x I014) Greifensee 3.5 20 8 4.6 Etang de la Gruyre 13 6 28 1.5 Rhine, Basel 3.2 27 5.6 3.6 Secondary effluent** 15 14 11 2.2 "Summer midday sunlight, 1 kW/m 2. ^Concentration at the surface. Concentration averaged over 1-m depth. ^Communal wastewater.

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  • 3 7 0 AQUATIC HUMIC SUBSTANCES

    1 and 8 indicate that a depth-averaged steady-state concentration for an ideally mixed top meter of the lake ( [ I 0 2 ] s s l m ) , within which most of the photoactive light is absorbed, is 5 10" 1 4 M . At greater depth, the average [ ^ I s s * simply be 5 10~14 M divided by the depth in meters. Surface values of [ 1 O 2 ] s s 0 increase proportionally with the rate of light absorption (i.e., with the D O C of the water) (Table II). For a water depth within which all light is absorbed, the area-based production of *0 2 is independent of D O C (see last column in Table II). However, a comparison of very different types of surface waters must consider that quantum efficiences and action spectra for producing l02 vary somewhat with the type of D O M (J).

    The occurrence of singlet oxygen in sunlit waters can be important for the elimination of cyclic 1,3 dienes, polynuclear aromatic hydrocarbons, organic sulfides, and phenolic compounds when the latter are significantly dissociated into the reactive phenolate anions (e.g., chlorophenols) (Figure 3). For phenols with high phenolic p K a values, correspondingly low reactivities are found. (For comprehensive literature, see ref. 21).

    O H Radical. Because O H radicals react with most organic substrates at nearly diffusion-controlled rates (H atom abstraction or *OH addition reactions), most organic substances can be applied as *OH probe molecules. Butyl chloride, in addition to other compounds, was often used as our probe substance of choice because of its easy analysis and its inertness against direct photolyis and other photoreactants (except eaq")(20).

    In fresh surface waters a large part of is produced from slow photolysis of nitrate (J 7, 20). The reaction is shown in equation 10.

    OH R A D I C A L F O R M A T I O N

    NO2 H* X~320nm > \ ^

    NO3- 0"

    P: e.g. butylchloride

    OH

    DOM

    - ) 5 S _ 1 *

    I I

    I I

    *oxid

    (10)

    * for Greifenseewater (D0O4mg/ l )

    The absorption of the relevant solar U V light by nitrate (at about 310 nm) is much smaller than that by D O M (Figure 2). As has been shown in preceding studies in which decomposed ozone has been used as an O H

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  • Tab

    le I

    I. P

    hoto

    chem

    ical

    Tra

    nsie

    nt R

    eact

    ants

    Pro

    duce

    d in

    Lak

    e G

    reife

    nsee

    and

    Fun

    ctio

    nal

    Dep

    ende

    ncie

    s of

    The

    ir S

    tead

    y-St

    ate

    Con

    cent

    ratio

    ns

    Rea

    ctan

    t P

    robe

    Mol

    ecu

    le,

    k P,

    (M-'

    s-1)

    kp

    .l[X

    i]u,

    [XU

    (%/h

    ) [

    XL

    0

    (M)

    [X]

    Jm

    a

    m

    Fu

    nct

    ion

    alit

    ies

    of [X

    ] M

    Su

    rfac

    e 1-

    m L

    ayer

    0

    2

    furf

    uryl

    alc

    ohol

    (1.2

    x 1

    08)

    3 8

    1

    01

    4 5

    x 1

    01

    4

  • 372 AQUATIC HUMIC SUBSTANCES

    R E A C T I O N R A T E S IN P R E S E N C E a ) b )

