Diss. ETH No. 9727
Photochemistry of Iron(Ill)/lron(II) Complexes in Atmospheric Liquid Phases and lts Environmental Significance - Formation of Hydrogen Peroxide and Oxidation of Oxalic Acid and Other Atmospheric Pollutants
A dissertation submitted to the SWISS FEDERAL INSTITUTE OF TECHNOLOGY ZURICH for the degree of Doctor of Natural Sciences
presented by Yuegang Zuo M.S. Environmental Chemistry Born September 17, 1958 citizen of the People 's Republic of China
accepted on the recommendation of Prof. Dr. Jürg Hoigne, examiner Prof. Dr. Werner Stumm, co-examiner
Zürich 1992 To my dear parents Abstract
Tue generation of hydrogen peroxide (H202) and depletion of oxalic and other carboxylic acids by photochemical/chemical cycling of Fe(lm/Fe(II) complexes in sunlight has been studied under conditions typical of acidified atmospheric water. H202 is produced through the reduction of oxygen by intermediates formed from photo-reactions of Fe(III)-oxalato complexes. Tue rate of H202 formation increases with sunlight intensity, and with both oxalate and Fe(III) concentration within the concentration range used. This rate is 3. 7 nM s-1 when a solution of 1 µM Fe(III) and 5 µM oxalate at pH 4 is exposed to September-noon-sunlight, and oxalic acid is photolyzed with a half-life of a few minutes. Tue dissolved iron is present as a photochemical catalyst. Fe(III) is reduced by photons to Fe(II). In the presence of Oz, the Fe(II) formed is reoxidized, leading to the reformation of Fe(III) complexes. At solar noon in September, the cycling time of Fe(III)-Fe(II)-Fe(III) is on the order of minutes. Speciation calculations based on the concentration range of Fe(III) and oxalic acid present in atmospheric water indicate that Fe(III)-oxalato complexes are often the predominant species of dissolved Fe(III). Tue concentrations of Fe(III)-oxalato complexes are sufficiently large to make their photolysis a dominant source of in-cloud H02·/02·- radicals and HzOz, and a major sink for atmospheric oxalic acid. Irradiation of authentic fog waters has shown a good correlation between the formation of HzOz and depletion of oxalic acid. Mechanisms and kinetics for the photochemical/chemical cycling of Fe(III)/Fe(II) complexes, and for the formation of H202 and the depletion of oxalic acid are discussed in detail. iii
Zusammenfassung
In Wasser gelöste Eisen(III) Verbindungen katalysieren den photolytischen Abbau von Oxalsäure und anderer komplexbildender Carboxylsäuren. Es bildet sich dabei Wasserstoffperoxid. Dieser Pr'ozess wurde in Lösungen, deren chemische Zusammensetzung mit denjenigen typischer atmosphärischer Wässer vergleichbar ist, untersucht. Das H202 entsteht dabei nach einer primären photochemischen oxidativen Bildung organischer Radikale, die ihrerseits gelösten Sauerstoff zu Superoxidionen (02·-), den Vorläufern von H202, reduzieren. Die Geschwindigkeit der Bildung von H202 steigt im interessierenden Konzentrationsgebiet (Mikromol pro Liter) mit der Konzentration des gelösten Fe(lll) sowie mit derjenigen des Oxalats an. Sie beträgt beispielsweise in einer Lösung, die bei pH 4 eine Konzentration von 1 µM Fe(lll) und 5 µM Oxalat enthält und dem Mittagslicht im September exponiert wird, 3,7 ·l0-9 M·s-1; die Halbwertszeit für die Umsetzung der Oxalsäure (zu C02, gemessen) liegt entsprechend bei nur einigen Minuten. Gleichzeitig wird Fe(III) zu Fe(II) reduziert, das in sauerstoffhaltigem Wasser durch Hydroperoxyradikale (H02·) bzw. durch H202 wieder zu Fe(III) aufoxidiert wird und wieder mit noch vorhandenem Oxalat komplexiert. Im Licht hohen Sonnenstandes wird ein solcher Redoxzyklus, solange genügend Oxalat vorhanden ist, innerhalb von Minuten durchlaufen. Die Geschwindigkeit dieses photolytischen Prozesses ist hoch, denn das UV-Spektrum aller drei Fe(III)-oxalato Komplexe überlagert das solare Bestrahlungsspektrum im Wellenlängengebiet von 300 bis ca. 380 nm recht intensiv, nämlich mit molaren Absorptionkoeffizienten von> 200 M-lcm-1. Die Quantensausbeute des photolytischen Redoxprozesses beträgt bei 313 nm 1,2 ± 0,1. iv
Für die in atmosphärischen Wässern zu erwartenden Konzentrationsverhältnisse von Oxalsäure zeigt die Berechnung der Speziierung der Fe(III) Verbindungen, dass in atmosphärischem Wasser gelöstes Fe(III) namentlich in Form eines der drei Fe(IIl)-oxalato Komplexe vorliegt. Die Konzentrationen in Nebeln und Wolkenwasser anthropogen belasteter Gebiete sind zudem genügend hoch, dass diese Prozesse einerseits eine dominierende Quelle für H202 und andererseits eine dominierende Senke für atmosphärische Oxalsäure bilden können. Diese photochemische Bildung von H202 kann die Kinetik der Bildung von saurem Regen, via Oxidation von aquatischem S02 zu Schwefelsäure, wesentlich beschleunigen. Auch eine Belichtung von gesammelten hiesigen Nebelwasserproben zeigte eine gute Beziehung zwischen der Bildung von H202 und der Abnahme von Oxalsäure. V
Acknowledgements
Completion of this thesis would not bave been possible witbout the support and assistance of a great many individuals. First, 1 would like to tbank Prof. Werner Stumm for promoting these study and for acting as co- examiner. lt is also Prof. Stumm wbo made it possible for me to come EA WAG and work on this thesis Special thanks go to my tbesis advisor, Prof. Jürg Hoigne, for bis invaluable advice, support and encouragement througbout this work. 1 particularly would like to thank him for bis ability to allow me a great deal of freedom to pursue this researcb. 1 would like to thank Prof. Geoffrey Hamer for stimulating discussion during my qualifying examinations. 1 thank Dr. Barbara Sulzberger, Prof. Laura Sigg and Prof. Aliatair Kerr for their useful discussions. 1 am especially indebted to Heinz Bader for bis expert tecbnical assistance and invaluable advice in practical problems. During my work 1 could take great benefit from our group of doctoral students, post-docs and friends. 1 would like to thank all of tbem for their contribution to a pleasant atmospbere, belpful discussion and comments on my ongoing work. Special thanks to the old grade students: Bamey, Cbristophe, Dieter and Agatha, Daniel, Albert and Barbara, Beat, Paul, Slavi, Francis, Micbeal, and Jürg; the youngsters: Annette, Tina, Madeleine, Yael, Pierre, Elke, Adrian, Gianluca and Sanja; The coming students: Norbert, Timo and Xiaoxiao. The post-docs and bigber positions: Paul, Litza, Janet, David, Pbilippe van Capellen, Tony, Urs and Birgit, Annette, Silvio, Steve McDow, Bruce, Pbilippe Bebra, Hanbin, Pat, Maria, Carrick, Natascba, Jennifer and Tom, Gorgos, Erich, Kerry, Jürg Zobrist. 1 vi especially thank David, Tina, Madeleine and Tony for their invaluable discussion, comments and reading the manuscript of this thesis. I could learn a great deal from the visiting professors: Jim Morgan, Charlie O'Melia, Jerry Schnoor, George Luther Appreciation also goes to all my co-workers who helped me tackle the many analytical and everyday problems, Especially Sonja Rex, Gerda Thieme, Irma Kipfer, Beatrice Schwertfeger, Renata Roy, Annemaria Vit, Diana Hornung, Verena Cajochen, David Kistler, Hans-Ueli Laubseher, Thomas Rüttimann, Claude Jaques, Werner Roth, Ursula Mohlberg and Maria Huber-Steiner. Also a Special thank goes to our efficient and helpful librarian Elisabeth Stüssi. Finally, thanks go to my parents, wife and son (Xiaoxiao) for their love, support and encouragement throughout all of my endeavors.
This thesis work was partly supported by the Presidential Foundation of the Swiss Federal Institute of Technology under the project "Kinetics of Oxidation Processes in Cloud-, Fog- and Rainwater: Role of Transition Metals (1987)". vii Table of Contents
Abstract .„„ ... „„ ... „.„„ ... „„„„.„ ... „„„.„„.„ ..... „ ...... „ ... „ ...... „ ...... „.i Zusammenfassung „ .... „. „ „ .... „. „ ... „. „ ..... „. „ „ „ .. „. „ „ .. „ ... „ .„. „. „ ... „iii Acknowledgements ... „ .. „ ...... „ ...... „„ ..... „ .. „ .. „ .. „ ..... „ ...... „ .. „.v Table of Contents .... „ ...... „ ...... „„.„ ...... „ ... „ .. „ ...... „„ ... „ ...... „ .. vii List of Figures .... „ ...... „ ...... „ ..... „ ...... „ .. „ ...... „ „ .... xi List of Tables „ .... „ ...... „ ..... „ .... „ ..... „ .... „ ..... „ ...... „ .... „xv
1. Synopsis .. „ ..... „ ...... „ ...... „ ...... „ ...... „ ..... „ .... „ ...... „ .... „ ... 1 1.1 Introduction and motivation .„ .. „„„ .. „.„„„„ .. „„„„.„ .... „1 1.2 Objectives of the present work ...... „„ „ ...... „ „ „ „ ..... „ ... „3 1. 3 Main results .... „ ... „ ... „ ... „.„.„ .... „„. „„ „ .. „ .. „ „ „ .. „„ ... „ .. „4
2. Introduction and background ... „„.„.„.„„ .. „„ .. „ ... „ .. „„„ .... „ .. „.10 2.1 Significance and sources of H202 in atmospheric liquids .... „ ... „.„„ ...... „ ...... „.„„ ...... „ ...... „ .... „ ... „ .... „ ... 10 2.2 Proposed mechanism for the photochemical formation of H202 and decomposition of oxalic acid in atmospheric droplets ... „ .. „.„„.„„ ...... „„13 2.3 Occurrence and speciation of Fe(III)-oxalato complexes in atmospheric waters „. „„ .... „ .. „ „ .. „„ „„ „„18 2.3.1 Occurrence of iron in atmospheric waters „„.„„.„.„„.„21 2.3.2 Occurrence of oxalic acid in atmospheric waters .„.... „ .. 24
3. Experimental methods ...... „ ...... „„. „„ .. „ .. „„„„ ..... „ .. „ .. „ .. „.29 3.1 Materials ...... „ .. „ ...... „ ...... „ .. „ ...... 29 3.2 Irradiation procedures ..... „ ...... „ ... „.„ ... „.„„„„„ .. „ ... „29 3 .2 .1 Solar irradiation .„.„.„„ ..... „. „. „. „ ... „. „ „ ... „ „. „. „. „ „30 viii
3.2.2 Photochemical reactor and Monochromatic irradiation .30 3.2.3 Actinometry and quantum Yields ...... 33 3.3 Analytical techniques ...... 34 3.3.1 H202 .„.„„„.„„.„„„.„„•• „„ .. „„.„„ .. „„.„.„ •.•• „.„.„.„.34
3.3.2 Fe(//) and Fe(///) „„„„„„„„„„„„„„„„„.„„„„„„„„„.36
3.3.3 Oxalic acid „„„„„„„„„„„„.„„„„„„„„„.„„„„„„„„„.37
3 .3 .4 Valerophenone actinometer „„„„„.. „„„„. „ „„„ „„„„„.38
3.3.5 pH .„„„.„„„„„„„.„„„„.„„.. „ .. „„„„.„„„ .. „„„.„„„„.38
3 .3 .6 Electronic absorption spectra „„ „„ .„„„„„.. „„„„„ „ „ „38
3.3.7 Equilibrium calculation „„„„„„„„„„„„„.„„„„„„„„„38
3.3.8 Experimental accuracy „„„„„„„„„„„„„„„„„„„„„„„38
4. Experimental results and discussion „„„„„„„ „ „ „ .„. „ „„„„„ „ „ .41 4.1 Formation of hydrogen peroxide by photolysis of iron(III)-oxalato complexes and other iron- carboxylate systems in aqueous solution „ „. „ „„ „ „ „ „ . .41
4.1.1 Stability of oxalate „„„„„„„„„„„„„„„„„„„„„„„„.„.41 4.1.2 Oxygen requirement for H202 formation .„„ .. „„ .. „„„.41
4.1.3 Photo-formation of H202 in sunlight „„„„„„„„„„„„„.42
4.1.4 Concentration effects of Fe(Ill) „„„„„„„„„„„„„„„„ . .43
4.1.5 Effects of oxalate concentration „„„„„„„„„„„„„„„„ . .44
4.1.6 pH effects „„„„„„„„.„„„„„„„„„„„„„„„„„„„„„„„„47 4 .1.7 Stoichiometry of H202 formed to oxalate decomposed „48
4.1.8 Stability of H202 in solution „„„„„„„„„„„„„„„„„.„ . .49 4.1.9 Photochemical production of H202 from F e(//)-oxalic
acid-02 system „.„ .. „.„ ...•..•..•..•...•. „ ....•.....•.•..•.•...•.... 52 4.1.10 Formation of H202 by photolysis of other carboxylic
acid-Fe(///)-02 systems „„„„„„„„„„„„.„ .. „„„„„.„„53 ix
4.1.11 Subsequent OH radical formation ...... 55
4. 2 Photochemical/chemical cycling of Fe(III) and Fe(II) complexes in atmospheric waters ..... „ ..•....••.... „56 4.2.1 Kinetics for monochromatic photoreduction of Fe(lll) in deoxygenated solutions ...... „ ... „„ .. „ ...... „ .. „.56 4.2.2 Effects of pH and the concentratian of oxalic acid on monochromatic photoreduction ...... „ ...... • „ ...... 58 4.2.3 lnfluence of dissolved dioxygen on monochromatic photoreduction ...... 61 4 .2 .4 Kinetics for the sunligh~ photqreduction in air-free solutions ...... „ ...... „ ...... „63 4.2.5 Effects ofpH and the concentration of oxalate on the sunlight-induced reduction of ferric ion ..... „ .....•...... 65 4.2.6 Sunlight photoreduction of Fe(lll) in the presence of oxygen ...... 67 4.2.7 Solar photoreduction of Fe(lll) associated with formic and acetic acid ...... 68 4.2.8 Photochemical!chemical cycling offerric andferrous complexes ...... •...... 10 4.3 Photolysis of oxalic, glyoxalic .and pyruvic acids catalyzed by iron in cloud· and fog-waters ...... 72 4 .3 .1 Kinetics of disappearance of oxalic acid under monochromatic and solar irradiation ...... 72 4.3.2 lnfluence of the initial concentration of Fe(Ill) ...... 13 4.3 .3 lnfluence of initial concentration of oxalic acid ...... 15 4.3.4 lnfluence of pH ...... 77 4 .3 .5 P hotodecomposition at more dilute concentrations X
of Fe(Ill) and oxalic acid ...... „„.„ •.•.•.• „.„ ...•.•.. „.„ .....79 4.3.6 lnfluence of sunlight intensity „„„„„„„„„.„„„„„„„„„80 4.3.7 Photodegradation of glyoxalic and pyruvic acids in the presence of iron „„„„„„„„„„„„„„.„„„„„„„„„81
4.4 Photochemical formation of H202, depletion of oxalic acid and oxidation of dissolved S02 „„„„„„„„.82 4.4.1 Fog water concentrations and absorption spectrum „„„.82 4.4.2 Irradiation of fog water with 313 nm light „„„„„„„„„83
S. General discussion and environmental significance „ „ „ „ „ „ „. 87 5.1 Mechanism for hydrogen peroxide formation „„„„„.„87 S. 2 Environmental significance „„„„„„„ „ „ „ „ „ „ „ „ „ „„ „„ „89 5.2.1 Photolysis of Fe(Ill)-organic complexes can be a major source of H202 in atmospheric waters „„„„„„89 5.2.2 Photolysis of Fe(Ill)-organic complexes can generate a significant amount of OH radicals in atmospheric
liquids „ .. „„.„„„„„„.„„.„.„.„„.• „„„.„.„ .... „„.„.„.„„90 5.2.3 lron catalyzed photodegradation of oxalic acid is a major sink of atmospheric oxalic acid „„„„„„„„„„„.91 5.2.4 Photochemicallchemical cycling of Fe(Ill)!Fe(II) complexes provides a significant source of dissolved iron in natural surface waters „„„„„„„„„„„„„„„„„„.91
6. Conclusions „„„„.„„.„„„„„.„„.„„.„„.„.„.„„.„„ .. „„„ ... „.„.„.„.„.93
List of references „„ •• „.„ .. „„„ .. „.„„„ ... „„„ .. „„ ... „ .. „ •. „ ...... „„„.95
Curriculum vitae „ „„ „„ „ „„ „ „„ „. „ „„ „„ „ .. „ „„ „ „ „„ „ „ „ ... „ „ „ „ „„ .107 xi
List of Figures
Figure 1.1 Mole-fraction distribution of the iron(ill)-oxalato complexes ...... 4 Figure 1.2 a) First-order treatment of phot-0reduction data of Fe(ill)-oxalato complexes in deoxygenated solutions ...... 5 b) Photo-production of Fe(m in air-saturated solutions ... 5 Figure 1.3 Photochemical formation of H202 and depletion of oxalic acid with sunlight ...... 6 Figure 1.4 a) Rate of decomposition of oxalic acid with sunlight as a function of the initial concentration of oxalic acid...... 8 b) Rate of decomposition of oxalic acid in sunlight as a function of the initial concentration of Fe(IIJ) ...... 8 Figure 1.5 a) Photochemical formation of H202 and decomposition of oxalic acid in authentic fog water ...... :...... 9 b) Photooxidation of dissolved S02 in fog water ...... 9 Figure 2.1 Scheme of possible photochemical reactions in the Fe(IlD-oxalato-oxygen system ...... 19 Figure 2.2 Absorption spectra of Fe(ill)-oxalato complexes, Fe(OH)2+ and Fe3+ in the region 290-500 ...... 20 Figure 2.3 Mole-fraction distribution of the iron(III)-oxalato complexes ...... „ ...... 27 Figure 2.4 Mole-fraction distribution of the iron(II)-oxalato complexes ...... ~ ...... ;...... 28 Figure 3.1 A schematic diagram of the MGRR ...... 31 Figure 3.2 UV and visible light transmissioii spectra of the solidex sleeve and a 2.0 mM K1Cr04 + 3% (w/w) xii
K2C03 aqueous filter solution ...... „ ...... 32 Figure 3.3 Calibration curves for the determination of H 202 in the presence and in the absence of Fe(III)-oxalato complexes ...... 36 Figure 4.1.1 Hydrogen peroxide formation as a function of sunlight illumination time ...... 42 Figure 4.1.2 Hydrogen peroxide formation at a lower concentration of Fe(IIl)-oxalato complexes ...... 43 Figure 4.1.3 Hydrogen peroxide formation in sunlight as a function of the initial Fe(III) concentration ...... 44 Figure 4.1.4 Photoformation of H202 in sunlight with various oxalate concentrations ...... „„. „. „„ „„. „ ... .45 Figure 4.1.5 Hydrogen peroxide formation as a function of oxalato concentration ...... 46 Figure 4.1.6 Effect of pH on the hydrogen peroxide formation ...... „47 Figure 4.1.7 Decomposition of oxalic acid ...... „ ...... 49 Figure 4.1.8 Inhibition of hydrogen peroxide decomposition and stabilization of hydrogen peroxide by oxalate ...... „„ .. „.52 Figure 4.1.9 Photoproduction of H202 from Fe(ll)-oxalic acid- 02 system ..... „ ...... • „ ...... „.„ ...... „„ ...... „.„ .. 53 Figure 4.1.10 Formation of H202 by photolysis of other carboxylic acid-Fe(III)-02 system ...... „„ ...... „ ..... 54 Figure 4.2.1 Photoreduction of Fe(III) with 313 nm monochromatic light in deaerated solutions ...... „ „ .. „. „ .56 Figure 4.2.2 Photoproduction of Fe(II) and depletion of oxalic acid with 313 nm light in deaerated solutions „ ...... 58 Figure 4.2.3 Photoreduction of Fe(III)-oxalate complexes at various pH values with 313 nm irradiation ...... „.„„.„ .. 59 xiii
Figure 4.2.4 Photoreduction of Fe(III) with various oxalatc concentrations under 313 nm light „. „„„ „„„ „„„„„„„„60 Figure 4.2.5 Influence of dissolved oxygen on the photoreduction of Fe(IIl)-oxalato complexes with 313 nm light „„„„„„62 Figure 4.2.6 Formation of H202 by photolysis of Fe(IIl)-oxalato complexes in air-saturated solution with 313 nm light .„.„ .. „.„„.„.„„„ ... „.„„„.„.„ .. „.„„.„ ..... „.„„.„„.63 Figure 4.2.7 Sunlight photoreduction of Fe(IIl)-oxalato complexes in deoxygenated solutions „„ „„„„ „„„„„„„.64 Figure 4.2.8 Sunlight induced photoreduction of Fe(IIl)-oxalato complexes at various pH values ..• „„.„„„„„„„„„„„„„.65 Figure 4.2.9 Sunlight induced photoreduction of Fe(III) with various concentrations of oxalic acid „ „„„„„ „„„„„ „„.66 Figure 4.2.10 Sunlight induced photoreduction of Fe(IIl)-oxalato complexes in the presence of oxygen „„„„„ „„„„„ „„„.68 Figure 4.2.11 Sunlight induced photoproduction of Fe(Il) from Fe(IIl)-formic acid"02 and Fe(III)-acetic acid- 02 system „„„„„„„„„„„„„.„„„„„„„ .. „ .... „„„.„.„„ .. 69
Figure 4.3.1 Photodecomposition of oxalic acid in the presence of ferric ion and oxygen under 313 nm light „„„„„„„„72 Figure 4.3.2 Photodecomposition of oxalic acid in the prescnce of ferric ion and oxygen under sunlight „„„„„„„„„„„.73 Figure 4.3.3 Rate of decomposition of oxalic acid in sunlight as a function of the initial concentration of Fe(III) „„„.74 Figure 4.3.4 Effect of the initial concentration of ferric ion on the rate of photodegradation of oxalic acid with 313 nm light „„„„„„„„„„„„„„„„„„„„„„„„„„„75 Figure 4.3.5 Rate of decomposition of oxalic acid with sunlight xiv 1 !
