Photochemical decomposition of oxalic, glyoxalic and pyruvic acid catalysed by iron in atmospheric waters

Pergamon

Atmospheric Enrimnmmr Vol 2% No. 7. pp. 1231 1239. 1994 Elsewr Science Ltd Primed m Great Britain. 135?-2310194 56.CH3+0.00

1352-2310(94)EOOO4-4

PHOTOCHEMICAL DECOMPOSITION OF OXALIC, GLYOXALIC AND PYRUVIC ACID CATALYSED BY IRON IN ATMOSPHERIC WATERS Federal

and JURGHOIGNE

Zuo*

YUEGANG Swiss

Institute for Water Resources and Water Pollution Control (EAWAG). Institute of Technology ETHZ, CH-8600 Diibendorf, Switzerland

(First receioed 3 August 1993 and infinaljorm

22 December

Swiss

Federal

1993)

Abstract-Kinetics of the iron(III) catalysed photochemical degradation of oxalic, glyoxalic and pyruvic acids has been investigated under conditions representative of atmospheric waters. Degradation rates for the organic acids are proportional to the concentration of iron(lII)-substrate complexes and sunlight intensity. 1 PM dissolved iron(Il1) has a Under September noon sunlight irradiation (0.65 m Einstein m -2s-1), potential tocatalyse thedegradation ofoxalicacid at arateof lo+ I nMs-‘;ofglyoxalicacid,7f 1 nM s-‘: and of pyruvic acid, 8 f I nM s- ‘. The half life for the photolysis of oxalic acid and a-keto acids catalysed by iron(lll) is in the order of a few minutes whenever the concentration of dissolved iron(lII) is comparable to that of the organic acids. A detailed reaction mechanism has been proposed in which the complex formation between iron(II1) and the substrate is a key step. The evidence. that such an iron(III) catalysed photochemical process serves as a major removal pathway for atmospheric oxalic and a-keto acids, has also\

been discussed. Key word inder: Cloud tropospheric chemistry

chemistry,

glyoxylic

acid, iron complexes,

INTRODUCTION Dicarboxylic constituents

and z-keto acids in the troposphere.

are important A large

chemical number of

dicarboxylic and r-keto acids has been identified in ambient air, precipitation and cloudwater throughout the world (Norton et nl., 1983; Steinberg et al., 1985; Kawamura et (II.. 1985; Kawamura and Kaplan, 1987; Baltensperger and Kern, 1988; Talbot et al., 1990; Sigg er al., 1993; for review see Zuo, 1992; Sedlak and Hoignt, 1993). Among these acids, oxalic, glyoxalic and pyruvic acids are the most abundant species. Due to their polar nature, dicarboxylic and a-keto acids are preferentially transferred to the liquid phase. They constitute a significant fraction of organic aerosol associated with photochemical smog, and in part are responsible for reduction in the visibility of urban atmospheres. Oxalic and r-keto acids may play a significant role in the photochemical formation of hydroperoxy radicals and hydrogen peroxide, which are important oxidants in the transformation of inorganic and organic substances in atmospheric liquids (Zuo and Hoignt, 1990, 1992, 1993; Zepp et al., 1992; ZUO, 1992; Hoignk el (II., 1994; Sedlak and Hoignb, * Present address: Southeast Program, Florida International 33199. U.S.A.

Environmental University,

Research Miami, FL 1231

oxalic

acid, photochemistry.

pyruvic

acid,

1993). Thus, studies on the formation and removal processes of dicarboxylic acids and r-keto acids could yield further information about the fate of atmospheric pollutants and photooxidants. Ozonolysis in the presence of atmospheric water and photochemical decomposition of hydrocarbons are the main sources of tropospheric dicarboxylic and a-keto acids (Norton et al., 1983; Grosjean, 1989; Satsumabayashi et al., 1989, 1990; Talbot et al., 1990). The oxidation of aromatic compounds to glyoxal and oxalic acid was observed in studies of ozonation, ozonation in combination with UV radiation and photooxidation of aromatic compounds in aqueous solution (Leitis, 1979; Ho, 1986; Kusakabe et al., 1990; Yamamoto et al., 1979). Further oxidation of glyoxal produces glyoxalic acid and oxalic acid (Leitis, 1979). Direct emission from incomplete combustion, wind blown dust and emission by vegetation may also act as sources of dicarboxylic and r-keto acids (Andreae et al., 1988; Kawamura and Kaplan, 1987; Kawamura and Gagosian, 1990; Talbot ef al.. 1990). Except theoretical models of atmospheric deposition, little research has been published on the distribution and removal processes of dicarboxylic and aketo acids in the atmosphere. Grosjean (1983, 1989) estimated that glyoxalic, pyruvic and oxalic acids were photodegraded by solar irradiation in a few hours. In clouds, these acids will mostly partition into the liquid