    J I I I I I L 6 8 10 pH

    Figure 3. Compilation of rate constants for reactions of 2 (left scale), and sunlight irradiation times required for solute eliminations (right scales, 1-m average depth) vs. pH. Data are from ref 21. (a) Half-life of selected pollutants in Greifensee water during exposure to June midday sunlight (1 kW/m2). For the case of l02 this yields a [ i 0 2 ] l m value of 4 14 M . (b) Scale of times for achieving the irradiation dose required for the reduction of the concentration to 50% of its initial value. These times are an estimate based on the real sum-curve of measured solar irradiations in Dubendorf, Switzerland,

    when starting on a clear summer day (i.e., June 2, 1985, 11 a.m.).

    radical source (22, 23), is primarily consumed by fast scavenging by D O M or, in waters with low D O M , even by carbonate anions. Different types of aquatic D O M exhibit comparable rate constants for trapping . Thus, we can assume that the value of [*OH] s s increases with the ratio of the concentration of N 0 3 " to D O C , as shown in Figure 4.

    For deep water columns, within which most photoactive light is absorbed, the light-screening factor increases with D O C (equation 8). Therefore depth-averaged [*OH] s s values for water columns in which all light is absorbed decrease with the square of the D O M concentration (see entry in Table II).

    On the basis of detailed measurements of O H information quantum yield, absorption spectrum of N 0 3 " , and experimentally calibrated rate constants for consumption of by D O M , we estimate the rate constant for elimination of compounds with reactivity similar to that of butyl chloride for the surface of Lake Greifensee ( [ N 0 3 - N ] / D O C = 0.4 mg/mg) to be

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 373

    [0H ] S S VS relative NO3" concentrt for June noon sunshine

    M

    ion

    10 14 |0Hl S S , max

    10 r15

    10 r16 / S'1" H10

    100

    a) Om M/2 ideal hr

    1000

    01 1 NO3-N

    DOC

    10 (mg) (mg)

    100

    Figure 4. Surface values of [0H]SS dunng summer midday sunshine (1 kW/m2) vs. the ratio of [NO3~]/DOC. Values are given (right scale) for a compound of mean reactivity (kon.e = 6 X JO9 M'1 exposed to ideal midday summer sunshine (1 kW/m2). Star indicates that the experiment was performed in Greifensee water. Bold line defines range within which most medium-sized Swiss nvers lie. Nonlinear relationships could prevail at high

    nitrate concentrations.

    0.2%/h in June midday sunshine (1 kW/m 2 ) . This corresponds to a [ 'OH] S S 0

    value of 2 x 1 6 M . Rates of reaction of O H * with many compounds in natural waters may

    be estimated from the calibrations performed with probe molecules and with information from tables of relative rate constants (Figure 5).

    Organic Peroxy Radicals. Compounds that are classified as antioxidants (such as alkylphenols, aromatic amines, thiophenols, and imines) generally exhibit moderately high reactivity toward organic peroxy radicals (ROO) (26). 2,4,6-Trimethylphenol, the most water-soluble representative of the antioxidant class of 2,4,6-trialkylphenols, was our preferred probe molecule for characterizing the organic peroxy radical reactivity of natural waters (7). Experiments have demonstrated that its direct sunlight photolysis is slow, and that possible interfering reactions with singlet oxygen are negligible.

    A plausible reaction sequence for the formation of ROO* is presented in equation 11.

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  • 374 AQUATIC HUMIC SUBSTANCES

    P E R O X Y R A D I C A L F O R M A T I O N :

    02

    (11)

    P:e.g. trimethylphenol

    ^oxid

    Experiments show that some electron-rich phenols are rapidly oxdized by a transient oxidant postulated to be an organic peroxy radical. A l l phenols exhibited apparent first-order oxidation kinetics, consistent with the assumption of the kinetic model described previously. 2,4,6-Trimethylphenol is photooxidized with an apparent first-order rate constant of 15% per hour in sunlit (midday, June, 1 kW/m 2 ) Greifensee water. Other phenols exhibit lower reactivity toward this transient oxidant, as shown in Figure 6.

    kE0H,P M-V 1

    R E A C T I O N R A T E S IN P R E S E N C E O F R A D I C A L S

    10,v

    10"