as a function of the initial concentration of oxalic acid ... „. „. „ „„„ „„ „ „„ „. „ „„ „. „ „„ .... „ ... „ „ ... „. 7 6 Figure 4.3.6 Influence of pH on the degradation of oxalic acid under sunlight ..... „ ...... „ ...... „ ...... „ ...... 77 Figure 4.3.7 Fraction of uncomplexed oxalic acid and oxalate ion in Fe(III)-oxalate system „ .. „.„„„„ ... „ ... „.„.„ .. „ .... 78 Figure 4.3.8 Photodepletion of oxalic acid and formation of H202 at more dilute concentrations of oxalic acid and ferric ion under sunlight „„.„„„.„ .. „ ... „ .. „„ .. „„ ... 79 Figure 4.3.9 Effect of sunlight intensity on photodegradation of oxalic acid „„„„.„„„„„„ „„„„„„„„ .„„ .. „ „„„„ „ .. „80 Figure 4.3.10 Sunlight-induced photodegradation of glyoxalic and pyruvic acids in the presence of iron „ ..... „„ ...... „81 Figure 4.4.1 Absorbance spectrum of fog water „.„„ .. „ ..... „„„„ .. „„83 Figure 4.4.2 Photochemical formation of H202 and decomposition of oxalic acid in fog water ... „ ... „.„ ... „ .. 84 Figure 4.4.3 Photooxidation of dissolved S02 in fog water with 313 nm light „„„„.„„ .... „„„ .• „.„„„„ .... „„„ .... „ .. 85 Figure 4.4.4 Total concentration of peroxide formed and degradation of DOC and oxalic acid „„.„„ .. „„ .. „„„ ..... 86 Figure 5.1 Scheme for the photochemical/chemical cycling of iron and the formation of H202 „„„.„„„„„ ..... „.„„ .. „ .. 88 XV
List of Tables
Table 2.1 Reactions and reaction rate constants for H02·/02·- and H102 with 03, Fe and Cu ions ...... 16
Table 2.2 Total iron concentrations in rain, fog and snow waters ...... 22 Table 2.3 Percentage of dissolved iron/total iron and dissolved Fe(Il)/dissolved total iron in Dübendorf fog, snow and rain waters ...... „ ...... 23 Table 2.4 Concentrations of oxalic acid in cloud, rain, fog and snow waters .. „ ...... 25 Table 3.1 Ultraviolet-visible emission spectrum of Hanau TQ 718 high pressure Hg lamp (290-600 nm) ...... „ ... 32 Table 3.2 Composition Matrix of the iron(III)-oxalate system in perchlorate solution and the equilibrium constants . .39 Table 3.3 Composition Matrix of the iron(II)-oxalate system in perchlorate solution and the equilibrium constants . .40 Table 4.1.1 Stability of photochemically formed H102 ...... 50 Table 4.2.1 Photolysis rate constants of Fe(III)-oxalato complexes at various values of pH under 313 nm monochromatic light ...... „ ...... „ ...... 61 Table 4.2.2 Photoreduction rate constants of Fe(III)-oxalato complexes at various values of pH under sunlight ...... 67 Table 4.4.1 The concentrations of relevant constituents in fogwater and liquid water content (LWC) ...... 82 - 1 -
1. Synopsis
1.1 Introduction and motivation Hydrogen peroxide (H202) is believed to be a key component in the a main conversion of dissolved sulfur dioxide (S02) to H2S04, which is much contributor to acid rain (Calvert et al., 1985). As a result, there is for the interest in the understanding of the mechanisms responsible of formation and degradation of H20z in the atmospheric liquid phase. Most and these works have been focused on exchange processes between gas phase water atmospheric liquid phase. Tue occurrence of H202 in atmospheric and droplets is currently explained through dissolution of gaseous H2Ü2 the gas disproportionation of H02 radicals, which are also scavenged from phase (Schwartz, 1984; Lelieveld and Crutzen, 1990; McElroy, 1986). Surprisingly little attention has been paid to the in-cloud photochemical formation of H202. Photochemical electron-transfer been processes involving 02, organic matter, and transition metals have Zika, postulated to generate superoxide ions (02·-) in sea water (Petasne and 1987). 02·- and its conjugate acid, the hydroperoxyl radical (H02·), undergo rapid disproportionation to yield H2Ü2 and 02 in aqueous solutions 1991). such as atmospheric waters (Weinstein-Lloyd and Schwartz, that Extensive field measurements within the past few years have indicated and transition metals such as Fe, Mn, Cu and organic ligands such as oxalic see pyruvic acid are common constituents of atmospheric waters (for review Chapter 2). Photolysis of the aqueous complex Fe(OH)2+ has been widely studied. in This chemical process has been proposed as a major source of OH radicals et rain in the absence of organic ligands (Faust and Hoigne, 1990; Graedel oxalic al. 1986). However, the presence of organic chelate ligands, such as - 2 - and pyruvic acids in acidic abnospheric water, can significantly modify the speciation and photochemical reactivity of Fe(III) ions because they may form thermal stable Fe(IIl)-organo complexes with strong absorptions overlapping the solar spectrum. The absorption of sun-light by these Fe(IIl)-organo complexes (Fe3+.org.) results in the electron transfer from the organic ligand (org.) to the central ferric ion and produces a ferrous ion and an organic radical (org. radical). In the presence of dissolved oxygen, the organic radical could then reduce 02 to form superoxide ion Oi--, leading to H202 formation. hv Fe(IIl)-org. ~ Fe(II) + org. radical (1.1)
org. radical + 02 ~ 02·· + oxidized org. (1.2)
(1.3)
Further, photolysis of Fe(IIl)-org. complexes produces ferrous ions. This can have important implications for the generation of OH radicals through Fenton's reaction (Walling, 1975):
Fe(II) + H202 ~ Fe(III) + OH· + OH- (1.4)
The photolysis of Fe(IIl)-complexes can provide efficient pathways for the transformation of organic pollutants. First, the sunlight induced charge transfer from the complexing ligand (potential pollutant) to the center ferric ion leads to oxidation of the complexing ligand; second, this photolysis produces photooxidants such as H202, 02··, H02·, OH· and various organic and inorganic radicals, which subsequently oxidize a variety of organic and inorganic pollutants. - 3 -
This photochemical process for the formation of H202 and depletion before. of organic substances in atmospheric waters has not been studied acid- These facts let us choose oxalate-Fe(III)-02 and other carboxylic Fe(III)-02 systems as models in order to better understand the involvement of this photochemical process in the transformation of atmospheric of the pollutants and to gain deeper insight into the kinetics and mechanisms photolysis of these systems.
1.2 Objectives of the present study I) The aim of the first part of this thesis has been to examine whether under Fe(III)-oxalato complexes could be the predominant species of Fe(III) the conditions representative of atmospheric water (pH 2.5-5). II) We have been bound (a) to establish the kinetics of p9otoreduction Fe(IIl)- of Fe(III)-oxalate complexes, and the photoreactivity of individual of oxalato species in dilute solutions, and (b) to elucidate the influence of dissolved oxygen and photochemical/chemical redox cycling Fe(IIl)/Fe(II) complexes. the Im The main experimental objectives have been (a) to investigate in production of H202 by photochemical cycling of Fe(IIl)/Fe(II) complexes commonly the presence of oxalate or other organic ligands which are also rate of observed in atmospheric waters; (b) to determine the formation to compare H202 and factors which may influence the formation rate; (c) potential the rate of production of H202 from this process with other in sources and to estimate its significance to the occurrence of photooxidants atmospheric waters. the IV) Tue other main objectives have been (a) to characterize of iron kinetics for the photodecomposition of oxalic acid in the presence and the and to observe the effects of the amount of initial ferric ion - 4 - concentration of substrate; (b) to assess the life-time of oxalic acid and other keto-acids in the lower troposphere. V) In order to relate the model experiments to atmospheric aqueous environments, a real fog water was irradiated. The kinetics for the formation of H202, decomposition of oxalic acid and oxidation of dissolved S02 were studied.
1.3 Main results I) Based on the published data on the concentration ranges of iron and oxalic acid in atmospheric waters and available stability constants, the speciation calculations indicated that Fe(III)-oxalato complexes can often
1.0 <1.1 ·o~ Fe(OHf2+ ~ „ c. 0.8 <1.1 ' ' f eox+ .... \ .... \ -...... 0.6 ' \ ! ~ \ .' ~ ,.1 ,'So ....0 0.4 ,·; \ ; ' = ; \ s:.... ./' ~ 0.2 ...C':S .,.~· " ~ 0.0 -7 -6 -5 -4 -3 Log [Ox] (total)
Figure 1.1 Mole-fraction distribution of the iron (J/l)-oxalato comp/exes based on the stability constant of Fe(JJJ)-oxalato species and Fe(JII)-hydroxo species (298 K, pH 4, [Fe(lll)] = 1.0 µM, ionic strength: 0.03 M NaC/04). Ox standsfor oxalate. - 5 -
is shown in become the predominant species of Fe(ID). A typical calculation Figure 1.1. Fe(IIl)-oxalato II) (a) In the absence of 02, the photoreduction of of Fe(III) complexes is first order with respect to the concentration ratio of 1:2 to Fe(II) (Figure l.2a). Oxalate is depleted in a stoichiometric photoreactivity. Tue formed. Various ferrioxalate species have similar the reduction change of pH in the range of 1 to 5 only slightly influences Fe(II) formed is rate. (b) In the presence of dissolved oxygen, the complexes. reoxidized to Fe(III), leading to reformation of Fe(III)-oxalato is rapidly A photostationary state for the redox cycle of Fe(III)/Fe(II) established (Figure l.2b).
0 10 ..-.,s; b ~ ..-. ~ ~ ~ 8 :::1. CU -r.x...... -1 ,....., ..-. 6 ,....., -~ ..-. ~ --~ '-' ~ CU ~ r.x.. 4 CU ...... -r.x.. -2 ...... 2 = - 0 20 40 60 80 0 100 200 300 400 Sunlight exposure time (sec) Sunlight exposure time (sec)
data of Figure 1.2 a) First-order treatment of photoreduction = 10 Fe(lll)-oxalato complexes in deoxygenated solutions. [Fe(lll)]o 0.03 M µM., [oxalate]o = 30 µM., pH = 4.00 ±0.05, ionic strength = performed at (NaCl04, HCl04), deaerated using Ni. Irradiation was solar noon on 13thAug.1989 (Io = 0.70 mE m-2s-1); was b) photoproduction of Fe(ll) in air-saturated solutions. Irradiation m-2s-1 ). performed at solar noon on 15th Aug. 1989 (Io = 0.74 mE . 6 . '' ''
III) (a) Hydrogen peroxide is generated by photolysis of Fe(III)- oxalato complexes in the presence of oxygen (Figure 1.3).
50 1.0 ,....,= 0.8 QJ 40 ....CO CO ~ -~ -:::1. 0.6 '-' 30 0...... ,...., ._,...., N QJ 0 CO N 20 0.4 -CO ...... -~ = 0...... 10 0.2
0.0 0.1 0.2 0.3 0.4 Sunlight exposure (Einstein m-2)
0 2 4 6 8 10 Sunlight exposure time (min)*
Figure 1.3 Photochemicalformation of H202 and depletion of oxalic acid with sunlight. The solution composition is the same as for Figure 1.2, except for [oxalate]o. Irradiations for [oxalate]o: 60 µM (squares) and 120 µM (circles) were peiformed on 18th Sept. 1989 (Io = 0.64 mE m-2s-l and 0.63 mE m·2s-l, respectively), for [oxalate]o: 10 µM (diamonds) and 240 µM (triangles) on 22nd Sept. 1989 (Io = 0.55 mE m·2s-l and 0.36 mE m·2s·l, respectively ). *Time scale corresponds to Io = 0.63 mE m·2s-l. - 7 -
(b) Tue rate of H202 formation increases with sunlight intensity in low intensity range, and with oxalate and Fe(ID) concentrations within the ranges used. These phenomena can be successfully explained by following mechanisms (taking monoferrioxalate as an example):
Fe(ill) + Ox -7 Fe(IIl)Ox (1.5)
hv Ox Fe(ill)Ox -7 FeOx* -7 Fe(II) + Ox·- (1.6)
JW Ox·- + 02 -7 02··JH02· + C02 (1.7)
Or/H02· + Fe(II) -7 H202 + Fe(III) (1.8)
02··1H02· + Fe(ill) -7 02 + Fe(II) (1.9)
(c) Tue rate of H202 formation is 3.7 nM s-1 when a solution of 1 µM Fe(III) and 5 µM oxalate at pH 4 was exposed to September noon sunlight. This rate is similar to or greater than that expected from the dissolution of gaseous H202 and disproportionation of H02 radicals scavenged by cloud droplets under similar environmental conditions.
IV) (a) Tue photolysis of oxalic acid in the presence of ferric ions shows the characteristics of a catalytic reaction. Tue decomposition is first order at low concentrations with respect to oxalic acid and becomes zero order at high concentrations (Figure l.4a). Tue rate of oxalic acid decomposition increases linearly with increasing of amount of initial Fe(III) - 8 - up to a level corresponding to complete complexation of oxalic acid as depicted in Figure 1.4b, or to complete absorption of the incident light potentially absorbable by Fe(ill)-oxalato complexes. (b) In September noon sunlight, oxalic acid is photolyzed with a half life of a few minutes.