1232

YIEGANG

ZIIO and J. HOIGNF

phase due to their large water solubility and deprotonation. Oxalic, glyoxalic and pyruvic acids have been detected in rain, fog- and cloudwater (Norton er al., 1983; Steinberg et al., 1985; Baltensperger and Kern, 1988; Behra and Sigg, 1990: Joos and Baltensperger, 1991; Husain and Husain, 1992; Zuo, 1992). In previous studies, it has been demonstrated that the iron-sensitized photochemical oxidation of oxalic acid produces hydroperoxide and superoxide ions (HO,./O,.-) and the iron and copper catalysed conversion of these species are important pathways for incloud production of hydrogen peroxide (Zuo and Hoigne, 1990, 1992, 1993; Sedlak and Hoigne, 1993; Faust and Allen, 1993; von Piechowski er al., 1993). Now, we present further kinetic information for the iron sensitized photochemical decomposition of oxalic, glyoxalic and pyruvic acids under conditions closely approximating those of acidified atmospheric waters. To relate the model experiments to reactions occurring in the environment, the photolysis of oxalic acid in real fogwater has also been studied.

EXPERIMENTAL Chemicals

and jog

SECTION

samples

Except where noted, all chemicals were analytical grade and were used as received. Fogwaters were collected during four fog events from September to December 1989 with a Teflon screen collector on the roof of EAWAG in the urban area of Diibendorf (Switzerland). After sampling. rogwaters were stored at 4‘C in the dark until used. The samples were filtered through 0.45 pm membrane filters (Millipore) and analysed immediately before irradiation. Analytical Oxalate

oxidation and the interference of Fe(Il) in 1he measurcmcnt radicals could not be significant in this by reducing DPD+. case. Analysis of Fe(I1) and speciation calculations were carried out as described by Zuo and Hoignc (1992 and references therein). Irradiations All irradiations were carried out in glass quartz tubes (1.5 cm i.d.) in sunlight or in a merry-go-round reactor (MGRR) equipped with a high-pressure Hg immersion lamp (Hanau TQ 718) operated at 500 W. In the MGRR. the 313 nm line or the Hg lamp was isolated with a Solidex borosilicate glass tube and a 2.0 mM K,CrO, in 0.22 M K2C0, aqueous filter solution. The minimum light path-length of the filter solution was 2.0 cm. The lamp was cooled with water and the reaction temperature was maintained at 2Ok I C by recircula1ion of the filter solution through a thermostat. Irradiarions in sunlight were performed using the same quartz test tubes held in rack at a 30’ angle from the horizontal and about I m above a black pavement. All irradiations were centered on solar noon at EAWAG (Diibendorf, Zurich; 47.4 N 440 m elevation). The ambient temperature was IO-30 C. The sunlight intensity (in mEinstein m - ‘s- ‘) was recorded every IO min using a quantum sensor (Li-COR. Li-185) that responds to light of 400-700 nm. Valcrophenone was also used as a sunlight acttnometer (Zepp et ul.. 1992). Solutions for irradtation were prepared from air-saturated stock solutions of Fe(CIO,), in 0.1 M HCIO, and potassium oxalate and adjusted to the desired pH with either HCIO, or NaOH solutions. Except where noted, all synthetic cloudwater solutions were saturated with air and contained 0.03 M NaCIO, (to adjust the ionic strength). O-10.0 /tM Fe(Ill) and various concentrations of organic substrates. De-aeration. when dcsircd. was accomplished by bubbling water-saturated high-purity N, through the solution for a1 least I5 min before irradiation.