    10e

    I 9

    a) b) 1m 1m

    /2 M/2 ideal real

    hr

    500 OH n-C8H190H

    C4H9CI

    CCl2=CCl2 n-C^gOH

    VC2H3-0H" H5-10J

    h- 3 CH3COOH 10'

    +

    |4months

    3 years

    30 years

    Figure 5. Compilation of rate constants for reactions of OH radicals (left scale), and sunlight irradiation times required for solute elimination (right scales) as in Figure 3. k 0 H . p values for tetrachloroethylene and butyl chloride are from our own measurements of competition kinetics (22); other data are from ref. 23. The mean half-life at the surface would be about 20 times shorter because

    of the absence of light screening by DOM. ['OH]ssIm = I 1017 M.

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  • 23. HOIGN E T AL. Photochemically Produced Transient Reactants 375

    a) b) 1m .1m

    REACTION RATES IN PRESENCE XM2 M/2 OF "ROO*" ideal real

    k r e lt OH J

    ^ 3 OH

    40

    20

    OH

    CH 3

    OH

    H3CC(H)CH3

    OH

    OCH3

    OH

    C2H5

    OH

    CgH19

    20

    50

    100

    200

    500H

    1 day

    2 days

    10 days

    30 days

    80 days

    Figure 6. Compilation of experimentally determined refotive rate constants for ROO' radicals (left scale), and sunlight irradiation times required for solute eliminations in Greifensee (right scales; see Figure 3 for sunlight intensity). Data are from ref 7. Half-life scales are depth-averaged values for a well-mixed 1-m water column, calculated by using equation 6 with 366 = 0.01 cm1. Calcuhtions using 3 = 0.022 cm1 predict values to be

    1.8 times the half-lives given.

    4-Alkylphenols exhibit the following relative reactivity sequence: methyl > ethyl > isopropyl > nonyl > H . The relatively low reactivity of nonylphenol might be caused by steric hindrance or hydrophobic interactions of the alkyl group with parts of the D O M , which could inhibit attack of the phenol by the transient oxidant. Benzoate and hydroxybenzoates (ortho, meta, and para forms) exhibited no measurable reactivity toward ROO*.

    From diagnostic tests and kinetic analyses, we conclude that D O M is the source of this transient oxidant, and we definitely exclude singlet oxygen, hydroxyl radical, and D O M triplets as responsible for the oxidation of these phenols (7).

    Steady-state concentrations of the transient oxidant (postulated ROO*) are linearly related to D O C (for D O C < 5 mg/L) and to sunlight intensity.

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  • 376 AQUATIC HUMIC SUBSTANCES

    This relationship indicates that the fate of this transient oxidant is neither controlled by D O M scavenging (for D O C < 5 mg/L) nor by reaction with itself. The absence of ROO* scavenging by D O M is not surprising because ROO* exhibits negligible reactivity toward benzoate and hydroxybenzoates. Therefore, for low D O C we assume that D O M is a source, but not a sink, of the transient oxidant. In such a case the depth-averaged [ROO*] s s for a well-mixed water column (aK-z > 1) is independent of the D O C concentration (Table II).

    Additionally, the area-based production rate of postulated ROO* is independent of D O C concentration (for D O C < 5 mg/L), provided that most photoactive light is absorbed within the water column. However, at higher D O C (DOC > 5 mg/L) scavenging reactions of D O M may begin to control the fate of the transient oxidant (proposed ROO') . Under these conditions D O M may function as both source and sink of the postulated ROO*.

    In contrast to the other photoreactants discussed here, ROO* are D O M -derived radicals and contain parts of the D O M . Consequently the ROO* speciation, like that of D O M , is complicated, unknown, and dependent on water source. Therefore, DOM-derived ROO* may well exhibit a wide range of reactivities. This variation introduces considerable uncertainty into the determination of their steady-state concentration and production rate by any competing kinetic method (7).