400 40 a b m ,.-. c:.i 30 Q> 300 100 200 300 2 4 6 8 10 12 [OX]o (µM) [Fe(III)Jo (µM) Figure 1.4 a) Rate of decomposition of oxalic acid with sunlight as a function of the initial concentration of oxalic acid. [Fe(llf)]O =10 µM, ionic strength =0.03 M (NaC/04, HCl04), pH = 4.00 ±0.05, air saturated, at 283-300 K. Exposure time: five minutes at solar noon on 28th March 1991 (Io = 0.70 mE m-2s-l) b) Rate of decomposition of oxaüc acid in sunlight as a function of the initial concentration of Fe(///). [Ox]o = 5 µM, ionic strength = 0.03 M (NaC/04, HC/04), pH = 4.00 ± 0.05, air-saturated, at 283-300 K. Exposure time: two minutes at solar noon on 22nd April 1991 (Io = 0.78 mE m-2s-l) - 9 - of H202, V) Irradiation of real fog waters has shown the fonnation DOC (dissolved oxidation of S02, and the depletion of oxalic acid and demonstrate the organic carbon) (Figure L5 a and b). These experiments the production of importance of in-cloud photochemical reactions in photooxidants and transfonnation of atmospheric pollutants. 4 ,-. ~ :::1. 2.0 3 -'N 0 2 N =...... , 1.5 1.0 .____,_,___. _ _.___._...... __..__...... __..__..___,_ 0 0 30 60 90 120 150 Time (min) 78 76 74 72....__...___._.....__..._..___.___._ _.___.,___. 0 30 60 90 120 150 Time (min) Figure 1.5 a) Photochemical formation of H202 and decomposition S02 in of oxalic acid infog water. b) Photo-oxidation of dissolved in MGRR fog water. Irradiation was performed with 313 nm light (lo = 1.2 µE 1-1 s-1 ). - 10 - 2. lntroduction and background 2.1 Significance and sources of H202 in atmospheric Iiquids Tue occurrence of hydrogen peroxide (H202) in atmospheric water has significant consequences for atmospheric chemistry. Numerous experiments have shown that H202 plays an important role in the acidification of cloud, fog and rain water through its participation in the oxidation of dissolved S02 to H2S04 in the atmospheric liquid phase (Calvert et al., 1985; Chandler et al., 1988; Gervat et al., 1988; Fung et al., 1991; Hoffmann and Edwards, 1975; Martin et al., 1989). This process is particularly important because H202 is highly soluble and the oxidation of S02 by H202 is relatively fast in acidic solution, whereas the oxidation of S02 by 03 and the catalytic oxidation of S02 by 02 in the presence of Fe and Mn is retarded at lower pH. The nonlinearity between S02 emission and the wet deposition of sulfate suggests that a limiting factor in the acidification of precipitation is often the amount of H202 available to the cloud-precipitation system (Meagher et al., 1990). In addition, H202 formation could be involved in radical chemistry, being a sink of H02· and a source of OH radicals chiefly through the reactions in gas phase: 2 H02·(g) ~ H202(g) + Ü2(g) (2.1) hv (A. ~ 350nm) H202(g) ------~ 2 OH·(g) (2.2) (2.3) and in liquid phase by the reactions: - 11 - (2.4) H102 + M(n-1)+ ~ Mn+ +OH· + OH- (2.5) hv (A. ~ 350nm) (2.6) H102 ------~ 2 OH· ion where M(n-1)+ is reduced metal ion and Mn+ is oxidized metal 1986; (Schwartz, 1984; Chameides and Davis, 1982; Graedel et al., 1991). McElroy, 1986; Zepp et al., 1991; Weinstein-Lloyd and Schwartz, of OH radicals are the most reactive oxidizing species for the transformation atmospheric inorganic and organic pollutants. in lt is now widely believed that the damage to forests by oxidants damage North America and Central Europe may be more significant than the 1986; caused by acid rain deposition (Mallant et al., 1986; Masuch et al., at Möller, 1989; Ennis et al., 1990). This includes damage by H102 concentrations typically measured in clouds, fogs and rainwater. All of these observations encourage much interest in the understanding H102 in of the mechanisms responsible for the formation and degradation of the gas the atmospheric liquid phase. H102 has been extensively measured in as phase and in the atmospheric aqueous phase, in cloud-, rain- and fogwater et al.; weil as in snow and ice (Daum et al., 1990; Olszyna et al., 1988; Barth reviews 1989; Sigg and Neftel, 1991; Neftel et al., 1984; Tsai et al. 1991; for 1992). see Kelly et al., 1985; Gunz and Hoffmann, 1990; Zuo and Hoigne, and Tue typical concentration range of H102 is 0.1-2 ppb in the gas phase 10-7-l0-4 M in rain and cloud water, with the highest values generally observed in the summer. - 12 - Tue presence of H202 in the atmospheric liquid phase is usually assumed to arise solely from the dissolution of gas phase H202, and the disproportionation of H02 radicals, which are produced in the gaseous phase and scavenged by cloud droplets (Lelieveld and Crutzen, 1990; Bufalini et al„ 1979; Simonaitis et al. 1991; Becker et al., 1990; Schwartz, 1984; Chameides and Davis, 1982; McElroy, 1986). However, Kleinman (1984, 1986) has pointed out that the amount of H202 predicted by existing gas phase photochemical models of the boundary layer (1 ppb) is insufficient in many instances to account for the in-cloud oxidation of S02 in summer rainwater. Zika et al. (1982) did not observe any time dependence of the H202 concentration in rain during storms in South Florida and the Bahamas. They concluded that a substantial fraction of H202 in this case was generated within the cloudwater. lt has also been reported that H202 can be produced in an aqueous solution through ozone decomposition, but these reactions are very slow in atmospheric water droplets (Graedel and Goldberg, 1983; Hoigne and Bader, 1975; Heikes et al., 1982; Zika and Saltzman, 1982). Kormann and co-workers (1988) have suggested the photo-sensitized production of H202 by certain sunlight-absorbing semiconductor solids acting as sensitizers in cloudwater. However, the limited concentration data available for solid photosensitizers in precipitation suggest that this mechanism of H202 formation would require exceptional circumstances to be significant. lt is clear that the sources of H202 in the atmospheric liquid phase are still far from being established. Surprisingly, the in-cloud liquid phase photochemical formation of H202 has received scant attention. Photochemical electron-transfer processes involving 02, organic substances, and transition metals have been postulated to generate superoxide ions (02·-) in sea water (Petasne and Zika, 1987). 02 .- and its conjugate acid, the hydroperoxyl radical (H02·), readily - 13 - such undergo disproportionation to yield H202 and 02 in aqueous solutions is well as atmospheric water (Weinstein-Lloyd and Schwartz, 1991). lt such as known that transition metals such as Fe, Mn, Cu and organic ligands (for oxalic and pyruvic acid are common constituents of atmospheric water Sigg, review see Graedel et al., 1986; Zuo and Hoigne, 1992; Bebra and et al., 1990; Joos and Baltensperger, 1991; Steinberg et al., 1985; Norton 1983; Johnson et al. 1987; Xue, 1991). In this study, we take aqueous Fe(III)-oxalic acid-02 as a model system to investigate the mechanisms and kinetics for the photochemical formation of H202 and the decomposition of oxalic acid in sunlight under experimental conditions which represent chemical conditions of acidified atmospheric water. of 2.2 Proposed mechanism for the photochemical formation H202 and decomposition of oxalic acid in atmospheric droplets Most of the proposed mechanisms for the production of H202 in the to aqueous phase involve the single electron reduction of molecular oxygen in form the intermediate superoxide ion (02·-). The 02·- is always equilibrium with its conjugate acid, the hydroperoxyl radical ( H02·)(Rabani and Nielsen, 1969; Behar et al., 1970; Bielski, 1978): (2.7) and the equilibrium constant KH02· = 1.6 X 10-5 mol /-1 . 14. Tue subsequent disproportionation of H02· and 02·- leads to H202 (2.4a) (2.4) (2.4b) where H02- is the conjugate base of hydrogen peroxide. Tue reaction rate constants for H02·/02·· and H202 with 03, Fe and Cu ions are given in Table 2.1. Tue rate of disproportionation depends largely upon pH because the reaction between two molecules of H02· proceeds more slowly than the reaction between H02· and 02··. Tue reaction between two superoxide anions was shown to be extremely slow. With specially purified reagents, and with the complexing agent EDTA present to sequester possibly catalytic metal ions, Bielski and Allen (1977) could not observe this reaction proceeding at all (k < 0.3 M-1 s-1). As a consequence, the disproportionation rate of these radicals, nearly constant at pH less than 2, rises to a maximum near the pKa 4.8, where the H02· and 02·· are of equal concentration, and falls an order of magnitude for each unit the pH is raised beyond the pKa due to the progressive decline in the concentration of H02'. Transition metal ions such as Fe, Cu and Mn can significantly influence this reaction. The oxidized metals react with H02· and 02·· to form dioxygen, and the reaction of H02· and 02·· with reduced metals produces H202. Tue reactions with Fe(III) and Fe(II) can be described by the equations: (2.8) - 15 - (2.8a) Oy + Fe(Ill) --+ 02 + Fe(II) (2.9) H02· + Fe(ID) --+ 02 + Fe(II) + H+ (2.9a) They show that in the acidic pH range, the reaction with Fe(ID) is very slow; the main fate of H02 radical would be to react with Fe2+ to produce H202 when the concentration of Fe(II) is higher than that of hydroperoxide radical (see table 4). In addition to the scavenging from the gaseous phase, H02· and 02·- could be formed by several photochemical mechanisms in the aqueous phases. Light-absorbing organic substances can reduce oxygen either by direct electron transfer from the excited state of the organic compounds or by the generation of free hydrated electrons via photo-ionization. However, the production of free electrons is usually very slow in natural waters (Zafiriou, 1977; Zika, 1981; Draper and Crosby, 1983; Sturzengger, 1989; Zepp et al., 1987). lt is also possible in natural waters that a photochemically induced electron transfer from complexing organic ligands (org.) to oxidized metals (Mn+) occurs and, subsequently, the electron deficient organic ligands further reduce 02 to 02·-. This pathway can be represented as: hv Mn+-org. --+ M(n-1)+ + org. radical (2.10) org. radical + 02 --+ 02·- + oxidized org. (2.11) In the present study, Fe(ID)-oxalato complexes were used as a model system to test this postulated mechanism for the photochemical formation of H202. - 16 - Table 2.1 Reactions and reaction rate constantsfor H02·/02·- and H202 with 03, Fe and Cu ions k ref. Conc.* k'** (M-ls-1) (M) (s-1) H02· reactions: H02·+HOi H202+02 8.3 X 105 [l) 10-8 8.3 X 10-3 H+ H02·+ 0 2·- H202+02 9.7 X 107 [l] 10-9 0.1 HOi + Fe(III) -H+ Fe(II) + 02 <104 [2) 10-6 02·- reactions: H+ Oi-- + H02· H202+02 9.7 X 107 [l] 10-8 1 02·- + Fe(III) Fe(II) + 02 1.5 X 108 [2) 10-6 150 H+ 02·- + Fe(II) Fe(III) + H20 2 1.0 X 107 [2] 10-6 10 02·- + Cu(II) Cu(I) + 02 8X109 [3] 10-9 8 H+ 02·- +Cu (1) Cu(II) + H20 2 7X109 [3] 10-9 7 H+ 02·-+Ö3 202 +0H· 1.5 X 109 [1] 109 1.5 H20 2 reactions: H202 + Fe(III) very slow (4] (4] 10-6 7.6 X 10-5 H20 2 + Fe(II) Fe(III) + OH· + OH- 76 * Assumed concentration values of the reactantsfor H02·/02·- and H202for noon summer scenario at pH4; values are given only for exemplification of reaction rates and their competition. ** Pseudo-first order reaction rate constant based on k and the arbitrary concentration values of the reactantsfor H02·/02·-, H202 and 03. [l] Bielski et al., 1985; [2] Rush and Bielski, 1985; [3] Von Piechowski, 1991; [4] Walling, 1975. - 17 - - 18 - The photochemistry of Fe(ID)-oxalato complexes has been widely studied, particularly in connection with UV and visible chemical actinometry (Balzani and Carassiti, 1970; Parker and Hatchard, 1959; Hatchard and Parker, 1956). The absorption of a photon by an Fe(ID)-oxalato complex results in an electron transfer from a complexing oxalate ligand to the central ferric ion, producing a ferrous ion and an oxalate radical anion. The oxalate radical anion could reduce a further ferric oxalato ion. In the presence of 02, however, the oxalate radical anion would react with 02 to produce 02·· in dilute solutions. Taking the mono-oxalato ferric complex as an example, the mechanism could be expressed by the scheme presented in Figure 2.1. Analogous reactions can be written for the di- and tri-oxalato ions. 2.3 Occurrence and speciation of Fe(IIl)-oxalato complexes in atmospheric waters The photochemistry of any aqueous Fe(ID) complex depends on its speciation because different aqueous Fe(III) species exhibit different absorption spectra and photoreactivities. For instance, although the hexaaquo Fe(III) complex (dominant at pH < 3) photolyzes to produce OH radicals indirectly, its charge transfer band does not significantly overlap the tropospheric solar spectrum. The Fe(OH)2+ complex (dominant at pH 3-5 in the 1;tbsence of organic ligands) and other hydroxide ion complexes have absorption bands that extend into the 290-400 nm region, overlapping the tropospheric solar spectrum. They are therefore considered to be a major photochemical precursor for the formation of OH radicals in the atmospheric liquid phase (Faust and Hoigne, 1990; Benkelberg et al., 1990). - 19 - + 0 CD hv o-c:-·o Fe{III),... o-c~1 ]+ ~ Fe{II)(' 1{ [ 'o-~o e- transfer [ ·o-o,.o ] © + + - ~ u. ~{ 1 ~ (.)- Figure 2.1. Scheme of possible photochemical reactions in the Fe(l/I)-oxalato-oxygen system. Fe(III)-oxalato complexes absorb even more strongly in the tropospheric solar UV-visible region (295-570 nm) (see Figure 2.2), and are also photochemically more reactive. - 20 - 6000 6 .,... ":- e -E \ c .,...(,) 5000 \ FeOx2-/FeOx33- 5 .,... ~ \/ N"' c \ FeOx2• 4 e(,) Q) 4000 \ - X ex> \ :::J C') ;g =o Q) 3 c..- -0 3000 (,) ,g >< c _g ..... 0 :;:::; 2 a...,..: (,) c >< ~ c X "ä) Q) 1 ,_ 1000 Cf) as -c ö jjj ~ 0 0 280 320 360 400 440 480 - Wavelength (nm) Fe(OH)2+ Figure 2.2. Absorption spectra of Fe(/1/)-oxalato complexes, of Fe(OH)2+ is and Fe3+ in the region 290-500 nm. The absorption spectrum complexes obtainedfrom Faust and Hoigne (1990). Thosefor Fe(Ill)-oxalato and Fe3+ were measured in this work using standard spectrophotometric and Pitts techniques. The solar spectrum is obtained from Fin/ayson-Pitts 1 st. The (1986), and refers to a /atitude of 47.4 ON at noon on September points indicate calculated averages over 5or10 nm intervals. 1.10.0 µM Fe(C/04)3 in 0.1 N HC/04 (Fe3+ dominant), 2. 10.0 µM Fe(C/04)3, pH 4 (Fe(OH)2+ dominant) dominant), 3. 10.0 µM Fe(C/04)3, 30 µM K1C204, pH 4 (dioxalato 4. 10.0 µM Fe(C/04)3, 120 µM K1C204, pH 4 (tri- and di-oxalato) All solutions contain 0.03 M NaC/04 adjusted ionic strength. - 21 - 2.3.1 Occurrence of iron in atmospheric waters Iron is a ubiquitous constituent of atmospheric waters and particles. lt may be derived from both natural and anthropogenic sources. Tue principal natural sources are wind-borne soil particles, volcanoes and forest fires. Tue anthropogenic sources mainly illclude industrial activities, fossil fuel and solid wastes burning. Among these sources, wind blown dust and fuel fly ash are the most significant. Tue average concentration of iron in soils is 3.8% (Bowen, 1966). Coupling this value to a flux of continental dust wafted into the atmosphere annually of 5 x 1Q14 g (Goldberg, 1971), one can estimate a flux of iron, 0.12 g m-2 yrl, transported from continental surfaces to the atmosphere. Tue content of iron in fuel fly ash often represents a significant fraction (2-15%) of the ash weight (Foster, 1969; Brimblecombe and Spedding, 1975; Hansen et al., 1984; Dlugi et al., 1985; Williams et al., 1988). Furthermore, iron is often enriched at the surface of fly ash particles (Hansen et al., 1984). Aerosol particles from Wuhan (China), Beijing (China) (Waldman et al., 1991), New York (U.S.), Deuselbach (Germany), Karlsruhe (Germany) (Dlugi et al., 1985), and the Shenandoah Valley (U.S.) (Tuncel et al., 1985) were also found to contain substantial amounts (0.6- 3.0% w/w) of iron. These iron-containing particles, like other airborne dust, are easily associated with atmospheric water droplets by serving as cloud condensation nuclei or by scavenging into cloud, fog and rain droplets. Table 2.2 summarizes collected data on the concentrations of iron in rain, fog and snow waters measured at a wide variety of locations. Concentrations of iron in rain and fog waters are most often determined by atomic absorption spectroscopy (AAS), neutron activation analysis (NAA), or inductively coupled plasma emission spectrometry (ICP). Tue published concentration results are mostly the total iron, including the dissolved and particulate species. - 22 - Table 2.2 Total iron concentrations in rain,fog and snow waters Conc. Sample Technique Location Time of yr. Ref. (µM) 2.0-220 Fog AAS &Col. Düb. Oct.-Feb. [l] [2] 0.3-2.3 Rain AAS &Col. Düb. Jan.-Sec. [1] [2] 0.1-1.6 Snow AAS &Col. Düb. Nov.-Apr. [2] 0.3-91.l Fog ICP NWZ Sept.-Dec. [3] 1.6-115 Fog AAS LA Nov.-Jan. [4] 1.0-29 Fog AAS Po Valley Oct.-Mar. [5] 2.3-55 Rain NAA Wray. Jan.-Dec. [6] 0.3-23 Rain AAS Leeds Jan.