RESULTS

techniques

was analysed

by the following

K ineticsjbr

procedures:

(a) The fi-counts of “C-labeled oxalate were determined on a liquid scintillation counter (BETAmatic I. Kontron Analytical, Zurich. Switzerland). One milliliter aliquots were taken from the sample or standard solutions, mixed with 9 ml xylene scintillation fluid (Luma Gel. LUMAC) and counted for 20 min. Calibration curves were determined from standards used in the photolytic experiments. The analytical error for oxalate concentrations above I PM is less than 5%. The t4C-labeled method was used for the experiments in which oxalic acid concentrations were present in low. (b) The DOC (dissolved organic carbon) was determined on a Dohrman DC-80 Carbon Analyser. The lower detection limit is about 4 FM oxalic acid (0.1 mg/’ carbon). (c) For experiments with real fogwater, ion-chromatography was applied with conductivity detection, A dionex

Ionpac AG 9 precolumn, a Dionex lonpac AS 9 column and a Dionex anion micromembrane suppressor were used. Aqueous bicarbonate-carbonate eluent. The detection limit deviation of less than 6%. Glyoxalate and method. The lower

pyruvate detection

solution was employed as was 0.2 /IM oxalic acid with a were measured by the DOC limit is about 4 PM glyoxalate

and 3 PM pyruvate. Hydrogen

peroxide

was measured

by the DPD

method

(Bader et al., 1988). Due to a fast dark reaction between Fe(I1) and H,O, in the presence of oxalate. Fe(lI) formed was reoxidized Reported

after samples were removed from the light source. H,O, concentrations were measured after this

photochemical

decomposition

ofosalic

acid

Fe(III) It has been shown previously (Zuo and Hoigne, 1992) that, in the absence of iron, oxalate was stable for at least I h in oxic solutions at the concentration range used when exposed to 313 nm monochromatic light or sunlight. Neither the radioactivity of the solutions of the “C-labelled oxalic acid nor the DOC of the solutions decreased during irradiation. In the presence of Fe(III), however, oxalic acid was quickly depleted under both light sources. Figure 1 shows the depletion of oxalic acid in the presence of Fe(lII) in de-aerated solutions as a function of time under 313 nm monochromatic radiation at pH 4. The stoichiometric ratio of Fe(ll) formed to oxalic acid depleted is approximately 2: I. The quantum yield (@) for the reduction of Fe(III) at 313 nm is 1.2kO.l and for decomposition of oxalate is 0.60+0.06. When the experimental results are plotted in the form log [Fe(III)]/[Fe(III)], against time, a linear relationship extending over one decade was obtained (Fig. 2). This confirms that the reaction is first-order with respect to [Fe(III)]. Figure 3a shows the photodecomposition of oxalic acid in air-saturated solution with 313 nm irradiation, in the presence

OJ

Photochemical

decomposition

of acids

1233

- 28

v0

2

4 Time

6

8

10 Time

(min)

I. Photoproduction of Fe(II) and depletion of oxalic acid lrom Fe(IlI)-oxalate with 313 nm light in de-aerated solutions. [Fe(lll)],= IO jcM, [oxalate], =30pM. pH=4.00+0.05. ionic strength=0.03 M (NaCIO,. HCIO,), de-aerated using N?. at 293 K. I 3,3nm- -060/IE/-‘s-i (valerophenone actinometer).

(min)

Fig.

0.8 0.6 0.4 0.2

o’°F7----l

o.o-

2

0

4 Time

O

2

4 Time

6 (min)

8

10

6

8

10

(min)

Fig. 3. Photodecomposition of oxalic acid (Ox) in the presence ol ferric ion and oxygen. The solution composition is the same as (or Fig. I. except for [Ox],,. (a) Under 313 nm light as in Fig. I, closed circles: [Ox], = 30 PM; open circles: [Oxlo =60 !tM. (b) Under sunlight, closed circles: [Ox],, = 60 PM; open squares: [Ox], = I20 PM. Irradiations were performed at solar noon on 16-18 September 1989 (1,=0.64~0.01 mEm-‘s-l).

Fig. 2. Photoreduction of Fe(lll) in the presence ofoxalic acid with 313 nm monochromatic light in de-aerated solutions, The data are derived from the experiments shown in Fig. I.

and Fig. 3b with sunlight. The reaction is no longer pseudo first-order in the aerated solutions containing a relatively low concentration of iron, but closer to apparent zero-order kinetics. This result reveals a catalytic nature of the reaction. 0

frflueuce

OJ the initiul

concemmion

of

2

Fe(III)