    Solvated Electron. Flash photolysis combined with kinetic spectroscopy of natural-water D O M has shown high concentrations of a transient whose kinetic and spectroscopic characteristics correspond with those of the solvated electron (10-12), e^". However, experiments with added probe compounds at lower light intensities, which are more representative of continuous (sunlight) irradiations, show that the chemical effect of such species is rather low.

    We assume that the discrepancy between the high apparent production rates of transient absorbers representative of solvated electrons and the low efficiency observed for typical chemical reactions of solvated electrons might be due to the fact that most of the transients remain trapped in the mac-romolecular structure. Therefore, transient recombination may occur before reaction with compounds of interest.

    The postulated formation sequence of e^" and its reactions with probe molecules are described in equation 12.

    Probe molecules of choice are chlorinated organic compounds. These rapidly scavenge free solvated electrons to produce chloride anion. However, the formation rate of solvated electrons appears so low, and the rate with which solvated electrons in aerobic water are scavenged by dissolved oxygen is so high (~5 10 6 s"1), that their steady-state concentration is very low. Consequently, probe molecules (present in low concentrations) are not significantly reduced by solvated electrons within feasible irradiation times.

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 377

    aq F O R M A T I O N

    DOM* H20 02 (0.3mM)

    DOM -1 f 0 2 " (1)

    P: e.g. chloroethanol

    c r

    Therefore, for the determination of these photoreactants, high concentrations of probe molecules were used to trap all solvated electrons before they reacted with oxygen. The total production rate of solvated electrons was determined from the formation rate of CI" in such experiments. Then the steady-state concentration of solvated electrons was determined by combining the results with the known rate constant of solvated electrons with oxygen.

    From these results we could deduce that the value of [ e ^ ^ 0 for Lake Greifensee water during June midday sunshine (1 kW/m 2 ) is about 1.2 X 10" 1 7 M . The effect of such a concentration on reactive micropollutants is exemplified in Figure 7 and in Table II. Compounds with the kinetic characteristics of trichloroacetic acid and chloroform, which exhibit maximal rate constants, would be degraded at a rate of only 0.13%/h at the water surface.

    Although compounds present in D O M are presumably the main precursor of solvated electrons, reactions with oxygen constitute the major sink. The value of [ ^ ] ^ 0 will therefore increase with the water concentration ratio of D O C to 0 2 . However, the mean concentration, when averaged over a water depth in which all photoactive light is absorbed, again becomes independent of the D O C (see last column in Table II).

    Most solvated electrons wil l add to oxygen and form 0 2 ~. Part of the 0 2 " seems to be converted to H 2 0 2 (14). However, because added electron scavengers did not affect the formation rate of hydrogen peroxide in i l luminated Greifensee water, the solvated electron itself is not considered a main precursor for this species (27).

    Humic-derived cation radicals must also be considered photoionization products occurring when electrons are released. Such cations wil l eventually lead to oxygenated products and presumably could be an additional source of 0 2 ~ or even hydrogen peroxide. But because the yield of solvated electrons

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  • 378 AQUATIC HUMIC SUBSTANCES

    REACTION RATES IN PRESENCE a) b) OF e aq

    NfV 1

    10

    A0m Om M/2 M/2

    ideal real

    i 1 0 - \CCl2=

    10-

    10*

    10'

    N03" o 2 RNO2 - | cci2 ecu CHCI3

    CC13COO" CI CH3CI

    OH

    0 CH3

    hr + I 9 months 103 H

    10 4 H

    105

    IOH

    7 V 2 years

    75 years

    750 years

    Figure 7. Compihtion of rate constants for surface-level reactions of solvated electrons (left scale) and sunlight irradiation times required for solute elimi-nation (right scales; sunlight intensity as in Figure 3). & e a q - ,p values from ref 25. The mean half-life within a steadily mixed 1-m layer would be about 5 times longer because of light screening by DOM. [eoq']**0"1 = 12 17 M.

    is so low, such reactions can be of only minor importance in producing further photooxidants.