-Dec. [7] AAS = atomic absorption spectroscopy, Col. = colorimetry, ICP = inductively coupled plasma emission spectrometry, NAA = neutron activation analysis, Düb. = Dübendorf (Zurich), NWZ = northwest of Zurich, LA= Los Angeles, Wray. = Wraymires. [l] Behra & Sigg, 1990; [2) this work,· [3] Joos & Baltensperger, 1991; [4] Munger et al, 1983; [5] Fuzzi et al, 1988; [6] Peirson et al, 1974; [7) Clarke & Radojevic, 1987. In this study, we restrict our investigation to the reactions occuring in homogeneous system. What portion of iron is soluble in moderately acidic clouddrops? In atmospheric waters particulate iron could be dissolved by several pathways. That the dissolution is promoted by protons, ligands, reductants and photons has been well established (Stumm and Morgan, 1981; Schneider, 1988; Sulzberger et al., 1989; Siffert and Sulzberger, 1991; Deng, 1992). Large fractions (12-56%) of the iron in fuel ash were found to dissolve rapidly (within 5-20 min) in water at pH 2 (Dlugi et al. 1985) and - 23 - pH 2.8-4.0 (Brimblecombe and Spedding, 1975). Similar results have been obtained by Williams et al. (1988) and by us (unpublished data). Recently, the fraction of Fe(II)dis/Fetot (subscript "dis" represents "dissolved" and "tot" is "total") (20-90%) in fog water collected in Dübendorf during the winter time has been established by Bebra and Sigg (1990). Tue fractions of FedisfFetot and Fe(II)dis/Fetot-dis for 8 rain events, soluble Fe/Fetot and soluble Fe(Il)/soluble Fetot for 5 snow events in Dübendorf were also measured (see Table 2.3). Table 2.3 Percentage of dissolved ironltotal iron and dissolved Fe(II) ldisso/ved total iron in Dübendorffog, snow and rain waters Sample Fedis!Fetot Fe(II)disfFetot Fe(II)dis/Fetot-dis Time ofyr. Ref. (%) (%) (%) Fog 20-90 20-90 Oct.-Feb. [l] Rain 49-92 10-82 49-95 Apr.-Dec. [2] Snow 21-32 15-18 74-98 Nov.-Apr. [2] subscript "dis" represents "dissolved", and "tot" is "total", but for snow samp/es "dis" expresses solub/efraction. [l] Behra and Sigg, [2] this work. Tue high percentage of Fe(II) in these precipitation samples is mainly due to photochemical reduction of particulate and dissolved Fe(III) species and occurrence of dissolved S02 (Siffert and Sulzberger, 1991; Faust and Hoigne, 1990; Bebra and Sigg, 1990; Zhuang et al. 1990 and 1992)~ - 24 - Although there is an abundance of Fe measurements, little information exists on the speciation of Fe(ill) in atmospheric waters. Calculations by Faust and Hoigne (1990) and by Weschler et al. (1986) indicate the pH dependent speciation of Fe(ill)-hydroxo complexes in the absence of organic ligands. 2.3.2 Occurrence of oxalic acid in atmospheric waters Chelate ligands are also abundant and ubiquitous in atmospheric droplets. Of them, oxalic acid is characteristic and important. Oxalic acid is mainly produced by ozonolysis and photooxidation of hydrocarbons in the gas phase or in atmospheric liquids. Tue oxidation of aromatic compounds to glyoxal and oxalic acid was found in studies of ozonation, ozonation in combination with UV radiation and photooxidation of aromatic compounds in aqueous solutions (Norton et al., 1983; Leitis, 1979; Ho, 1986; Kusakabe et al., 1990; Yamamoto et al., 1979). Further oxidation of glyoxal produces glyoxylic acid and oxalic acid (Leitis, 1979). Direct emission from incomplete combustion is another important source of oxa lic acid. Tue concentration of oxalic acid (3-5 nmol m-3) was found even slightly higher than those of formic and acetic acids (1-4 nmol m-3) in the haze layers (at 1-4 km altitude) derived from biomass-buming. This oxalate concentration is 14-24 times higher than that in the regional background air over Amazonia (Andreae et al., 1988). Tue concentrations of oxalic acid in motor exhaust, 182-482 nmol m-3, were also established (Kawamura and Kaplan, 1987), which are similar to the concentration of fonnic acid, 378 nmol m-3, in the motor exhaust. These concentrations are 28-144 times higher than the concentration of oxalic acid (2.12-8.65 nmol m-3) in Los Angeles ambient air (Kawamura et al., 1987). lt is worth noting that the concentration of oxalic acid in Los Angeles ambient air is comparable to that - 25 - expected for gaseous H102. Wind blown dust and emission by vegetation may be additional sources of atmospheric oxalic acid. Table 2.4 Concentrations of oxalic acid in cloud, rain, fog and snow weiters Ref. Conc. Sample Technique Location Time of yr. (µM) 3.4 Rain IC Divide, CO* July-Aug. [1] 4-23.3 Mist D & PHLC Los Angeles June [2] 17.3-18.7 Fog D & PHLC Los Angeles June [2] 2.8-20.7 Rain D & PHLC Los Angeles June [2] 1.27-28.3 Rain D & GC Los Angeles Apr.-Nov. [3] 20-25 Fog D & GC Los Angeles June [3] 6.0 Mist D & GC Los Angeles June [3] 2.5-5 Cloud IC NW Zurich Aug. [4] 7.5-24 Hili fog IC NW Zurich Oct. [4] 5.5-14 Ground fog IC NW Zurich Oct. [4] 1.1-24.4 Fog IC NW Zurich Sept.-Dec [5] 0.11-10 Fog IC Dübendorf Oct.-Feb. [6][7] IC: Ion chromatography; D: Derivatization; GC: Gas chromatography; HPLC: High presure liquid chromatography. [l] Norton, 1983, These samples were collected from a remote mountain area in Colorado. [2] Steinberg, 1985; [3] Kawamura, 1985; [4] Baltensperger and Kern, 1988; [5] Joos and Baltensperger, 1991; [6] Sigg et al„ 1991; [7] This work The decomposition of oxalic acid in the gas phase is very slow since it does not react with ozone (Leitis, 1979; Hoigne and Bader, 1978). lt reacts very slowly with OH radicals (Buxton et al., 1988), and its photolysis is very slow . 26. under tropospheric conditions (Y arnarnoto, 1985). On the other band, oxalic acid is highly polar. lt will be preferentially transferred into the liquid phase (at pH 4, the apparent Henry's law constant is 2.4 x 1011 M atm-1 ). Table 2.4. presents the data on concentrations of oxalic acid in rain, fog, mist and snow waters measured at several locations. The presence of chelate ligands, such as oxalate, has a significant effect on the speciation and photoreactivity of Fe(Ill) ions in acidic atmospheric water since they may form stable complexes with Fe3+ ions: (2.12) (2.13) (2.14) (For simplicity, water molecules in the coordinated sphere are not shown in the chemical formulae throughout this paper.) The speciation of dissolved Fe(III) in aqueous solution depends on the competition between the formation of Fe(Ill)-oxalato complexes and of Fe(Ill)-hydroxo complexes. Speciation calculations show that under the conditions of acidic atmospheric water droplets, where the pH ranges from 3-5, Fe(Ill)-oxalato complexes could be the predominant dissolved species (Figure 2.3). In contrast, Fe(II) mainly occurs as the hydrated cation (Figure 2.4 ). 1.0 „„ .,.,.. -- ,/ FeOx33- ;'\ I f/) Q) ,' 0.8 •i \ \ I "ö Q) ; \ ,' c. FeOx+/ \ / f/) A l \ l " i \ ' 0.6 l \ ; \ ,' : \; \ ' -Q) ! ~i \1 u.. : '1' 1, : ,1. ' 0 ; ., ' ' - 0.4 • ..j\ : \ c: . ' \ 0 . . :i .l '· ,1 \ \ ·u .. i \ : \ ro / 1 \ 1 \ .... • i \ ' \ U... 0.2 / \ \ : "· ~,'/ .\, ' „., Fe3+ ,, \1 .„ ·'·„. ··-··-··-··-··-·····-··__ .„ ,::„._____ I ,•„„ "'-... „ •·„.... „ ••„„. 0.01.---...-=.::.:...... ~...L.-.-..:::::::::::-..~.;:::.:....._~--~--=.:.:::::::aa.~--_J -8 - 7 -6 -5 -4 -3 -2 - 1 Log cOx 1 (total) Figure 2.3. Mole-fraction distribution of the iron(lll)-oxalate complexes based on the stability constants of the Fe(l/1)-oxalato species and Fe-hydroxo species (298 K, pH 4, [Fe(lll)] = 10.0 µM, 0.03 M NaC/04). Ox standsfor oxalate. - 27 - - 28 - 1.0 <12 ·o~ Fe 2+ ~ Cl. 0.8 ... <12 ~ ~ . 0.6 ~ -~ eo.. Q 0.4 ....=Q ~ -= 0.2 FeOx ~"" 0.0 ------0 50 100 150 200 250 [Oxalate]total (µM) Figure 2.4 Mole-fraction distribution of the Fe(J/)-oxalato complexes based on the stability constants of the Fe(J/)-oxalate species and Fe(/1)- hydroxo species (298 K,pH 4, [Fe(ll)] = 10.0 µM, 0.03 M NaC/04) - 29 - 3. Experimental methods 3.1 Materials 14C-labeled oxalic acid with a radiochemical purity of 98% and specific All other activity 4.1 mCi/mg was obtained from New England Nuclear. chemicals were of reagent grade and were used without further purification. doubly Tue water employed in all preparations was first deionized and then in the distilled. All stock solutions were prepared frequently and were stored solution dark. Tue photolysis tubes were thoroughly cleaned with 1 M HCl füll and then finally with deionized-doubly distilled water. They were stored with of deionized-doubly distilled water. Before irradiation they were rinsed the solution to be irradiated. 3.2 Irradiation procedures stock Solutions for irradiation were freshly prepared from air-saturated to solutions of Fe(Cl04)3 in 0.1 M HCl04 and potassium oxalate, and adjusted noted, all the desired pH with either HCl04 or NaOH solutions. Except where the ionic solutions for irradiation contained 0.03 M NaCl04 (to adjust and were strength), 10.0 µM Fe(III), various concentrations of oxalate, bubbling saturated with air. Deaeration, when desired, was accomplished by for at water-saturated high purity N1 through the solution in photolysis tubes least 15 min before irradiation. Photolyses of individual samples and actinometers were performed tubes simultaneously in separate glass-stoppered cylindrical quartz photolysis (1.5 cm i.d., containing 20 ml of solution) in the sunlight or in the merry-go- round reactor. - 30 - 3.2.1 Solar irradiation Irradiations in sunlight were carried out in the quartz photolysis tubes held in a rack at a 30° angle from the horizontal and about 1 m above a black pavement. With this rack, the photolysis tubes could be put in and removed from sunlight within one second. All irradiations were centered on solar noon at EA W AG (Dübendorf, Zürich; 47.4 °N 440 m elevation). The ambient temperature was 10-30 °C. Tue sunlight intensity (in mE m-2 s-1, where E stands for Einstein) was recorded every 10 minutes using a quantum sensor (Li-COR, LI-185) that responds to light of 400-700 nm. This measurement gives a value of 1.8 mE m-2 s-1 for a clear sky at solar noon in June. Valerophenone was also used as a sunlight actinometer (responding to the UV B region) (Faust and Hoigne, 1990; Zepp, 1989; also see § 3.2.3). 3.2.2 Photochemical reactor and monochromatic irradiation A merry-go-round reactor (MGRR) equipped with a Hanau TQ 718 high pressure Hg immersion lamp (500 W) was used for all indoor photolytic experiments. In this MGRR, the Hg lamp (light source) is in the center of the apparatus with photolysis tubes arranged in a ring around it. The ring is fixed with a Dewar vessel. lt rotates around the light source at a speed of 11 rpm to give a uniform irradiation of the photolysis tubes. When the stirring is desired, a magnetic stirring bar is put in the photolysis tube. lt is lifted up on each revolution by the suspended magnet. A schematic diagram of this MGRR is shown in Figure 3.1, and the emission spectrum of the Hanau TQ 718 high pressure Hg lamp from 290-600 nm is given in Table 3.1. The 313 nm line of the Hg lamp was isolated by filtering the lamp output through a solidex borosilicate glass sleeve and a 2.0 mM K1Cr04 in 3% K1C04 aqueous filter solution present in the Dewar vessel. The minimum light pathlength of the - 31 - Figure 3.1 A schematic diagram of the merry-go-round reactor photolysis tube support / Magnet Magnetic stirring bar UV-lamp Drive box - 32 - Table 3.1 Ultraviolet-visible emission spectrum of Hanau TQ 718 high pressure Hg lamp (290-600 nm given by manufacturer) Wavelength (nm) Intensity (µEs-1) Wavelength (nm) Intensity (µEs-1) 297 8 405-408 36 302 16 436 51 313 38 492 1 334 5 546 78 366 66 577-579 76 390 1 100 ___ ...... „-„ - „„ ... „„„„... ------/ ,....._ 80 ,/ ~ ,, '-' ,, 1 8 60 ,. ~ , '§ • 40 .• e"'c ,,• ~ ,• 20 ,,. ,1 ,• 0 200 300 400 500 600 700 Wavelength (nm) Figure 3.2 UV and visible light transmission spectra of the solidex sleeve (dash Une) and a 2.0 mM K1Cr04 + 3% K1C03 aqueausfilter solution (solid line measured in this work) - 33 - filter solution was 2.1 cm. Figure 3.2 illustrates the UV and visible light transmission spectra of the solidex sleeve and the filter solution. Irradiations at 313 nm were performed using the saine photolysis tubes as for solar irradiation. Tue reaction temperature was maintained at 20 ± 1 ° C by recirculation of the filter solution through a thermostat, and the lamp was cooled with tap water passed through the space between the lamp and the solidex borosilicate glass sleeve. Tue correction factors for the above mentioned photolysis tubes were determined by a valerophenone actinometer. 20 ml of 12 µM valerophenone in water in each tube was subjected to photolysis for 45 min at 313 nm light, and the loss of reactant was quantified by HPLC. For the photolysis with white light in the MGRR or under sunlight, the deviation between different tubes was less than 2%, and therefore no intercalibration was made. 3.2.3 Actinometry and quantum yields Aqueous valerophenone (12 µMin air-saturated solution) was used as a chemical actinometer for solar irradiations and 313 nm monochromatic photolyses. This actinometer is reactive under ultraviolet light and very stable under common indoor illumination conditions. Tue reaction was followed by determining the concentration of va'lerophenone remaining, Cact· Tue intensity of monochromatic light (Einstein J-1 s-1) was calculated by equation (3.1) dCactfdt (3.1) IA.o = 2.303DEA,act Cact d[Fe(Il)]/dt (3.2) cl>A.Fe(II) = 2.303De1..[Fe(ID)Oxn]I1..o where d[Fe(II)]/dt is the initial rate of Fe(II) formation at the wavelength A., Dis the light path length (cm), e1.. is the decadic molar extinction coefficient of Fe(ID)-oxalato complexes at wavelength A. (M·l cm·l), [Fe(ID)Oxn] is the molar concentration of Fe(III)-oxalato complexes, and 11..o is the volume- averaged incident light intensity at wavelength A. (Einstein z-1 s-1). I313nm,O was determined with the valerophenone actinometer. 3.3 Analytical techniques H202 analysis was begun immediately after irradiation. All samples were analyzed within 40 min of irradiation. No significant concentration changes were detected during the analyses. The determinations were - 35 - performed using a method based on the horseradish peroxidase (POD)- catalyzed oxidation by H202 of N,N-diethyl-p-phenylenediamine (DPD) (Bader et al., 1988). The radical cation formed, DPD·+, is stabilized by resonance and has a fairly stable absorption with a maximum at 551 nm and molar extinction coefficient (log base 10) e = 21,000 ± 400 M-1 cm-1. This method has been successfully applied to the analysis of hydrogen peroxide in freshwater, rainwater and fogwater samples as well as during several laboratory research projects. For the case of clean water, the standard deviation for this method is 1% and the lower limit of detection is 10 nM. Tue method is also somewhat sensitive to organic peroxides. Tue latter can be distinguished from H202 by examining the reaction kinetics of the radical cation DPD·+ formation, or by the use of catalase. Application of these techniques during our experiments indicated that no detectable organic peroxides were present in the irradiated samples. lt is important to test whether or not the presence of Fe(II) and Fe(III) interfere with the analysis of H202. As shown in Figure 4.1.8, the reaction between H202 and Fe(II) was virtually instantaneous. Tue ferrous ion was oxidized to ferric iron. All the concentrations of H202 measured are accumulated concentrations. That is, these do not include the H202 decomposed either during irradiation or by subsequent rapid dark reaction with residual Fe2+. Tue reason for this is explained in detail in the section of Experimental results and discussion. Tue presence of Fe(III) does interfere with the measurement of H202. The calibration curves for the determination of H202 in the presence and in the absence of Fe(III)-oxalato complexes are given in Figure 3.3. Tue difference between the absorbance values measured in the presence and in the absence of Fe(III) species is negligible. - 36 - 1.0 y =l.0155e-2 + 0.20567x R"2 =0.999 E 0.8 .