A series of experiments was carried out to test the influence of the concentration of iron on the photooxidation of oxalic acid. In these experiments, 5 PM oxalate was used and the initial Fe(II1) concentration was varied from 0 to 10 /IM. The reaction rate, r,,,, for the disappearance of oxalic acid against the initial concentration of Fe(II1) is presented in Fig. 4. The rAor is the average rate for 2 min of sunlight irradiation. It increases with the concentration of Fe(II1) over the range O-5 /tM. Thereafter, where the concentration of Fe(II1) exceeds the concentration required for the formation of an Fe(III)-mono oxalato complex, the transformation rate of oxalic acid assumes a plateau

4 tFe(lll)lo

6

8

10

MM)

Fig. 4. Rate of decomposition of oxalic acid as a timction of the initial concentration of Fe(II1). Ionic strength =0.03 M (NaCIO,, HCIO,), pH =4.00 +_0.05. [Oxlo= 5 FM. air-saturated, at 283-3000 K. Two minutes at solar noon on 22 April 1991 (lo =0.78 mE rn-*s-‘).

value. The maximum degradation rate corresponds complete complexation of oxalic acid. lnjluence

of initial

concentration

of oxalic

to

acid

The complex formation constants of oxalic acid and ferric ion are large and the ligand exchange process in

1234

YUEGANG

Zuo

this system is very fast (Baxendale and Bridge, 1955). This makes the photochemical reaction a rate-controlling step. Therefore, the overall rate of the photolytic reaction must be proportional to the concentration of the mono-, di- and tri-oxalato complexes of Fe(II1). This dependence of oxalic acid is shown in Fig. 5. The rate of disappearance of oxalic acid increases with the oxalate concentration in the low concentration range (approaching first-order kinetics, data not shown here), and becomes independent of oxalate concentration (zero-order kinetics) at high oxalate concentrations exceeding five-fold those required for the formation of the tri-oxalato complexes. Influence of pH The rate of the photodegradation of oxalic acid is strongly pH dependent. When the pH is varied from I to 6.4, the rate of degradation passes through a maximum value at about pH 3.0 (see Fig. 6).

and J. HOIGNE

Photodecomposition Fe(II1) and oxalic

at acid

more

dilute

concentration

01

In atmospheric liquids, the concentrations of iron and organic ligands are usually lower than those used in the experiments described above. To examine the kinetics for the decomposition of oxalic acid under these more representative conditions, the concentration of oxalic acid was held at IO /tM and the initial Fe(II1) concentration was varied from 0.4 to 1 PM. Typical results are shown in Fig. 7. As at higher concentrations of oxalic acid and Fe(III), the rates of degradation of oxalic acid are higher at the beginning of the irradiation and reach more constant rates within a few minutes. Higher concentrations of Fe(II1) yield faster decompositions of oxalate. Figure 7 also shows that the formation of H202 follows the same kinetic pattern as the decomposition of oxalate. Ittfluence

ofsunliyht

intensity

In oxygen-free solutions, the photolysis of ferrioxalate is proportional to light intensity (Zuo. 1992). This is also the case if air-saturated solutions are exposed to low light intensities. Representative results are given in Fig. 8, which involved experiments with two different sunlight intensities. However, when a very high intensity of irradiation was applied, this relationship was no longer observed. This is due to the fact that the acceleration of photoreduction shifts the steady-state of the photoredox equilibrium to Fe(I1) and reduces the steady-state concentration of Fe(II1). Photolysis “0

50

loo

150 IOx10

200

250

300

OIM)

Fig. 5. Effect of the initial concentration of oxalic acid on the rate of decomposition of oxalic acid. [Fe(lII)], = 10 /IM. ionic strength =0.03 M (NaClO,, HCIO,). pH =4.00~0.05. air-saturated. at 283-300 K. lllumination time: 5 min at solar noon on 28 March 1991 (I,

of oxalic

acid

in authenticJog

waters

To test the proposed mechanism, fogwater collected in Diibendorf (Switzerland) was filtrated and irradiated. The photochemical depletion of oxalic acid in these fog samples was observed (Zuo and Hoignt, 1993). A typical result is given in Fig. 9. It should be noted that

8 6 4

0

2

4

6 Time

“I

2

3

4

5

6

-7

PH Fig.

6. Erect

of pH

on the rate of photodegradation

of oxalic acid. [Fe(llI)], =0.03 M saturated,

= 10 FM, ionic strength

(NaCIO,, HCIO,), [Ox],= 120 FM, airat 283-300 K. 5 min at solar noon on 14 March 1991 (I,=O.48 mEm-*s-‘).