    Conclusions

    Table II summarizes the apparent rate constants with which appropriate probe or reference substances are transformed by the various photoreac-tants produced in eutrophic lake water under summer midday sunshine (1 kW/m 2 ) . Steady-state concentrations of the photoreactants are calculated by using absolute or estimated reaction-rate constants for reactions of photoreactants with the listed probe substances. The functions in the last two columns describe the dependence of the photoreactant steady-state concentration on D O C or other relevant water parameters.

    The dimensions of the functions describing steady-state concentrations are different for different photoreactants; they depend on the photolytic source and on the lifetime-controlling sinks to be considered. But in spite of this circumstance, which makes comparisons difficult, the results show that l02 and ROO* radicals are efficient photoreactants for transformations of some specific types of micropollutants. However, both of them react highly

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 379

    selectively and wil l degrade only a few types of chemical structures within a reasonable time.

    On the basis of steady-state concentrations given in Table II and known lifetimes of the various photoreactants in these waters, we can also estimate the total photoreactant production per kilowatt-hour of absorbed light or per integral amount of sunshine throughout 1 year (about 1300 kW-h/m 2 ) . The resulting values reported in Table III for the yearly production of photoreactants per square meter are remarkably high. Comparisons between these yearly production rates and the amount of D O M present per unit of area in the water column indicate that these reactants may also be important for the aging of D O M . For example, a lake containing 4 mg/L of D O C as " D O M molecular units" (molecular weight = 120 g of C/mol of such units) contains 0.3 mol of D O M molecular units per square meter over a 10-m depth. Comparison of this concentration with entries in Table III shows that 10% of the D O M molecular units would react, for example, with radicals in 1 year.

    Table III. Production Rates of Photoreactants (rx) in Greifensee Water

    Reactant mol/kW mol/(m2'yr) 2 50 x - 3 m 20 x 6 0.025 ROO* eq 70 x 6 0.1 NOTE: For water containing 4 mg of DOC/L, we calculate 0.03 mol of DOM units per square meter and per meter depth, assuming a molecular weight of 120 g of C/mol. This means that DOM/m 2 = 0.03 mol z(m) (molecular units: 120 g of C/mol).

    Aging effects of D O M induced by the other photoreactants cannot be quantified as easily because these do not significantly react with the D O M under the conditions considered in this overview. However, for each pho-toproduced solvated electron, a corresponding D O M cation radical is also produced. We expect the D O M cation radical to undergo molecular rearrangements that may either split out a cationic entity (possibly a proton) or add an anion (e.g., OH") and 0 2 (a biradical). Therefore, D O M might also undergo significant transformations by such reactions.

    On the basis of the data given in Table III for the production rate of solvated electrons, we deduce that about 40% of the D O M present in a 10-m-deep water layer can be converted in 1 year, if we accept the same assumptions. Although these latter ideas are somewhat speculative, they show that photoreactants could have large effects on the chemical transformations of humic substances. Data are not available to make additional comparisons of these effects with other possible direct photolytic reactions of D O M .

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  • 380 AQUATIC HUMIC SUBSTANCES

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    1977, 267, 421-423. 4. Wolff, C. J. M . ; Halmans, M . T. H . ; van der Heijde, H . B. Chemosphere 1981,

    10, 59-62. 5. Mil l , T.; Hendry, D. G.; Richardson, H . Science (Washington, DC) 1980, 207,

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    T.-W.; Bomberger, D. C.; Mil l , T. Environmental Pathways of Selected Chem-icals in Freshwater Systems: Part II. Laboratory Studies; U.S. Environmental Protection Agency: Athens, GA, 1978; EPA Report EPA-600/7-78-074.