-1= II) II) ~ 0.6 ~ ~ .c=CIS 0.4 =~""' .c < 0.2 y =8.9074e-3 + 0.20578x R"2 =1.000 0.0 0 1 2 3 4 5 Concentration of H202 (µM) Figure 3.3 Calibration curves for the determination of H202 in the presence and in the absence of Fe(lll)-oxalato complexes. Analysis procedures: Various volumes of H202 standard solution (150 µM) were pipetted into 25 ml of distilled water (open circles) or 25 ml of the solution containing 10 µM Fe(lll), 30 µM oxalic acid, 0.03 M NaC/04 at pH 4 (crosses), and leftfor 30 minutes. Then 3 ml of 0.5 M Na2HP04-NaH2P04 buffer solution, 50 µ11% DPD in 0.1 N H2S04 and 50 µ1 POD (specific activity of 100 units mz-1) in water were added, the mixture was shaken and the absorption spectrum from 650-450 nm was taken 30 ± 5 s after the addition of the PO D. Analysis of Fe(II) was carried out using the modified 1,10- phenanthroline colorimetric method of Tamura (1974), using a value of es10 = 1.105 x 104 M-1 cm·l for the Fe(II)-phenanthroline complex formed. Fe(III) stock solutions were standardized by reducing Fe(III) to Fe(II) with a solution of 1% (w/w) ascorbate, in place of the NH4F solution in Tamura's protocol, and measuring Fe(II). - 37 - In Tamura's procedure the fluoride is added before the phenanthroline. affects the His procedure is successful to a certain extent since daylight presence of determination. However, if the determination is carried out in the zero value reactive oxygen-containing species such as H202, a lower or even of fluoride of the concentration of Fe(II) is obtained because the presence accelerates the oxidation of ferrous ion. For accurate determinations of ferrous ion the following procedure solution, 3 was used: 2 m1 of 0.25% (w/w) 1,10-phenanthrolinium chloride 2 ml of 2 M ml of 0.063 M acetic acid/0.1 M acetate buffer solution and solution was NH4F solution were premixed in a 25-ml brown flask. Sample water. All then transferred into a flask and diluted to 25 ml with distilled influence of stock solutions were stored in brown bottles in order to avoid the kinetic light. This improved procedure is specially useful for fast redox studies of Fe(ffi)/Fe(II). Oxalate was analyzed by the following procedures: liquid a) The ß-counts of 14C-labeled oxalate was determined on a Zurich, scintillation counter (BET Amatic I, Kontron Analytical, or standard Switzerland). One ml aliquots were taken from the sample (Luma Gel, solutions and mixed with 9 ml xylene scintillation fluid determined LUMAC) and counted for 20 minutes. Calibration curves were was used from Standard solutions prepared from the same stock solution that for the in the photolytic experiments. 14C-labeled method is very useful experiments low in concentration of oxalic acid (< 10 µM). b) The DOC (dissolved organic carbon) was determined on a Dohrman oxalic acid DC-80 Carbon Analyzer. The Iower detection limit is about 4 µM (0.1 mg/l carbon). was c) For experiments with real fog water, ion-chromatography a applied with conductivity detection. A Dionex Ionpac AG9 precolumn, - 38 - Dionex Ionpac AS9 colunm and a Dionex anion micro membrane suppressor were used. Aqueous bicarbonate-carbonate solution was employed as eluent. With this system, it is difficult to achieve a good separation of oxalate from the perchlorate ion we used to adjust the ionic strength in the model experiments. Valerophenone (used for chemical actinometry) was analyzed by high pressure liquid chromatography using a methanol/water mixture as the mobile phase and ODS-II as the stationary phase. Measurement of the pH values was made with a Metrohm micro- combination glass electrode with 3 M KCl as the inner reference solution, calibrated with Merck Titrisol pH 2, pH 4 and pH 7 standard buffers. Electronic absorption spectra were obtained on a UVIKON Model 810 UV spectrophotometer. Equilibrium calculation. All speciation calculations in this study were made with MICROQL, a chemical equilibrium program in BASIC developed by Westall (1979). Values of the equilibrium constants were obtained from the literature and corrected for differences in ionic strength using the Davies equation. The composition of the matrix for the Fe(ffi) complexes and the stability constantS are collected in Table 3.2 and for the Fe(II) complexes in Table 3,3. Experimental accuracy. The given points in the figures represent the data from single series of experiments. All experimental points are presented. The error within the series is not greater than 4% for H202 determination, 4% for Fe(II) and 5% for oxalate. Due to variations of sunlight intensities, the errors between series are larger. But under the experimental conditions employed, the experimental results between series are also well correlated to the irradiation intensities. - 39 - Table 3.2. Composition Matrix of the iron (lll)-oxalate system in perchlorate solution and the equilibrium constants Species Components: lgßt (Ref.) Fe3+ C2042- Cl04- H+ Fe3+ 1 0 0 0 0 Fe(OH)2+ 1 0 0 -1 -2.57 [1] Fe(OH)2+ 1 0 0 -2 -7.89 [1] Fe2(0H)24+ 2 0 0 -2 -3.22 [1) FeC204+ 1 1 0 0 9.40 [2] Fe(C204)z- 1 2 0 0 16.20 [2] Fe(C204)33- 1 3 0 0 20.78 [2] FeHC2042+ 1 1 0 1 4.39 [3] FeCl042+ 1 0 1 0 -0.32 [4) C2042- 0 1 0 0 0 HC204- 0 1 0 1 4.21 [5) H2C204 0 1 0 2 5.40 [5) HCl04 0 0 1 1 -1.58 [6] Cl04- 0 0 1 0 0 OH- 0 0 0 -1 -14 H+ 0 0 0 1 0 [l] Faust and Hoigne, 1989; (2) La.croix 1949; (3) Bauer and Srnith, 1965; (4) Sutton 1952; (5) Schapp etal.1954. [6) Hood andReilly, 1960. - 40 - Table 3.3 Composition Matrix of the iron (1/)-oxalate system in perchlorate solution and the equilibrium constants Species Components: logßt (Ref.) Fe2+ C2042- Cl04- H+ Fe2+ 1 0 0 0 0 Fe(OH)+ 1 0 0 -1 -9.6 [1] FeC204 1 1 0 0 3.05 [1] Fe(C204)i2- 1 2 0 0 5.15 [2] FeCl04+ 1 0 1 0 -0.92 [2] C2042· 0 1 0 0 0.0 HC2,04· 0 1 0 1 4.21 [3] H2C204 0 1 0 2 5.4 [3] HCl04 0 0 1 1 -1.58 [4] Cl04· 0 0 1 0 0.0 OH- 0 0 0 -1 -14 H+ 0 0 0 1 0.0 [l] Davison 1979; [2] Martell and Smith, 1976; [3] Schapp etal., 1954; [4] Hood and Rei/ly, 1960. - 41 - 4. Experimental results and discussion 4.1 Formation of hydrogen peroxide by photolysis of Iron(llI)· oxalato complexes and othei" 'iron-carbbxylate systems in. aqueous solutions 4.1.1 Stability of oxalate Oxalate was stable when exposed td 313 nm radiation under anoxic in the conditions. lt was also stable for at least 1 hr in oxygenated solutions or under concentration range used, under 313 nm monochromatic light of Fe(ill) sunlight. In the dark oxalate was not depleted even in the presence nor under air saturated conditions. Neither DOC of solutions decreased hydrogen peroxide was forrned under these conditions. 4.1.2 Oxygen requirements for H202 formation A number of experiments were carried out in a nitrogen atmosphere, that of in which the initial concentration of Fe(ill) was 10 to 100 µM , and 313 nm oxalate 10 to 240 µM. Irradiations were carried out with °C, or in monochromatic light in the MGRR at temperatures of 20 and 30 was sunlight at ambient temperature. Under these conditions no peroxide amount forrned. Tue amount of oxalic acid decomposed was equivalent to the of Fe(ill) reduced, in agreement with the following overall reaction·: hv ~ ,(2n-l)C2042- + 2C02 (4.1) 2(Fe(C204)0 ) (3-2n)+ 2Fe2+ + solutions Tue fact that hydrogen peroxide was only produced in oxygenated aqueous indicates that oxygen is necessary for its forrnation. In air-saturated - 42 - solutions, the oxygen concentration is about 250 µM, much greater than that required for the formation of H202. 4.1.3 Photo-formation of H202 in sunlight Figure 4.1.l shows that in presence of 10 µM Fe(III) and 120 µM oxalate at pH 4, the H202 formation was a linear function of the sunlight illumination time over a wide concentration range of H202. The observed rate of fonnation of H202 in September noon sunlight was 80 nM s-1. ...... = 1 Q,> 40 = Q,> .== s:i. 30 ~ ...=CU -::1. ._, -= ...... >...... N 20 ...... 0 Q,> N - ..... = = Q,>= 10 .== s:i. ...=CU .1 -; 2 4 6 8 10 >..... Time (min) Figure 4.1.1 Hydrogen peroxide formation as a function of sunlight illumination time. [Fe(IIl)]o = 10 µM, [Oxalate]o = 120 µM, pH =3.96 ±0.05, lonic strength = 0.03 M (NaCl04, HCl04), air saturated solution, at 283-300 K, at solar noon on 18th Sept.1989 (lo =0.63 mE m·2s-l ). Valerophenone actinometer in deionized- tridistilled water (Co= 12 µM) was photolyzed simultaneously, in a separate photolysis tube. - 43 - In such a solution the trioxalato and dioxalato complexes are the predominant iron species (see Figure 2.3). Figure 4.1.2 shows the photo-formation of H202 with a lower concentration of Fe(III) (1 µM) and oxalate (5 µM) at pH 4. In this case, the dominant iron species are dioxalate and monoxalate. They contribute 85% and 12% of the total iron concentration, respectively. Tue initial rate of formation of H202 in September noon sunlight was 3.7 nM s-1. 1.2 1.0 ,-., :g .._,::1. 0.8 ...... N 0 0.6 N ...... == 0.4 0.2 0.0 0 2 4 6 8 10 12 Time (min) Figure 4.1.2 Hydrogen peroxide formation at a lower concentration of Fe(Jll)-oxalato complexes. The solution composition is the same as for Figure 4.J .1, except [Fe{lll)}() = 1.0 µM, [Oxalate]o = 5 µM, pH = 3.92 ±0.05. Irradiation was carried out at solar noon on 14th Sept. 1990 (lo = 0.62 mE m-2s-l ). 4.1.4 Concentration effects of Fe(III) To determine the dependence of H202 formation on the Fe(III) concentration, solutions at pH 4 containing 0 to 10 µM Fe(III), 120 µM - 44 - oxalate, and 0.03 M NaCl04 were irradiated. The fonnation of H202 in the solutions exposed to sunlight increased with increasing concentration of Fe(ffi) over the range of 0-10 µM. In the absence of Fe(ffi) no H202 was detected (see Figure 4.1.3). The results are in accordance with speciation calculations which show that the total concentration of all species of ferric oxalato complexes increases linearly with Fe(ffi) under the above conditions. 8 ~ -=-,...... 6 N 0 N =...... 4 2 2 4 6 8 10 [Fe(Ill)Jo (µM) Figure 4.1.3 Hydrogen peroxide formation in sunlight as a function of Fe (III) concentration. Samp/es exposedfor three minutes on 7th Feb. 1990 (lo = 0.39 mE m·2s-1 ). Solution composition is the same as for Figure 4.1.1, except for Fe(Ill) concentration. 4.1.5 Effects of oxalate concentration To test the oxalate concentration effects, Fe(III) was held at 10 µM and the total oxalate varied from 0 to 240 µM. - 45 - 50 40 -~ :::!.. 30 -...... 0""' 20 =...... ""' • 10 lOµMOx 0 0.0 0.1 0.2 0.3 0.4 Sunlight exposure (Einstein m-2) 0 2 4 6 8 10 Sunlight exposure time (min)* Figure 4.1.4 Photoformation of H202 in sunlight with various oxalate concentrations. The solution composition is the same as for Figure 4.1.1, except for [oxalate]o. Irradiations for [oxalate]o: 60 µM (closed circles) and 120 µM (open squares) were performed on 18th Sept.1989 (Io = 0.64 mE m-2s-l and 0.63 mE m-2s-l, respectively),for [oxalate]o: 10 µM (open circles) and 240 µM (closed squares) on 22nd Sept. 1989 (lo = 0.55 mE m·2s-l and 0.36 mE m-2s-l, respectively). *Time scale corresponds to Io = 0.63 mE m-2s-l. The results, presented in Figures 4.1.4 and 4.1.5, indicate that the accumulation of H202 in sunlight increased linearly with time and with - 46 - oxalate concentration. Tue formation of H202 can be corrected for different light intensities. In the absence of oxalate, no H202 was observed. [Oxalate]o (µM) Figure 4.1.5 Hydrogen peroxide formation as a function of oxalate concentration. (derived from Figure 4.1.4, sunlight exposure = 113 mE m-2). Cooper and DeGraff (1971) have shown in their flash photolysis studies that excess oxalate retards the decay of the transient species produced in the primary photoreaction of Fe(III)-oxalato complexes. If the precursor to H202 formation, the superoxide ion 02·-, is generated through the interaction of 02 with the transient species, the formation rate of H202 would increase with oxalate concentration. This is in agreement with our experimental results and supports the mechanism proposed above. - 47 - 4.1.6 pH effects Figure 4.1.6 demonstrates the effect of pH on the fonnation of H202. 10 m 8 m -~ ::i. '-'..... 6 ~ 0 ~ =...... 4 2 0 0 2 4 6 8 pH Figure 4.1.6 Effect of pH on the hydrogen peroxide formation. Samples exposedfor three minutes on 5th Feb. 1990 (afternoon lo = 0.22 mE m-2s-l ). The solution composition is the same as for Figure 4.1.1, except for pH. rather At the lower pH range 1.5 - 4.0, the H202 fonnation rate stayed at a H202 high and constant value. Above pH 4.0, the rate of formation of decreased with increasing pH. As discussed in the Background· Section, H02·/02·-radicals undergo disproportionation to yield hydrogen peroxide and dioxygen by reactions (6)-(7) in Figure 2.1. However, these reactions is the are strongly pH dependent. In the lower pH regime, H02· (pKa = 4.8) - 48 - dominant species. Its reaction with Fe(II) to produce H202 is 2-3 ord~rs of magnitude faster than with Fe(ID) to form 02. if comparable concentrations of Fe(II) and Fe(III) are assumed. In the higher pH regime, however, 02·- becomes the dominant species. The reaction of 02·- with Fe(II), forming H202. is much slower than that with Fe(III) leading to 02 (see Table 2.1, in which we assume comparable concentrations of Fe(III) and Fe(II)). Therefore, at the higher pH values the main product will be 02 rather than H202. In addition, above pH 5.5, the fraction of Fe(III)-oxalato complexes decreases sharply, which further reduces the possibility of H202 formation. 4.1.7 Stoichiometry of H202 formed to oxalate decomposed Figure 4.1.7 shows the photodecomposition of oxalate in the presence of 10 µM Fe(III), and 120 or 240 µM oxalate at pH 4. Under these conditions, the photodegradation of oxalate bad a half-life of a few minutes in September noon sunlight. Comparing the corresponding slopes in Figures 4.1.4 and 4.1.7 gives the quantitative relationship between the amount of H202 formed and the amount of oxalate consumed, .1.H20z/.1.oxalate = 0.45 ± 0.05. This value decreases with increasing pH and with increasing iron concentrations. At lower pH and lower iron concentration, the ratio approaches the theoretical limit, 1 (data are not shown). Assuming a mechanism consisting of reactions (1)-(3) and (5)-(8) in Figure 2.1 with a Fenton reaction occurring simultaneously, the values of the mole ratio of hydrogen peroxide formed to oxalate consumed should lie between 0 and 1. The qualitative agreement between these calculated values and the experimental results supports the proposed mechanism. - 49 - 1.0 ,_.,=> ~ ....= 0.8 -=~ 0...... 0.6 ,_., -....~ = 0.4 -=~ 0...... 0.2 0.0 0.0 0.1 0.2 0.3 0.4 Sunlight exposure (Einstein m-2) 0 2 4 6 8 10 Sunlight exposure time (min)* Figure 4.1.7 Decomposition of oxalic acid. The solution composition is the same as for Figure 4.1.1, except for oxafote concentration. lrradiations of solutions containing 120 µM oxalate (open aquares) and 240 µM oxalate (closed squares) were performed on 18th Sept. 1989 ), (lo = 0.64 mE m-2s-l) and on 22nd Sept. 1989 (lo = 0.36 mE m-2s·l respectively. *Time scale corresponds to Io = 0.63 mE m·2s-l. 4.1.8 Stability of H202 in solution The results in Table 4.1.1 show that the H202 formed is quite stable well after completion of irradiation. This is not surprising, although it is - 50 - known that both Fe2+ and fe3+ are capable of catalyzing the decomposition of hydrogen peroxide. Table 4.1.1 Stability of photochemically formed hydrogen peroxide*. expt. no. irradiation K2C204 H202 found at following times time (min) (I0-5 M) after completion of irradiation 10 min. lh. lh/lOmin (µM) (%) 1 2 6 4.9 5.0 102 2 4 6 9.6 9.6 100 3 9 6 14.7 14.3 97 4 0.5 12 2.5 2.6 104 5 2 12 9.5 9.2 97 6 4 12 19.9 19.6 98 7 9 12 36 35 96 *Except where noted, all irradiations were performed at solar noon ( maximum solar altitude ). The experiment 1, 2, 3, and 4, 5, 6, 7 were performed on l 8th Sept. 1989, the sunlight intensity Io = 0.64 mE m-2s-1 and lo = 0.63 mE m-2s-l, respectively. The photolytic solution composition: [Fe(IIJ)]o = 10.0 µM, pH = 4.00 ± 0.05, ionic strength 0.03 M (NaC/04, HC/04), [Oxalate]o = 30.0, 120 µM, air saturated, at 283-300 K. Tue catalytic process involving fe2+ ion can be described by reactions (4.2) and (4.3): - 51 - (4.2) Fe2+ + HO· ~ Fe3+ + HO- (4.3) In the presence of Fe3+, an additional reaction is involved according to the Ha:ber-Weiss mechanism (Walling, 1975): (4.4) However, the ferric ion reacts much slower with H102 than the ferrous ion (see Table 2.1). In addition, catalysis by Fe3+ is sensitive to changes in the coordination environment of the ferric ion (Jones et al., 1959). Tue formation of the Fe(IlI)-oxalato complex inhibits the decomposition of hydrogen peroxide (Weiss, 1947). In the Fe2+-H202-oxalate system, the decomposition of hydrogen peroxide was virtually instantaneous; before an initial sample could be taken, the hydrogen peroxide bad been decomposed, and Fe2+ simultaneously oxidized to Fe(III) (lt should be noted that all values of the H1ü2 concentration measured in this study are the values remaining after this reaction). Subsequently, H102 was stabilized by the residual oxalate via complex formation. To confirm this point, the dark experiments were carried out with initial concentrations of 5.8 µM H102, 30 µM oxalate and 10 µM Fe2+ or Fe3+ in the solutions saturated with air at pH 4. Tue results are shown in Figure 4.1.8. - 52 - 8 7 ..-. ::; :j. 6 '-' 'N 0 5 N ~ 4 3 10 20 30 40 50 60 70 Time (min) Figure 4.1.8 Inhibition of hydrogen peroxide decomposition and stabilization of hydrogen peroxide by oxalate ( dark experiments). [H202Jo = 5.8 µM, 0.03 M NaC/04, 30 µM oxalate at pH 4, air saturated (02 0.21 atm). Open squares: 10 µM Fe(III), closed squares: 10 µM Fe(II). 