8

IO

12

2 t 7i C ii

14

(min)

Fig. 7. Photodepletion of oxalic acid and formation of H,O, at more dilute concentrations of oxalic acid and ferric ion under sunlight. For all solutions [Ox],= 10 /cM, ionic strength =0.03 M (NaClO,. HCIO,). air-saturated. [Fe(IIl)],=0.4 ItM pH=4.00+0.05, at 283-300 K. (squares), 0.8 PM (triangles) and 1.0 ILM (circles). trradiations were performed at solar noon in August 1991 (I, = I.1 kO.1 mEm-‘s-l).

Photochemical decomposition of acids

I235

1 5 3

0.6

0.0 Time

(min)

Sunlight

Fig. 8. Effect of sunlight intensity on photodegradation of oxalic acid. [Ox],=lO~M, [Fe(W)], = 1.0 FM, ionic strength =0.03 M (NaCIO,, HCIO,). air-saturated. pH =4.00+0.05, at 283-300 K. Irradiations were performed on 6 August 1991 (open circles: 1,=0.43mEm-‘s-‘, closed circles: 10 =0.90 mEm-‘s-l).

0.2

0.1 exposure

4

Time

0.3

(Einstein

m-2)

6

8

(min)*

Fig. 10. Photodecomposition of glyoxalic and pyruvic acids in the presence of iron: compared with that of oxalic acid. [Fe(lIl)],= 10 FM. ionic strength =0.03 M (NaCIO,. HCIO,). pH =4.0 +O.OS, [glyoxalic acid], =60 PM (open squares). [Pyruvic acid], = 60 FM (closed diamonds), [oxalic acid], = 60 PM (open circles). Irradiation was performed on 24 October 1991 (I =0.44 mEm-‘s-l for glyoxalic acid and 0.21 mEm-‘s-r for pyruvic acid and 0.64 mE rn-‘s- I for oxalic acid on I8 September). Time scale corresponds to I,, (*) =0.64mEm-*s-‘.

DISCUSSION

and kinetics Most dicarboxylic acids and cr-keto acids do not significantly absorb light at wavelength longer than 295 nm (the tropospheric solar wavelength cutoff). Nevertheless, they may undergo fast photochemical decomposition in the atmospheric aqueous phase. Such reactions occur when these organic acids, as ligands, form complexes with dissolved Fe(III), a common constituent of atmospheric waters: Mechanisms

0

30

60 Time

90

120

150

(min)

Fig. 9. Photochemical decomposition of oxalic acid and formation of HzOz in fog water (collected in Diibendorf). [Fe(III)],= 1.7 FM, [Ox],=4pM. pH 1 s- f (adapted from = 3.70 f 0.05, I 313nm= l.ZpEfZuo and Hoigne, 1993).

the disappearance of oxalic acid in irradiated fogwater is the net decomposition, that is, the sum of the production and decomposition reactions maintaining the solution concentration of oxalic acid. Photodegradation presence of iron

of glyoxalic

and pyruoic

acids

in the

The sunlight-induced decomposition of glyoxalic and pyruvic acids in the presence of Fe(II1) shows kinetic characteristics similar to those found for the depletion of oxalic acid.The disappearance of these chelating compounds is also a linear function of exposure time (Fig. 10). However, the rates of the decomposition of glyoxalic acid, 7f 1 nM s-l, and pyruvic acid, 8 f 1 nM s- ‘, at September noon in the presence of 1 PM Fe(II1) are lower than that of oxalic acid, lo+ 1 nM s- I, due to their smaller complex formation constants with Fe(II1) and lower light absorptivity of their Fe(II1) complexes.

Fe(II1) +org Fe(III)-org

-+ Fe(III)-org complex

-

hv

complex Fe(II)+

(1) org radical

(2)

Fe(I1) + org radical + O2 -+ Fe(II1) + oxidized org.

(3)

A scheme of the reactions is presented in Fig. Il. The absorption of sunlight by Fe(III)-organ0 complexes results in the electron transfer from organic ligand to the Fe(II1) and produces an Fe(I1) and an organic radical. In the presence of molecular oxygen, the organic radical could be further oxidized by O,, leading to the formation of superoxide ion (0,. -). 0,. - and its conjugate acid, the hydroperoxyl radical (HO,.), may reoxidize Fe(II), and lead to the reformation of Fe(III)-organ0 complexes and the generation of H,O,. In this process, organic compounds are degraded by the iron sensitized photolysis and by the Fenton reaction subsequently. Fe2++H202+FeOH2’+OH.