    7. Faust, B. C.; Hoign, J. Environ. Sci. Technol. 1987, 21, 957-964. 8. Cooper, W. J.; Zika, R. G. Science (Washington, DC) 1983, 220, 711-712. 9. Draper, W. M . ; Crosby, D. G. J. Agric. Food Chem. 1981, 29, 699-702.

    10. Fischer, A. M.; Kliger, D. S.; Winterle, J. S.; Mill T. Chemosphere 1985, 14, 1299-1306.

    11. Power, J. F.; Sharma, D. K.; Langford, C. H . ; Bonneau, R.; Joussot-Dubien, J. In Photochemistry of Environmental Aquatic Systems; Zika, R. G.; Cooper, W. J., Eds.; ACS Symposium Series 327, American Chemical Society: Washington, DC, 1987; pp 157-173.

    12. Zepp, R. G.; Braun, A. M . ; Hoign, J.; Leenheer, J. A. Environ. Sci. Technol. 1987, 21, 485-490.

    13. Baxter, R. M . ; Carey, J. H . Nature (London) 1983, 306, 575-576. 14. Petasne, R. G.; Zika, R. G. Nature (London) 1987, 325, 516-518. 15. Zepp, R. G.; Schlotzhauer, P. F.; Sink, R. M . Environ. Sci. Technol. 1985, 19,

    74-81. 16. Zafiriou, O. C.; True, M . B. Mar. Chem. 1979, 8, 9-42. 17. Zepp, R. G.; Hoign, J.; Bader, H . Environ. Sci. Technol. 1987, 21, 443-450. 18. Faust, B. C.; Hoffmann, M . R. Environ. Sci. Technol. 1986, 20, 943-948. 19. Waite, T. D. In Geochemical Processes at Mineral Surfaces; Davis, J. .; Hayes,

    K. F., Eds.; ACS Symposium Series 323, American Chemical Society, Washington DC, 1986; pp 426-445.

    20. Haag, W R.; Hoign, J. Chemosphere 1985, 14, 1659-1671. 21. Scully, F. E. ; Hoign, J. Chemosphere 1987, 16, 681-694. 22. Hoign, J.; Bader, H. Ozone Sci. Eng. 1979, 1, 357-372. 23. Hoign, J.; Bader, H . In Organometals and Organometalloids, Occurrence and

    Fate in the Environment; Brinckman, F. E.; Bellama, J. M . , Eds.; ACS Symposium Series 82; American Chemical Society, Washington, DC, 1978; pp 292-313.

    24. Farhatazis; Ross, A. B. In Selected Specific Rates of Reactions of Transients from Water in Aqueous Solution. III. Hydroxyl Radical and Perhydroxyl Radical and their Radical Ions; National Bureau of Standards, Report No. NSRDS-NBS 59; Washington DC, 1977.

    25. Howard, J. .; Scaiano, J. C. In Landolt Brnstein New Series, Kinetic Rate Constants of Radical Reactions in Solution: Part d, Oxyl-, Peroxyl- and Related Radicals, Springer Verlag: Berlin, 1984; Series II/13d.

    26. Anbar, M . ; Bambenek, M . ; Ross, A. B. Selected Specific Rates of Reactions of Transients from Water in Aqueous Solution. 1. Hydrated Electron, National

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  • 23. HOIGN ET AL. Photochemically Produced Transient Reactants 381

    Bureau of Standards, U.S. Department of Commerce: Washington DC, 1973; National Standard Reference Data Series Report NSRDS-NBS 43.

    27. Sturzenegger, V. Ph. D. Thesis, Federal Institute of Technology, Zurich; in prep-aration.

    28. Zepp, R. In Humic Substances and Their Role in the Environment; Frimmel, F. H . ; Christman, R. H . , Eds.; Wiley: New York, 1988; pp 193-214.

    RECEIVED for review July 24, 1987. ACCEPTED for publication July 26, 1988.

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    23 Aquatic Humic Substances as Sources and Sinks of Photochemically Produced Transient ReactantsCharacteristics of Various PhotooxidantsConclusionsReferences

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