4.1.9 Photochemical production of H20 2 from the Fe(II)-oxalic acid-02 system Figure 4.1.9 shows the typical concentration profiles for the sunlight induced formation of H202 from the Fe(II)-oxalic acid-02 system. A striking feature is the presence of an induction period in the formation of H202 and depletion of oxalic acid, whereas none was observed in the reactions with the Fe(III)-oxalic acid-02 system. During this induction period ferric ions were generated. These facts indicate that the Fe(II)/Fe(III) - 53 - couple is involved in the photochemical conversion of oxalic acid into H202, with Fe(Ill) being the active species. Spectroscopic evidence reveals that the iron is predominantly in the divalent state under sunlight irradiation. 40 1.2 1.0 ,..._ 30 0.8 ~ :i 0 '-' ,...... , 20 0.6 u 0""' --u 0.4 ...... ""' = 10 0.2 0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 Sunlight exposure (Einstein m·2) Figure 4.1.9 Photoproduction of H202 from Fe(Il)~oxalic acid-02 system. [Fe(Il)]O = 10.0 µM, [Ox]o = 120 µM, ionic strength = 0.03 M (NaCI04, HC/04), pH = 4.0 ±0.05.1rradiation was performed at solar noon on 27th May 1991(10 = 0.99 mE m·2s-1 ). 4.1.10 Formation of H202 by photolysis of other carboxylic acid -Fe(IIl)-02 systems To explore the formation of H202 from other organic ligand-Fe(ffi) complexes by the proposed mechanisms, some photochemical experiments were carried out with the Fe(ffi)-glyoxalic acid-02 and Fe(III)-pyruvic acid- 02 systems. Both glyoxalic and pyruvic acid are found in atmospheric liquids in a concentration range similar to oxalic acid. They may also form complexes with ferric ion. The experiments were carried out with [Fe(III)]o - 54 - = 10 µM, [organic ligand]o = 60 µM, ionic strength = 0.03 M (NaCl04, HCI04) at pH = 3.75 ± 0.05 . ...... N 0 N =...... Sunlight exposure (Einstein m-2) Figure 4.1.10 Formation of H202 by photolysis of other carboxylic acid-Fe(lll)-02 system. [Fe(Ill)]o = 10.0 µM, ionic strength = 0.03 M (NaC/04, HC/04), pH = 4.0 ±0.05, [glyoxalic acid]o = 60 µM (open squares), [pyruvic acid]o = 60 µM (closed diamonds), Irradiation was performed on 24th Oct. 1991 (lo = 0.44 mE m-2s-lfor glyoxalic acid and 0.21 mE m-2s-l for pyruvic acid). lt is seen from Figure 4.1.10 that with pyruvic acid the rate of H202 formation is of a similar order of magnitude as that of oxalic acid, and with glyoxalic acid it rises by a factor of 2. With pyruvic acid, a stoichiometric ratio of 0.46 ± 0.05 for H202 formed to pyruvic acid depleted was obtained. This ratio is similar to that with oxalic acid. With glyoxalic acid a ratio of 1.0 ± 0.05 for H202 formed to glyoxalic acid depleted was obtained. - 55 - 4.1.11 Subsequent OH radical formation Preliminary experiments were carried out with 1-chloro-butane as a reference compound for OH radicals, in the presence of an excess of octanol to control the lifetime of the OH radical (Haag and Hoigne, 1985). The results show that irradiating aqueous solutions containing Fe(III)-oxalato complexes at pH 4 produces OH radicals; This production is due to the preceding formation of Fe(Il) and H102. The addition of catalase inhibits the OH radical formation. - 56 - 4.2 Photochemical/chemical cycling of Fe(llI) and Fe(II) complexes in atmospheric waters 4.2.l Kinetics for monochromatic photoreduction of Fe(III) in deoxygenated solutions When the experimental results are plotted in the form log [Fe(III)]/(Fe(IIl)]o against time, a linear relationship is obtained (Figure 4.2.1). This is in agreement with the theory of photochemical kinetics. 0.0 ...... ~ .....,...._ -0.2 ..... -(,) ;::::.:=. -0.4 ,...._...... (,) -0.6 -...... ~ ~ ~ -0.8 -1.0 0 2 4 6 8 10 Time (min) Figure 4.2.1 Photoreduction of Fe(IJJ) with 313 nm monochromatic light in deaerated solutions. [Fe(Il/)]o = 10.0 µM, [oxalate]o =30 µM, pH = 4.00 ± 0.05, ionic strength = 0.03 M (NaC/04, HC/04), deaerated using N2, at 293 K, /313 nm = 0.60 µE /-ls-1 (valerophenone actinometer ). - 57 - The rate for photoreduction of Fe(III) at a fixed wavelength A. is expressed by : -(d[Fe(III)]/dt h„ =cl>A Im„ (1-lQ-E).CD)(SN) (4.5) where d[Fe(IIl)]/dt is the average rate for disappearance of Fe(III) (M/min), where k is the first-order photoreaction rate constant. In this study, the photolyses of Fe(ill)-oxalato complexes were carried out under conditions of such low absorbances. Hence, the disappearance rate of Fe(III) in the deoxygenated solutions should be first order with respect to concentration of Fe(ill)-oxalato cornplexes [Fe(III)( Ox )n]. Figure 4.2.2 presents the fonnation of ferrous ion and the depletion of oxalic acid as a function of time under 313 nm monochromatic radiation at pH4. - 58 - 31 10 30 8 ,-., :; 29 ,-., ::i. :; .._, 6 ::i. 28 .._, ...... ,-., ..... ~ '-' 4 c:.I 27 Q. r...... 2 26 25 2 4 6 8 10 Time (min) Figure 4.2.2 Photoproduction of Fe(ll) and depletion of oxalic acid with 313 nm Light in deaerated solutions. The data are derived from the same experiments as those for Figure 4.2.1. 4.2.2 Effects of pH and concentration of oxalic acid on monochromatic photoreduction The photolysis of Fe(IIl)-oxalato complexes by 313 nm monochromatic radiation at various values of pH are illustrated in Figure 4.2.3. At each pH, the plots of In [Fe(ill)]/[Fe(III)]o versus exposure time are linear. The reaction rate constants were derived from the slopes of these lines and are listed in Table 4.2.1. - 59 - -0.5 -1.0 -= -1.5 -2.0 ,__..__...__....__.___.__....__.___....__....._...._ ...... _._~ 0 1 2 3 4 5 6 7 Time (min) Figure 4.2.3 Photoreduction of Fe(Ill)-oxalato complexes at various pH values with 313 nm irradiation. The irradiation conditions and solution composition are the same as for Figure 4.2.1, except for the pH of the solutions. Open squares: pH 1; triangles: pH 2; open circles: pH 3; closed circles: pH 4. Although the predominant species of Fe(IIl)-oxalato complexes shifts from monooxalato at pH 1 to dioxalato complexes at pH 4, the photoreductive rate is very similar. - 60 - 0 c= -...... '-' -Q.I -1 r;i;. -'-"' ---...... '-' -Q.I -~ -2 -= -3'--_._~...___.~_._~...__._~_.___.~_....__. 0 2 4 6 8 10 Time (min) Figure 4.2.4 Photoreducti'on of Fe(///) with various oxalate concentrations under 313 nm light. The irradiation conditions and solution composition are the same asfor Figure 4.2.1, except for the concentrations of oxalic acid. Squares: 30 µM; open circles: 60 µM; closed circles: 240 µM oxalic acid. Figure 4.2.4 depicts the photoreduction of Fe(III) at concentration of oxalate of 30, 60 and 240 µM. Here the photochemically reactive species change from a di-oxalato complex dominant at the lower concentration (30 µM) to a tri-oxalato complex, which dominates at the higher concentration of oxalate (240 µM). The small reduction in reaction rate constants with higher concentrations of oxalate is not significant. These results demonstrate that various oxalato complexes of Fe(III) have similar photoreactivity. The high photochemical reactivity of the Fe(IIl)-monooxalato species is in - 61 - contrast to the zero quantum yield for this species obtained by Vincze and Sapp (1987). Tab/e 4.2.1 Photolysis rate constants of Fe(l/1)-oxalato complexes at various values of pH under 313 nm monochromatic light* pHvalue [Fe(III)]o (µM) [OX]o(µM) k (s-1) 1.1 10 30 4.5 X 10-3 2.1 10 30 4.8 X 10-3 3.0 10 30 4.6 X 10-3 4.0 10 30 4.6 X 1Q-3 4.0 10 60 4.5 X 1Q-3 4.0 10 240 4.4 X J0-3 *The photoreductive rate constants are derivedfrom the data in Figures 4.2.3 and 4.2.4. Errors for the given va/ues of k are in the range of ± 0.3 x J0-3. 4.2.3 Influence of dissolved dioxygen on monochromatic photoreduction To test the influence of molecular oxygen on the photoreduction of Fe(III)-oxalato complexes, air-saturated solutions were used. Aliquots of irradiated solution were pipetted within one second into a brown volumetric flask, containing 1,10-phenanthroline, NH4F and acetate-acetic acid buffer solution for analysis of Fe(II). Typical results are given in Figure 4.2.5. - 62 - 8 m 6 m -~ m :::!.. -...... lo-4 4 lo-4 -~ -i;;,;...... 2 011-__,___._____.~.___,__.____._~.__...____.___._____.~.___. 0 5 10 15 20 25 30 35 Time (min) Figure 4.2.5 Influence of dissolved oxygen on the photoreduction of Fe(//1)-oxalato complexes with 313 nm light. The irradiation conditions and solution composition are the same as for Figure 4.2.1, except for air-saturation. The rate for production of Fe(II) decreased compared with that in the deaerated solution. In a few minutes, a photostationary state for the reduction and reoxidation of Fe(III) and Fe(II) was established. When the irradiated solutions were put in the dark, Fe(II) was rapidly reoxidized to Fe(III). The previous section of this work has shown that hydrogen peroxide was formed by sunlight irradiations of such air-saturated ferrioxalate solutions. With 313 nm radiation, H202 was also produced (see Figure 4.2.6). The H202 is responsible for the dark reoxidation of Fe(II). - 63 - 10 8 '1- ._,:::1. 6 ....., N 0 N 4 =...... 2 5 10 15 20 25 30 Time (min) Figure 4.2.6 Formation of H202 by photo/ysis of Fe(lll)-oxalato comp/exes in air-saturated so/ution with 313 nm light. The irradiation conditions and solution composition are the same as for Figure 4.2.1, except for air-saturation. 4.2.4 Kinetics of sunlight photolysis in air-free solutions For sunlight photolyses of dilute Fe(ID)-oxalate solution, the reaction quantum yield is nearly independent of wavelength in the range of 285 < A. < 440 nm. Above 440 nm the absorbance of Fe(ID)-oxalato complexes is very weak. The photoreaction in this region is not significant compared to the photolyses at shorter wavelengths. Thus, the complete rate equation for the sunlight photolyses ~an be also written in the fonn of first-order kinetics -(d[Fe(ID}]/dt)s =ks [Fe(ID)] (4.7) where - 64 - (4.8) ks is the first-order rate constant for sunlight photolysis and represents the sum of the kA. values over all wavelengths of light absorbed by the Fe(III)- oxalato complexes. The representative results are plotted in Figure 4.2.7a. The linear relationship between ln [Fe(IIl)]/[Fe(IIl)]o and illumination time confirms that the photolysis is first order with respect to the concentration of Fe(IIl)-oxalato complexes. The production of Fe(II) and decomposition of oxalic acid versus sunlight exposure time is shown in Figure 4.2.7b. 31 0 10 ...... = to-C b 30 to-C -to-C 8 '-' ~ ::a -1 -::i. 29 ::a ~ '-' -::i. 6 '-' ';:::::, ...... 28 ...... to-C Q ~ to-C to-C -to-C '-' 4 0 '-' ~ .__. ~ -2 27 .__.r... r....__. 2 26 = - 25 20 40 60 80 50 100 150 Sunlight exposure time (sec) Sunlight exposure time (sec) Figure 4.2.7 Sunlight photoreduction of Fe(J/1)-oxalato complexes in deoxygenated solutions. a) First-order treatment of the photolysis data, b) production of Fe(Il) and depletion of oxalic acid. The solution composition is the same as for Figure 4.2.1. Irradiation was performed at solar noon on 13th Aug. 1989 (Io = 0.70 mE m-2s-1 ). - 65 - 4.2.5 Effects of pH and the concentration of oxalate on the sunlight-induced reduction of ferric ion Photolysis of Fe(III)-oxalato complexes under sunlight at various values of pH is presented in Figure 4.2.8, and at various concentrations of oxalate in Figure 4.2.9. In each case, a straight line was obtained when ln[Fe(III)]/[Fe(III)]o was plotted against irradiation time. 0.0 ...... -= -.... -0.5 ....'-' Figure 4.2.8 Sunlight induced photoreduction of Fe(IIl)-oxalato complexes at various pH values. The solution composition is the same as for Figure 4.2.1, except for the pH values. Irradiations for pH 1.1 (open squares) and 2.1 (closed squares) were peiformed at solar noon on 16thAug.1989, for pH 4.1 (open circles) and 5.0 (closed circles) at solar noon on 13 Aug. 1989 (lo = 0.70 ±0.10 mE m-2s-l ). - 66 - 0 ,...... ~ ..... -..... '-' Cl> -1 ...... r.. ,...... -..... -..... '-' Cl> -2 ...... r.. -= 20 40 60 80 Time (sec) Figure 4.2.9 Sunlight induced photoreduction of Fe(Jll) with various concentrations of oxalic acid. The solution composition is the same as for Figure 4.2 .1, except for the concentration of oxalic acid. Irradiations for oxalate concentration of 30 µM (open squares) were performed at solar noon on 13th Aug. 1989, for oxalate concentrations of 60 (open eire/es) and 120 µM (closed squares) at solar noon on 31Aug.1989, and 240 µM (closed eire/es) at solar noon on llth Sept.1989 (Io = 0.71 ±0.07 mE m-2s-1). The photoreduction rate constants are collected in Table 4.2.2. These results are in good agreement with the results obtained by monochromatic photolysis. - 67 - Table 4.2.2. Photoreduction rate constants of Fe(lll)-oxalato complexes at various values ofpH under sunlight* pHvalue [Fe(III)]o (µM) [OX]o(µM) k (s-1) 1.1 10 30 3.0 X 10-2 2.1 10 30 4.4 X lQ-2 4.1 10 30 4.0 X lQ-2 5.0 10 30 3.4 X lQ-2 4.0 10 60 3.9 X lQ-2 4.0 10 120 4.0 X lQ-2 4.0 10 240 3.2 X lQ-2 *The photoreductive rate constants are derived from the data in Figures 4.2.8 and 4.2.9. 4.2.6 Sunlight photoreduction of Fe(III) in the presence of oxygen Figure 4.2.10 illustrated the photochemical reduction of ferrioxalate in air-saturated solutions at pH 4. Tue concentration of Fe(II) rapidly reach the steady state. When a solution containing 10 µM Fe(ill) and 30 µM oxalic acid at pH 4 was exposed to sunlight at solar noon in August, about 70% of Fe(ill) was converted to Fe(II) at steady state. When the photolytic tubes were removed from the sunlight and placed in the dark for a few seconds, no or only a small amount of Fe(II) remained, depending on the concentration of hydrogen peroxide formed (see Section 4.1). - 68 - 10 8 -~ ::t -...... 6 .-c -.-c ~ -i;;i;...... 4 2 50 100 150 200 250 300 350 Time (sec) Figure 4.2.10 Sunlight induced photoreduction of Fe(IIl)- oxalato complexes in the presence of oxygen. The solution composition is the same as for Figure 4 .2 .1. except for air-saturation. Irradiation was peiformed at solar noon on 15th Aug. 1989 (lo = 0.74 mE m-2s-l ). 4.2. 7 Solar photolysis of Fe(Ill) associated with formic and acetic acid Some photochemical experiments were carried out with the air- saturated systems Fe(III)-oxalate-monocarboxylic acid, Fe(III)-mono- carboxylic acid, and Fe(Cl04)3 alone. For these experiments 100 µM formic and acetic acid were used. In the first system, no appreciable influence of monocarboxylic acids was found on the reduction of Fe(III) and formation of hydrogen peroxide. In the absence of oxalate, the photoproduction of Fe(II) from Fe(III) with and without monocarboxylic acid is shown in Figure 4.2.11. - 69 - 6 I I 5 Acetic acid ·! ~ ::1. I 4 ...... Formic acid ~ ~ '-'- ~ 3 ...... ~ 2 1 0 0 5 10 15 20 Time (min) Figure 4.2.11 Sunlight induced photoproduc~ion of Fe(Il) from Fe(Ill)-formic acid-02 and Fe(lll)-acetic acid-02 systems. Open squares: [Fe(Ill)Jo = 10 µM, 100 µMformic acid, ionic strength = 0.03 M (NaCI04, HCI04), air-saturated, pH = 3.9; irradiation was peiformed in afternoon on 15th Oct. 1990; diamonds: the solution composition is the same asfor the Fe(lll)jormic acid-02 system but with 100 µM acetic acid instead of formic acid at pH 4.0; irradiation was peiformed in the afternoon on 14 Sept. 1990. For both irradiations Io = 0.28 mE m-2s-1. lt was observed that both formic and acetic acids accelerate the photoreduction of Fe(III). lt was demonstrated by a spectrophotometric examination that complex formation between ferric ion and formic and acetic acids is negligible in the concentration range in which these experiments were carried out. Therefore, the photochemically active species - 70 - must be the Fe(OH)2- complex. The acceleration of photoproduction of Fe(II) could be explained by the interactions of carboxylic acids with the excited product of Fe(OH)2- and with the OH radicals produced in the photolysis of Fe(OH)2- (Evans et al., 1951; Carey and Langford, 1975): Fe2+0H* + RCOOH ~ Fe2+ + H20 + RCOO· (4.9) OH· + RCOOH ~ Off + RCOO· (4.10) These reactions suppress the back reaction of the excited complex to the unexcited reactant and the reaction between OH· and Fe(II) ion. 4.2.8 Photochemical/chemical cycle of ferric and ferrous complexes The overall stoichiometry for the photolysis of ferrioxalate in deaerated solution is expressed by equation 4.1. lt indicates that for every ferrioxalate decomposed, there should be a corresponding gain of one Fe(II). Figure 4.2.2 depicts the photo-production of Fe2+ with monochromatic light (313 nm) in the MGRR. The quantum yield determined from the initial slope is 1.2 ± 0.1 in good agreement with literature values (Hatchard and Parker, 1956). In the presence of 02 (0.21 atm.), we observed that 70% of the total iron at steady-state became Fe(II) during solar irradiation of a solution containing 10 µM iron ion, 30 µM oxalic acid at pH 4. However, when the solution was placed in the dark for a few seconds after irradiating for a few minutes, only small amount of Fe(II) remained. This observation could be significant, because of possible photochemical/chemical cycling between ferric and ferrous complexes if ferric oxalato complexes, light, and 02 are present. - 71 - As illustrated in Figure 5.1, this cycling includes; (i) the direct photolysis of Fe(IIl)-oxalato complexes; (ii) the reaction of the organic radical photoproducts with oxygen leading to H02·/02·· radical formation, which leads to H102 production and (iii) the subsequent reaction of these photooxidants with Fe2+ (aq), leading to the reformation of Fe(III) and Fe(IIl)-oxalato complexes. The direct oxidations of photo-generated Fe2+(aq) by 02 and by H102 have not been considered in Figure 5.1. The reaction with 02 is very slow at pH < 6. In the absence of complexing ligand and light, the half-life for Fe(II) oxidation by 02 at pH 4 is longer than 285 days (Stumm and Morgen, 1981). Although oxalate accelerates the oxidation of Fe(II), we did not observe any loss of Fe(m after keeping a solution of 10 µM Fe(II), and 30 µM oxalate in 0.03 M NaCl04 at pH 4 in the dark for 24 hr. Even in the presence of 240 µM oxalate, less than 2% of the Fe(II) was oxidized after 1 hour in the dark. This indicates that the direct oxidation of Fe(II) by 02 is insignificant under the conditions used. The reaction of H102 with Fe(m is also relatively slow when compared to the oxidation of Fe(II) by H02· and Or radicals. But in the dark, H102 is the dominant oxidant for Fe(II) because the concentrations of other photooxidants such as hydroperoxyl and superoxide radical are very low. -72 - 4.3 The Photolysis of oxalic, glyoxalic and pyruvic acids catalyzed by iron in cloud· and fog-waters 4.3.1 Kinetics of the disappearance of oxalic acid under monochromatic and solar irradiation Tue kinetic data for the photolysis of Fe(ill)-oxalato complexes in the absence of dioxygen have been given in section 4.2. Tue production of Fe(II) is first order with respect to the concentrations of ferrioxalate, and the stoichiometric ratio of Fe(II) formed to oxalic acid depleted is 2 : 1. Figures 4.3.1 shows the photodepletion of oxalic acid in air-saturated solutions with 313 nm irradiation, and Figure 4.3.2 with sunlight. 