(4)

YUEGANG

1236

Zuo

and J. HOIGNC‘

0 Hz02

+ organic

compounds

-+ products.

f

(5)

q =

system acts merely as a catalyst for the photochemical oxidation of organic substances by oxygen. If conditions (light intensity, pH, concentra-

tions) are such that the rate of Fe(I1) oxidation is slow in comparison to the Fe(II1) reduction by the photochemical reaction, a relatively high steady-state concentration of Fe(I1) can be maintained in the system. In the case of the Fe(III)-oxalato complex, the oxalate radical, C,O;-. formed may reduce a molecular oxygen or an additional Fe(Ill))oxalato complex; it may also undergo a decarboxylation reaction yielding CO, and the C02. - radical, which is also a strong reductant, and can also reduce O? and Fe(II1) complexes. A detailed discussion on the relative importance of the reactions involved in this system have been given elsewhere (Zuo and Hoigne. 1992: Hoigne er al.. 1994; Zuo, 1992). According to the above mechanisms, the rate of the photochemical decomposition of Fe(III)-organ0 complexes can be expressed by:

k=2.303DrIQ1f1a

-k[Fe(III))org] (Ds[Fe(III)-org]

(6) ~0.1)

(7)

where d[Org]/dr is the initial rate of decomposition of the organic compound at the wavelength i, D is path length (cm), E>, is the decadic molar extinction coefficient of Fe(III)-organ0 complex at wavelength i. (M-‘cm-‘), @, is the quantum yield for the decomposition of the Fe(III))organo complexes at wavelength i., [Fe(III)-org] is the molar concentration of the Fe(III)-organ0 complex, and I,, is the volume-averaged incident light intensity at wavelength 1 (Einstein (s- i ). In the presence of a stoichiometric excess of oxalic acid, the Fe(III)-oxalato species, containing mono-, di- and tri-oxalato complexes, becomes the predominant form of Fe(II1) (see Fig. 12a). The rate for depletion of oxalate increases with the concentration of the Fe(II1). At higher [Fe(III)],, i.e. [Fe(III)],

lo-

=L -2 G C x

The Fe(III)-Fe(I1)

d[Org]/dt=

4

6

8

IFeW)lo

Fig. II. Scheme for the iron catalysed photochcmical decomposition of organic substrate in the cloud system.

OH.

2

equilibr. with gaseous phase

2

,,” .4

*-

10

12

14

PM

r,___._.-.-.-.___._.-.-.-.-.-.-.~

,/.’ i

6-

’

4-

/

: ”

2- ’ ,;’ 07 0

n IO

c 20

’ 30

4 50

’ 40

12. The total concentration Of Fe(III)~oxalato complexes (mono-, di- and trioxalato ions) for 298 K. pH 4, 0.03 M NaCIO,. (a) as a function of the concentration of Fe(II1); [Ox], = 5 /cM: (b) as a function of the concentration of oxalic acid; [Fe(IlI)],= 10 PM. Fig.

>[C20:-],,, the reaction rate remains nearly constant with the increase of [Fe(III)], as indicated in Fig. 4. If the initial concentration of oxalate is less than the initial concentration of the Fe(III), the [Fe(III)-ox], and the rate of reaction is proportional to the concentration of oxalate. At higher oxalate concentrations, i.e. if [C,O:-],~[Fe(I11)],, the total concentration of Fe(IlI)-organ0 oxalate complexes containing one, two or three ligands is independent of [C,O:-1, (Fig. 12b). However, the rate of reaction still increases with increasing of [CzO:-], even where [C,O:-1, is three times higher than [Fe(III)], (Fig. 5). This may be explained by the interaction the intermediates formed considering that elevated

ofexcess

oxalate

ion with

during the irradiation concentrations

of

or by oxalate

accelerate the reoxidation of Fe(I1) by O2 or H,O, (Zuo ,~ncl tloignt?. 1092: Sedlak and Hoigne, 1993). of pH on the photodeyrudation of oxalic acid The pH effect presented in Fig. 6 could be due to the fact that the pH influences the speciation of Fe(III)-oxalato complexes and oxalate ions. In deaerated solutions, experimental results have indicated that the variation of Fe(II1) oxalato species only Effect

Photochemical

decomposition

0.6

PH

Fig. 13. Fraction of uncomplexed oxalic acid and oxalate ion in Fe(III)-oxalate system ([Fe(III)] = 10 gM, [Ox] = 120 /cM, 0.03 M NaClO,/HCIO,, 298 K).

slightly affects the photoproduction of Fe(II) (Zuo, 1992) and are not sufficient to explain the strong pH influence in the oxygenated solutions, It is worth noting that the HC,O,. ion also achieves a maximum concentration near pH 3 in the uncomplexed oxalic acid system (Fig. 13). Its concentration decreases sharply at higher and lower pH values. It could imply that some photochemically produced intermediates react with HC,O, - ions much more easily than with nondeprotonated oxalic acid and oxalate ion. The change of pH, however, affects the protonation states of the intermediates formed. For example, the CO,H radical has a dissociation constant (pK,) of 1.4. Rowan et al. (1974) have estimated the pk, of the HC,O, radical to lie somewhere between the pK, of HC,O; (4.1) and that of the CO,H radical. The protonated radicals exhibit a redox ability which is very different from that of the corresponding deprotonated species. This could also be a reason for the pH effect observed in Fig. 6. Environmental

consideration

In order to determine the significance of iron catalysed photochemical decomposition as a major removal process for oxalic acid and r-keto acids in the atmosphere, it is necessary to evaluate both the reaction rate constants and the concentrations of dissolved Fe(H), oxalic acid and a-keto acids in the atmospheric waters. During the past two decades, extensive field measurements have provided strong evidences that iron is a common component in atmospheric waters (Behra and Sigg, 1990; Pehkonen ef al., 1992; Weschler er al., 1986; Graedel er a/., 1986; Zuo, 1992 and references cited therein. For a synoptic comparison, see also Fig. 1 in Sedlak and Hoigne, 1993). The concentration of dissolved iron species ranges from IO-’ to 10m4 M. The speciation distribution of iron is a complex function of sunlight intensity, and the concentrations of oxidants (HO,., 0,. -, H,O,, 0,, OH., 0,), reductants (HO,., O,.-, S (IV)) and complexing agents. Model calculations indicate that

of acids

1237

the Fe(III)-oxalato complexes and other Fe(III)organo-complexes can be the predominant species of Fe(II1) in cloud, fog and rain waters (Zuo and Hoignt, 1992; Hoigne er al., 1994; Zuo, 1992; Sedlak and Hoigne, 1993). Iron catalysed photodegradation may play a major role in the removal of atmospheric oxalic and a-keto acids. For instance I PM Fe(II1) has a potential to catalyse the decomposition of oxalic acid at a rate of lo+ 1 nM s-‘, of glyoxalic acid, 7 fl nMs-‘;andofpyruvicacid,8*1 nMs-‘atsolar noon in September (I, = 0.65 mE m ’ s- ’ ). If the concentration of dissolved iron species is of the same order of magnitude as that of oxalic and r-keto acids, the half life for the photolysis is on the order of a few minutes. During the daytime, in-cloud or in-fog depletion of oxalic, pyruvic and glyoxalic acids by this iron catalysed photochemical process is therefore much faster than that by other removal pathways, such as dry deposition, rainout or the slow reaction with OH radicals. On the other hand, atmospheric oxalic acid and rketo acids are mainly derived from photochemical oxidation of hydrocarbons in the gas phase and in tropospheric liquids. Therefore, formation and degradation of these acids during cloud events are taking place simultaneously, and under continuous irradiation the reactions will reach a steady-state which is controlled by the overall formation rate and degradation rate constants. If the concentration of Fe(II1) and oxalic and z-keto acids at steady-state as well as the depletion rate are known, the formation rate of these compounds in the atmosphere can be estimated. This would become useful to understand the mechanisms and kinetics for the transformation of hydrocarbons and the significance ofdicarboxylic and r-keto acids in the troposphere. However, uncertainties for quantifying the efficiency and degree of in-cloud events limit the accuracy of such models. Field studies involving clouds should, therefore, be used for calibrating such models. Consequently, the experimental results given here show which parameters and rate equations have to be considered when formulating appropriate submodels for interpreting such field data. It is worth noting that Fe(II1) also forms coordination complexes (or, at least, ion pairs) with a great number of other organic anions. Formic and acetic acids can form ion-pair complexes with Fe(II1) (Baxendale and Bridge, 1955; Carey and Langford. 1975). Our preliminary experimental results haye shown that the photolysis of these ion-pair complexes may have significantly been depleting the formic and acetic acids in atmospheric waters (Zuo, 1992). In this study, we have focussed on iron sensitized reactions which lead to a photodegradation of iron complexing compounds dissolved in atmospheric waters. The results of this study are rather complimentary to those found for photodegradations of oxalic acid occurring on surfaces of particulate iron oxides which could be further relevance for the aerosol chemistry (Sulzberger er al., 1994). Photosensitization