0.8 c ~ 0 ...... 0.6 --~ 9. 0.4 0.2 0 10 20 30 40 Time (min) Figure 4.3.1 Photodecomposition of oxalic acid in the presence of ferric ion and oxygen under 313 nm light. The irradiation conditions and the solution composition is the same as for Figure 4.2.1, except for the concentrations of oxalic acid. Closed circles: 30 µM oxalic acid; open circles: 60 µM oxalic acid. - 73 - 1.0 0.8 ,...... Q > 0.2 0.0 0 2 4 6 8 10 Time (min) Figure 4.3.2 Photodecomposition of oxalic acid in the presence offerric ion and oxygen under sunlight. The solution composition is the same as for Figure 4.2.1, except for the concentrations of oxalic acid (Ox). Closed eire/es: 60 µM Ox; open squares: 120 µM Ox. Irradiations were peiformed at solar noon on 16-1 Bth Sept. 1989 (lo = 0.64 ± 0.01 mE m·2s-l) The reaction is no longer pseudo-first-order but closer to apparent zero-order kinetics. This result is consistent with the previous results of photoreduction of ferrioxalate in air-saturated solution, where the steady- state for photoredox was rapidly established and then the concentration of Fe(III) complex remained constant. 4.3.2 Influence of the initial concentration of Fe(Ill) Two series of experiments were carried out in order to test the influence of the amount of iron on the oxidation of oxalic acid. In the first - 74 - series of experiments, 5 µM oxalate was used and the initial Fe(Ill) concentration was varied from 0 to 10 µM. Tue reaction rate, rÄOx• for the disappearance of oxalic acid against the initial concentration of Fe(III) is presented in Figure 4.3.3. 6 8 10 [Fe(III)] o (µM) Figure 4.3.3 Rate of decomposition of oxalic acid in sunlight as afunction ofthe initial concentration of Fe(III). [Ox]o = 5 µM, ionic strength = 0.03 M (NaCl04, HCl04), pH = 4.00 ±0.05, air-saturated, at 283-300 K. Exposure time: two minutes at solar noon on 22nd April 1991 (Io = 0.78 mE m·2s-1) Tue rate rÄOx increases linearly with the concentration of Fe(III) over the range 0 to 5 µM. Thereafter it remains constant. Tue maximum degradation rate corresponds to complete complexation of oxalic acid. In the second series of experiments, a higher concentration of oxalic acid, 120 µM, was - 75 - employed. In this case, rßOx increases with increasing amount of Fe(III) up to a level which corresponds to complete absorption of the incident light potentially absorbable by ferrioxalate. Beyond this concentration, rßox is independent of the initial Fe(III) concentration (Figure 4.3.4). 25 (,,1 20 -Q,I ~ ~ = 15 -~ 0 5 [Fe(Ill)]o (µM) Figure 4.3.4 Effect of the initial concentration of Ferric ion on the rate of photodegradation of oxalic acid with 313 nm light. [Ox]o = 120 µM, ionic strength = 0.03 M (NaCl04, HCl04), pH = 4.20 ±0.05, air-saturated, at 293 K. Exposure time: 7 min in MGRR !313 nm = 1.60 µE [-1 s·l (valerophenone actinometer ). 4.3.3 lnfluence of initial concentration of oxalic acid Tue complex formation constants of oxalic acid and ferric ion are !arge and the ligand substitution process in this system is very fast. Tue photochemical reaction is the rate-limiting step. Therefore, the overall rate of photolytic reaction must be proportional to the concentration of the - 76 - Fe(ill)-oxalato complexes Fe(ill)(OX)n. This dependence on oxalic acid is shown in Figure 4.3.5. 350 300 ,__ (,,1 c:.I 250 "' --::E 200 ,_.c ;.( 150 0 ~ .... 100 50 [Ox]o (µM) Figure 4.3.S Rate of decomposition of oxalic acid with sunlight as a fanction of the initial concentration of oxalic acid. [Fe(IIl)]o = 10 µM, ionic strength = 0.03 M (NaC/04, HC/04), pH = 4.00 ± 0.05, air- saturated, at 283-300 K. Exposure time: five minutes at solar noon on 28th March 1991 (Io = 0.70 mE m·2s-1) Tue rate of disappearance of oxalic acid varies linearly with the oxalate concentration at low concentrations (first-order kinetics), and becomes independent of oxalate concentration (zero-order kinetics) at high concentrations where complete complexation of Fe(III) occurs. - 77 - 4.3.4 Influence of pH The rate of the photodegradation of oxalic acid is strongly pH dependent. When the pH is varied from 1 to 6.4, the rate of degradation passes through a maximum value at about pH 3.0 (see Figure 4.3.6). 200 ,-, CJ 150 C.I -.!!: ~ .._, = 100 ~ 0 pH Figure 4.3.6 lnfluence of pH on the degradation of oxalic acid under sunlight.[Fe(/II)]o = 10 µM, ionic strength = 0.03 M (NaC/04, HCZ04), [Ox]o = 120 µM, air-saturated, at 283-300 K. Exposure time: five minutes at solar noon on 14th March 1991 (lo = 0.48 mE m·2s-l ). This pH effect could be due to the fact that the pH influences the speciation of oxalato complexes and oxalate ions. In nitrogen atmosphere, experimental results have indicated that the variations of Fe(III) oxalato species only slightly affects the photoproduction of Fe(II) and are not sufficient to explain the strong pH influence in the oxygenated solutions. Figure 4.3.7 presents speciation distribution of uncomplexed oxalic acid as a function of pH. - 78 - 1.0 HOx- Qx2- 0.8 ·o"'C.I C.I Q. "' 0.6 ....0 .!:= 0.4 (j -ca ~""' 0.2 0.0 1 2 3 4 5 6 7 pH Figure 4.3.7 Fraction of uncomplexed oxalic acid and oxalate ion in Fe(IIJ)-oxalate system ([Fe(III)] = 10.0 µM, [Ox] = 120 µM, 0.03 M NaCl04, 298 K). lt is worth noting that the HC2 0 4 - ion also achieves a max im um concentration near pH 3. lts concentration decreases sharply at higher and lower pH values. This pH function is very similar to that observed for the depletion of oxalic acid. lt could imply that some photochemically produced intermediates react with HC204- ions much more easily than with oxalic acid and other oxalate ions. In addition, the change of pH can affect protonation states of the intermediates formed. For example, C02H radical has an acid dissociation constant (pKa) of 1.4. Rowan et al. (1974) have estimated the pKa of HC204 to lie somewhere between the pKa of HC204- (4.1) and pKa of C02H radical. The protonated radicals exhibit a redox ability which is very different from that of their deprotonated partners. This could contribute to the pH effect observed in Figure 4.3.6, although the axact acid-base and redox properties of the intermediates involved are not known. - 79 - 4.3.5 Photodecomposition at more dilute concentrations of Fe(llI) and oxalic acid In atmospheric liquids of cleaner areas the concentrations of iron and organic ligands are usually lower. To examine the kinetics for the decomposition of oxalic acid under these cleaner conditions, the concentration of oxalic acid was held at 10 µM and the initial Fe(III) concentration was varied from 0.4 to 1 µM. Typical results are shown in Figure 5.3.8. 10 10 8 8 -~ ::1.. ~ 6 -::1.. 6 -~ 0"" -~ 4 ::r.: 0...... 4 ...... "" 2 2 0 0 0 2 4 6 8 10 12 14 Time (min) Figure 4.3.8 Photodepletion of oxalic acid and formation of 11202 at more dilute concentrations of oxalic acid and ferric ion under sunlight. For all solutions [Ox]o = 10.0 µM, ionic strength = 0.03 M (NaC/04, HC/04), air- saturated, pH = 4.00 ± 0.05, at 283-300 K. [Fe(II/)]O = 0.4 µM (squares), 0.8 µM (triangles) and 1.0 µM (circles). Irradiations were performed at solar noon in Aug. 1991 (lo = 1.1 ± 0.1 mE m·2s-l ). - 80 - As at higher concentrations of oxalic acid, the rates of degradation of oxalic acid are higher at the beginning and gradually decrease to constant rates in a few minutes. Tue increase in the amount of Fe(III) corresponds to a faster decomposition. Figure 4.3.8 also gives the concentration of H202 as a function of the time. A similar kinetic tendency was observed with H202. 4.3.6 Influence of sunlight intensity In oxygen-free solutions the photolysis of ferrioxalate is proportional to light intensity. This is also the case if air-saturated solutions are exposed to low light intensities. Representative results are given in Figure 4.3.9. Time (min) Figure 4.3.9 Effect of sunlight intensity on photodegradation of oxalic acid. [Ox]o = 10.0 µM, [Fe(IIl)Jo = 1.0 µM, ionic strength = 0.03 M (NaCl04, HCl04), air-saturated, pH=4.00 ± 0.05, at 283-300 K. Irradiations were performed on 6th Aug. 1991 (open circles: lo = 0.43 mE m·2s-1, closed circles: lo = 0.90 mE m-2s-1 ). - 81 - When a high intensity of irradiation was applied, this correlation did not exist any more. This might be explained by the fact that the acceleration of photoreduction shifts the steady-state of photoredox to Fe(II) and reduces the steady-state concentration of Fe(ID). 4.3. 7 Photodegradation of glyoxalic and pyruvic acids in the presence of iron Tue sunlight-induced decomposition of glyoxalic and pyruvic acids in the presence of Fe(ill) has shown a kinetic characteristics similar to that for the depletion of oxalic acid. Tue disappearance of substrates is a linear function of exposure time (Figure 4.3.10). 1.0 0.9 0.8 u= --u 0.7 0.6 0.5 0.0 0.1 0.2 0.3 Sunlight exposure (Einstein m-2) Figure 4.3.10 Sunlight-induced photodegradation of glyoxalic and pyruvic acids in the presence of iron. [Fe(I/I)Jo = 10.0 µM, ionic strength = 0.03 M (NaC/04, HC/04), pH = 4.0 ±0.05, [glyoxalic acid]o = 60 µM (open squares), [pyruvic acid]o = 60 µM (closed diamonds). Irradiation was performed on 24th Oct. 1991 (Io = 0.44 mE m·2s-lfor glyoxalic acid and 0.21 mE m·2s-1 for pyruvic acid). - 82 - 4.4 Photochemical formation of H202, depletion of oxalic acid and oxidation of dissolved S02 in real fogwaters 4.4.1 Fog water concentrations and the absorption spectrum In order to relate the model system to tropospheric water environments, fog waters were collected in Dübendorf. After collection, fog waters were stored at 4 ·c. Tue samples were filtered at 0.45 µm and analyzed immediately before irradiation. Tue concentrations of relevant constituents are given in Table 4.4.1. Table 4.4.1 The concentrations of relevant constituents in fogwater and liquid water content (LWC) Fe(III) 1.7 µM Sulfite 79µM Fe(II) 14 µM sulfate 86µM H2()z l.2µM formaldehyde 59µM DOC 12 mg/l pH 3.7 oxalate 4.8µM LWC 2.6 X 10-3 )/m3 Figure 4.4.1 shows the absorption spectrum of fogwater that was used for irradiation. Tue absorption gradually decreases with increasing wavelength with essentially no absorption above 600 nm. Tue absorption could be caused by organic chromophores and transition meta! complexes. - 83 - 1.0 0.8 ~ ~ =CO 0.6 .l:l Q ""'{ll .l:l 0.4 < 0.2 0.0 240 340 440 540 640 740 Wavelength (nm) Figure 4.4.1 Absorbance spectrum (pathlength of 10 cm) of fog water used for irradiation. 4.4.2 Generation of hydrogen peroxide, oxidation of dissolved sulfur dioxide and degradation of oxalic acid in fog water by irradiation with 313 nm light Irradiations were carried out in separate quartz photolysis tubes containing 20 ml of fogwater sample solution saturated with air, inside a water-cooled merry-go-round-reactor (MGRR), using 313 nm line of the 500 W Hg lamp. Fig.4.4.2 summarizes the accumulated concentration of H202 as a function of irradiation time. The concentration is the net accumulated concentration of H202 formed, that is, the sum of the production and decomposition reactions maintaining the solution concentration of H202. The concentration increases with irradiation time. Oxalic acid is depleted simultaneously. - 84 - 2.5 5 4 ,-... ,-.. ~ :i ~ 2.0 3 :i ---...... ~ ---N .... 0 2 = N -II<= =...... 1.5 0...... 1.0 .___._ _ _.___.._...... ___. _ _.__~_.__...... ___._ 0 0 30 60 90 120 150 Time (min) Figure 4.4.2 Photochemica/ formation of H202 and decomposition of oxa/ic acid infog water. Irradiation was performed with 313 nm light in MGRR Io = 1.2 µE /-ls-l(va/erophenone actinometer) As listed in Table 4.4.1, the fogwater sample contains sulfite. Sulfite is oxidized in the presence of H202. Fig.4.4.3 shows the disappearance of sulfite during the irradiation. The amount of sulfite oxidized increases with exposure time. Experimental and theoretical evidence currently indicates that several oxidants are important in the conversion of dissolved sulfur dioxide (S02) to sulfuric acid in the aqueous phase (i.e., cloud-, fog- and rain-drops). The oxidation of S02 can occur through several pathways. If the pH of water is greater than 5, ozone (03) is probably the major oxidant. If the pH is less than 4.5, however, hydrogen peroxide (H202) is most likely the dominant oxidant (Penkett et al., 1979; Calvert et al., 1985). Computer simulations suggest that at pH 3, H202 pathways account for more than 99% of the oxidation of S(IV) (sulphur dioxide and related sulfites) to S(VI) (sulphuric acid and related - 85 - sulfates) (Graedel et al., 1985). The pH of the fogwater sample used in this experiment is 3. 7, and the concentration of ozone in the air which was equilibrated with the fog is only 1 µg!m3. Furthermore, in the dark (thermal control) experiments, no measurable oxidation of dissolved S02 was observed. This indicates that the oxidation of S02 by 03 and by Mn, Fe catalyzed oxygenation is negligible. 78 -~ ::!. '-" ,....., 76 -.....> '-" r'-l...... 74 72'---'-~...._~....__._~....._~..____._~...._~....__. 0 30 60 90 120 150 Time (min) Figure 4.4.3 Photo-oxidation of disso/ved S02 in fog water with 313 nm light We can consider that all the S02 converted into sulfate was oxidized by H102. Therefore, the total formation of peroxide must correspond to the sum of the net accumulated concentration of H102 formed and the concentration of dissolved S02 converted into sulfate. This is presented in Fig.4.4.4 against irradiation time. The degradation of DOC and oxalic acid is also illustrated in Figure 4.4.4. - 86 - 12 10 * Q -.;:= 8 Cl:I ...."" Q,I 6 =C.J u=Q 4 2 Oxalate 0 0 30 60 90 120 150 Time (min) Figure 4.4.4 Total concentration of peroxide formed and degradation of DOC and oxalic aeid. Closed eire/es: total peroxide (µM) (the sum of H202formed and bisu/fite dep/eted); open eire/es: oxalic aeid (µM); open squares: DOC (mgll). The total fonnation of peroxide reaches 9.3 µM (1 µM = l0-6 mol/l) after 150 min irradiation. lt has been shown that the accumulation rate of H202 increases with the concentration of oxalic acid and with DOC. In fogwater, the maximum values of oxalate concentration and DOC can be 10 times higher than those in the fog sample used. Thus, a higher photochemical fonnation rate of H202 is expected. - 87 - 5. General discussion and environmental signiticance 5.1 Mecbanism for tbe pbotocbemical formation of hydrogen peroxide and decomposition of oxalic acid The above experimental results are consistent with the proposed mechanisms. H202 is produced via photochemical/chemical cycling of Fe(III)/Fe(II) oxalato complexes as shown in the schematic diagram in Figure 5.1. lt is not surprising, however, that in previous studies of the photolysis of the solutions of potassium ferrioxalate the formation of H202 in the presence of 02 was not observed. These studies involved such high concentrations of oxalate that the oxygen, at relatively low concentrations in aqueous solution, was unable to compete for the photochemical intermediate. Thus the oxalate radical formed in the primary photolysis would reduce another Fe(III) oxalato complex rather than reduce oxygen (Figure 2.1): In our system, the concentrations of Fe(Ill)-oxalato complexes (~10 µM) relative to oxygen (250 µM) are much lower, so reactions (2) and (5) (in Figure 2.1) must be taken into account. Thus an oxygen molecule reacts either with an excited complex or with an oxalate radical to form 02 ·-, which subsequently leads to H202 formation. The extent of H202 formation shows that reactions (2) and (5) compete efficiently with reaction (4) (see Figure 2.1). That the amount of H202 fonned can be many times greater than the initial concentration of Fe(III) in the presence of high oxalate concentrations (Figure 4.1.4), is additional evidence for the occurrence of redox cycling - 88 - between Fe(III) and Fe(II) (Figure 5.1 ). This cycling continues until the oxalate is consurned. nOx Fe{III)Oxn a Fe(!Il) Ox Fe(II) ox·- nOx 02 Figure 5.1 Scheme for the photochemical/chemical cycling of iron and the formation of H202. - 89 - lt also should be noted that other transition metals, particularly Cu(II)/Cu(I), could also play a significant role in cloud water H202 formation, by catalyzing the disproportionation of H02·/02·- involving a pathway analogous to that described for iron (von Piechowski, 1991). For comparison, the relevant reactions and rate constants are presented in Table 2.1. lt should be kept in mind, however, that complex formation might greatly change the reactivity of transition metal ions. Further identification of organic and inorganic ligands for the transition metals in atmospheric water and studies of the interactions of the ligands with transition metals and the interaction of the complexes formed with sunlight are required. 