1238

YUEGANG

Zuo

by iron species is also of importance for many pollutants present in surface water;. For example, aminopolycarboxylic acids. such as EDTA (ethylenediaminetetraacetic acid). are also quite selective to complex to ferric iron, and they even form iron complexes in surface waters. Due to concern over the increased introduction of these compounds into the environment through both industrial and domestic uses, and their high stability against the transformation by microorganisms and by thermal chemical reactions, much attention has been put on their sunlight degradation

in

the

presence

of

Fe(III)

in

natural

waters

(Frank and Ru, 1990; Lockhart and Blakley, 1975; Trott et al., 1972). The absorption spectra of these Fe(II1) complexes usually show a high absorption intensity in the visible and/or near UV spectral regions. Such absorption is generally attributed to LMCT (ligand to metal charge transfer) or IPCT (ionpair charge transfer) transitions. The photochemical process of these Fe(II1) complexes always consists of the reduction of ferric iron and the generation of a corresponding organic radical, which can then be further oxidized by oxygen. leading to evolution of carbon dioxide and formation of other oxidation products of the acids. The Fe(I1) formed can be reoxidized process

by the is the

intermediates

analog

formed

of equations

SUMMARY

AND

absorb

sunlight. and

ligand

to

Fe(I1)

and

reduces

an organic

dissolved acid by

HOz..

acid

electron

I I.

transfer

radical.

Fe(III)-

complexes

The

organic

to superoxide

ion

HO?..

The

Fe(I1)

is formed

undergo

reaction,

Oz

which

This

Fig.

Fe(II1) catalysed a major removal acid and a-keto a-keto acids and with the correscomplexes easily

photoexcitation,

Fe(III)+keto metal

conjugate oxidized

Upon

and

CONCLUSIONS

This study demonstrates that photochemical degradation can be mechanism for atmospheric oxalic acids. In atmospheric waters, oxalic, ferric iron are in a fast equilibrium ponding Fe(III)-complexes. These oxalato

or by air.

(l)-(3)

radical Oz.formed

in this

a

yielding then and

its

is

re-

process

or

transferred from the gaseous phase into the atmospheric droplet. This leads to the generation of H,O, and the reformation of Fe(II1) which becomes complexed by further organic acids. In this photochemical process, the Fe(III)-Fe(I1) system acts merely as a catalyst for the sunlight induced oxidation of organic substance by oxygen. If the concentrations of Fe(II1) are below those of the organic ligand, the rate of photochemical degradation of oxalic and a-keto acids is directly proportional to the concentration of the Fe(II1) complexes. The experimental results show that the degradation rate of the substrates increases with sunlight intensity, with the concentration of Fe(II1) (when [Fe(III)] <[substrate]), and with the concentration of substrate. The Fe(III)-Fe(II)-Fe(II1) cycling time is in the order of only 100 s. At the solar

and

J. HOIGNC

noon in September, I PM ferric ion can photosensitize the depletion of oxalic, glyoxalic and pyruvic acids at

a

rateoflOfInMs-‘,7klnMs-‘and8flnMs-’, respectively. minutes

These for

the

concentrations Interpretations characterizing depletion rates day time, must

rates

depletion

will

give

of oxalic

a half and

life

a-keto

of a few acids

at

observed in atmospheric waters. of field data, on clouds and fogs the ratio of Fe(I1) to Fe(II1) and the of oxalic acid and a-keto acids during take account of such reactions.

Acknowledgements--We thank David Sedlak for discussions and reviewing the manuscript. Our thanks also go to Heinz Bader for technical instruction and assistance. This work is part of our contribution project entitled: “Kinetics

to the HALIPPiEUROTRAC of Oxidation Processes

subin Cloud

and Fog Waters: Influence of Transition Metal Ions”. and has been supported by the Presidential Foundation of the Swiss

Institute

of Technology

(ETH).

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