5.2 Environmental Significance 5.2.1 Photolysis of Fe(III)-organic complexes can be a major source of H202 in atmospheric waters. · Although direct dissolution of gaseous H202 and the disproportionation of H02 radicals scavenged by cloud droplets are currently considered to be the major sources of H202 in the atmospheric liquid phase, our results show that the photochemical generation of H202 from aqueous Fe(III)-oxalate complexes could be important. If, for example, 1 µM Fe(III) and 5 µM oxalate occur in an atmospheric droplet, their net contribution would be 3.7 nM s-1 of H202 at midday in autumn (Figure 4.1.2). This value is similar to the calculated transfer rate of gas phase H202 and the disproportionation of H02 free radicals taken up by small cloud droplets ( In many atmospheric water droplets, dissolved iron, and oxalic, glyoxalic and pyruvic acids are present in concentrations typically ranging from 0.1 to a few tens of micromoles (Zuo and Hoigne, 1992; Joos and Baltensperger, 1991; Steinberg et al„ 1985; Norton et al„ 1983). These facts indicate that the photolysis of Fe(IIl)-oxalato and other Fe(III)-organic complexes can be a major source of HzOz in cloud, fog and rainwater. 5.2.2 Photolysis of Fe(III)-organic complexes can generate a significant amount of OH radicals in atmospheric liquids. Photolysis of Fe(IIl)-oxalato complexes produces Fe(II) ions. This has important implications for the production of OH radicals through the Fenton reaction as assumed in previous studies (Graedel et al, 1986). At night the Fenton reaction can be a major source of OH radicals in cloud droplets. Our preliminary experimental results have also shown that OH radicals are formed when high concentrations of Fe(III)-oxalate are illuminated. The formation rate of OH radicals is a few times higher than that in the absence of oxalate. Recently, Zepp et al. (1992) have also quantified the formation of OH radicals by photo-Fenton reactions in aqueous solutions containing Fe(IIl)-oxalate complexes (100 µM) and HzOz. Based on the rate constants given in Table 2.1, however, Fe(II) is reoxidized mostly by HOz·/Oz·- and leads to Hz02 formation under conditions of day-time irradiation. The production of OH radicals in atmospheric water droplets has considerable significance, as these radicals can oxidize a wide variety of natural and anthropogenic organic and inorganic substances. lt would be interesting to further investigate the mechanism and kinetics of OH radical formation in the concentration ranges of dissolved iron and oxalate associated with atmospheric waters. - 91 - 5.2.3 Pbotolysis of oxalic acid catalyzed by iron as a major sink for atmospheric oxalic acid. Removal of oxalic acid from the atmosphere has received little attention to date. Grosjean (1983,1989) studied the photolysis of pyruvic acid in the atmospheric gas phase. He estimated that glyoxalic, pyruvic, and oxalic acids were photodegraded in a few ho'urs. In the presence of cloud droplets, however, these acids will mostly transfer into the liquid phase due to their large Henry's coefficients. Oxalic, glyoxalic and pyruvic acid have been found at significant concentrations in rain, mist, fog and cloud water (Zuo and Hoigne, 1992; Joos and Baltensperger, 1991; Steinberg et al„ 1985; Norton et al., 1983; Kawamura et al„ 1985; Baltensperger and Kern, 1988). Based on the experimental results presented in Chapter 4, the fast conversion of atmospheric oxalic and keto acid by photochemical/chemical redox cycling of iron becomes a major sink for these compounds. Tue half life for the photolysis of oxalic acid in the presence of iron(ill) is on the order of a few minutes. lt should be noted that photolysis of the dissolved oxalic and keto acids can be expected to become nearly continuous, because these acids are continuously resupplied by the "fresh" air and by in-cloud oxidation of other organic compounds. Correspondingly, the photochemical formation of H202 also becomes continuous. 5.2.4 Photochemical/chemical cycling of Fe(Ill)/Fe(II) complexes provides a significant source of dissolved iron in natural surface waters. Tue presence of oxalic acid, even at micromolar levels, considerably accelerates the photochemical/chemical cycling of iron. Tue complexation by oxalate not only increases the quantum yield for the photoreduction of Fe(ill) from 0.14 (Faust and Hoigne, 1990) to 1.2 at 313 nm, but also extends the absorption band into the visible region and increases - 92 - the absorption coefficient (see Figure 2.2). Under conditions of similar sunlight intensities, the ferric oxalate in this study photolyzed about 2 orders of magnitude faster than the Fe(OH)2+ previously studied (Faust and Hoigne, 1990). In September noon sunlight, the photoreduction of Fe(III) oxalato complexes have a half-life of 17 seconds. Combining this half-life time with the steady-state ratio of Fe(III)/Fe(II), we estimate that the photochemical /chemical cycling time for Fe(III)-Fe(Il)-Fe(III) is on the order of minutes on a clear September day at noon. This cycling time is expected to decrease by a factor of three for a sunny June day at noon, even for in-cloud conditions as light-screening effects are compensated by light scattering. This photochemical/chemical redox cycling of complexed iron species could also involve iron (hydr)oxide surfaces. Iron is photodissolved (Siffert and Sulzberger, 1991) and subsequently precipitated to the ocean and fresh surface water where it becomes a source of dissolved iron. - 93 - 6. Conclusions (a) This study has shown that the formation of hydrogen peroxide by the sunlight induced photolysis of Fe(IU)-oxalato complexes could be a maj~r source of H202 in the atmospheric liquid-phase. Both dissolved iron and oxalic acid are ubiquitous pollutants in cloud, fog and rainwater. The photolysis of Fe(III)-oxalate complexes produces oxalate radicals, these radicals subsequently reduce oxygen to the superoxide ion leading to the formation of H202. About one H202 molecule is generated by every 1.8 ± 0.7 molecules of oxalate consumed under the conditions used. The formation rate of hydrogen peroxide is related to the sunlight intensity, pH and the concentrations of iron and oxalate. Increasing the concentrations of either oxalate or Fe(III) within the concentration range typically found in atmospheric waters accelerates the formation of H202. Under conditions typical of the cloudwater, H202 was produced at a significant rate, for instance, when a few micromolar oxalic acid and 1 µM dissolved iron are present, the accumulation rate of H202 is 3.7 nM s-1 in September-noon- sunlight. This rate is similar to or greater than that expected from the absorption of gaseous H202 and disproportionation of H02 radicals scavenged by cloud droplets under similar environmental conditions. (b) The preliminary experiments have also shown that these photochemical processes can produce a significant amount of OH radicals in cloudwater via subsequent photo-Fenton reactions. (c) The iron catalyzed photolysis of oxalic acid provides a major sink for atmospheric oxalic acid. The rate of depletion of oxalic acid increases linearly with increasing concentration of oxalic acid for low concentrations, bot remains constant at higher concentrations. Tue rate also increases as the - 94 - initial concentration of ferric ion is increased up to a level corresponding to the complete complexation of oxalic acid. At solar noon in September, oxalic acid is photolyzed with a half life of a few minutes even in clean cloud droplets. ( d) In the presence of oxalic acid, the photochemical/chemical cycling of Fe(ill)/Fe(II) complexes is very fast. In September-noon-sunlight, the photoreduction of Fe(ill)-oxalato.complexes has a half-life of 17 s in the absence of 02. Varying the ferrioxalate speciation and pH (from pH 1 to 5) only slightly influences the rate of reduction of Fe(ill) complexes. In the presence of 02, the Fe(II) formed is rapidly reoxidized to Fe(III). Tue cycling time for Fe(ill)-Fe(II)-Fe(III) is in the order of minutes on a clear September day at noon. - 95 - 7. References Andreae M. 0., Browell E. V., Garstang M., Gregory G. L., Harriss R. C., Hill G. F., Jacob D. J., Pereira M. C., Sachse G. W., Setzer A. W., Silva Dias P. L., Talbot R. W., Torres A. L., and S. C. Wofsy (1988) Biomass-buming emissions and associated haze layers over Amazonia. J. geophys. Res. 93, 1509-1527. Bader H., Sturzenegger V. and Hoigne J. (1988). Photometrie method for the determination of low concentrations of hydrogen peroxide by the peroxidase catalyzed oxidation of N,N-diethy-p-phenylenediamine (DPD). Wat. Res. 22, 1109-1115. Baltensperger U. and Kern S. (1988) Determination of mono~ and divalent cations and anions in small fog samples by ion chromatography. J. Chromatogr. 439, 121-126. Balzani V. and Carassiti V. (1970) Photochemistry of coodination compounds. Academic press, London. Barth M. C., Hegg D. A., Hobbs P. V., Walega J. G., Kok G. L., Heikes B. G. and Lazrus A. L. (1989) Measurements of atmospheric gas-phase and aqueous-phase hydrogen peroxide concentrations in winter on the east coast of the United States. Tellus 41B, 61-69. Bauer R. F. and Smith W. M. (1965) The mono-oxalato complexes of iron(ill). Can. J. Chem. 43, 2755-2762. Becker K. H., Brockman K. T. and Bechara J. (1990) Production of hydrogen peroxide in forest air by reaction of ozone .with terpenes. Nature 246, 256-258. Behar D., Czapski G., Rabani J., Dorfman L. M. and Schwarz. H. A. (1972) The acid dissociation constant and decay kinetics of the perhydroxyl radical. J. phys. Chem. 74, 3209-3213. - 96 - Bebra P. and Sigg L. (1990) Evidence for redox cycling of iron in atmospheric water droplets. Nature 344, 419-421. Benkelberg H.J. Deister U. and Wameck P. (1990) OH quantum yield for the photodecomposition of Fe(Ill) hydroxo complexes in aqueous solution and the reaction of OH with hydromethanesulfonate. p263- 269, in Physico-chemical behaviour of atmospheric pollutants. edited by G.Restelli and G.Angeletti, Kluwer Academic Publishers. Dortrecht. Bielski B. H. J. (1978) Reevaluation of the spectral and kinetics properties of 645-649. the H02 and o2- free radicals. Photochem. Photobiology 28, Bielski B. H. J.; Cabelli, D.; Arudi, R.; Ross, A. (1985) Reactivity of HOz/02- radicals in aqueous solution. J. Phys. Chem. Ref. Data 14, 1041-1051. Bielski B.H.J. and Allen, A.O. (1977) Mechanism of the disproportionation of superoxide radicals. J. phy. Chem. 81, 1048-1050. Bowen H. J. M. (1966) Trace Elements in Biochemistry, 241 pp., Academic, San Diego, Califomia. Brimblecombe P. and Spedding D. J. (1975) The dissolution of iron from ferric oxide and pulverized fuel ash. Atmos. Environ. 9, 835-838. Bufalini J. J., Lancaster H. T., Namie G. R. and Gay B. W. (1979) Hydrogen peroxide formation from the photooxidation of formaldehyde and its presence in rainwater. J. Environ. Sei. Health A14, 135-141. Calvert J. G., Lazrus A., Kok G. L., Heikes B. G., Walega J. G., Lind, J. and Cantrell C, A. (1985) Chemical mechanism of acid generation in the troposphere. Nature 317, 27-35. Chameides W. L. and Davis D. D. (1982) The free radical chemistry of cloud droplets and its impact upon the composition of rain. J. geophys. Res. 87, 4863-4877. - 97 - Chandler A. S., Choularton T. W., Dollard G. J., Eggleton A. E. J., Gay M. J., Hill T. A , Jones B. M. R., Tyler B. J., Bandy B. J. and Penkett S. A. (1988) Measurements of H202 and S02 in clouds and estimates of their reaction rate. Nature 336, 562-565. Clarke A. G. and Radojevic M. (1987) Oxidation of S02 in rainwater and its role in acid rain chemistry. Atmos. Environ. 21, 1115-1123. Cooper G. D. and de Graff B. A., (1971) On the photochemistry of the ferrioxalate system. J. phys. Chem. 75, 2897-2902. Daum P. H., Kleinman L. 1., Hills A. J., Lazrus A. L., Leslie A. C. D., Busness K. and Boatman J. (1990) Measurement and interpretation of concentrations of H202 and related species in the upper midwest during summer. J. geophys Res. 95, 9857-9871. Davison W. (1979) Soluble inorganic ferrous complexexs in natural waters. Geochemica et Cosmochimica Acta 43, 1693-1696. Deng Y. (1992) Formation and dissolution of aquatic iron(lll)(hydr)oxides -implications for redox cycling of iron in natural waters Ph.D. Dissertation, ETH, Zürich. No. 9724. Dlugi R., Jordan S. and Manegold S. (1985) Chemische Reaktionen und Aerosolverhalten in Rauchfahnen mit kondensierender Atmosphäre. Kernforschungszentrum Karlsruhe, KfK 3817, Karlsruhe, F.R.G. Draper W. M. and Crosby D. G. (1983b) Photochemical generation of superoxide radical anion in water. J. Agric. Food Chem. 31, 734- 737. Ennis C. A., Lazrus A. L. Zimmerman P. R. and Monson R. K. (1990) Flux determinations and physiological response in the exposure of red spruce to gaseous hydrogen peroxide, ozone, and sulfur dioxide. Tellus 42B, 183-199. - 98 - Evans M. G., Santappa M. and Uri N. (1951) Photoinitiated free radical polymerization of vinyl compounds in aqueous solution. J. Polymer Sei. 7, 243-260. Faust B. C. and Hoigne J. (1990) Photolysis of Fe(IIl)-hydroxy complexes as sources of OH radicals in clouds fog and rain Atmos. Environ. 24A, 79-89. Finlayson-pitts B. J. and Pitts J. N. Jr. (1986) Atmospheric Chemistry: fundamentals and experimental techniques. John Wiley & Sons lnc. New York. Foster P. M. (1969) The oxidation of sulfur dioxide in power station plumes. Atmos. Environ. 3, 157-175. Fung C. S., Misra P. K., Bloxam R. and Wong S. (1991) A numerical experiment on the relative importance of H102 and 03 in aqueous conversion of S02 to S042-. Atmos. Environ. 24A, 411-423. Fuzzi S., Orsi G., Nardini G., Facchini M. C., McLaren S., Mclaren E. and Mariotti E. (1988) Heterogeneous processes in the Po Valley radiation fog. J. geophys. Res. 93, 11,141-11,151. Goldberg E. D. (1971) Atmosphere dust, the sedimentary cycle, and man. Geophysics 1, 117-132. Gunz D.W. and Hoffmann R. (1990) Atmospheric chemistry of peroxides: a review. Atmos. Environ. 24A, 1601-1633. Gervat G. P., Clark P. A., Marsh A. R. W., Teasdale 1., Chandler A. S., Choularton T. W., Gay M. J., Hill M. K., and Hill T. A. (1988) Field evidence for the oxidation of S02 by H102 in cap clouds. Natuer 333, 241-243. Graedel T. E., and Goldberg K. 1. (1983) Kinetic studies of raindrop chemistry 1. Inorganic and organic processes. J. geophys Res., 88, 10,865-10,882. - 99 - Graedel T. E., Mandich M. L. and Weschler C. J. (1986) Kinetics Model studies of atrnospheric droplet chemistry 2. Homogeneous transition metal chemistry in raindrops. J. geophys. Res. 91, 5205-5221. Graedel T. E., Weschler C. J. and Mandich M. L. (1986) Influence of transition metal complexes on atrnospheric droplet acidity. Nature 317, 240-242. Grosjean D. (1983) Atrnospheric reactions of pyruvic acid. Atmos. Envrion. 17, 2379-2382. Grosjean D. (1989) Organic acids in southem Califomia air: ambient concentrations, mobile source emissions, in situ formation and removal processes. Environ. Sei. Techno/. 23, 1506-1514. Haag W. R. and Hoigne J. (1985) Photo-sensitized oxidation in natural water via OH radicals. Chemosphere 14, 1659-1671. Hansen L. D., Silberman D., Fisher G. L. and Eatough D. J. (1984) Chemical speciation of elements in stark-collected, respirable-size, coal fly ash. Environ. Sei. Techno/. 18, 181-186. Hatchard C. G., Parker A. C. (1956) A new sensitive chemical actinometer. II. Potassium ferrioxalate as a standard chemical actinometer. Proc. Roy Soc. Lon. Ser. A, 235, 518-536. Heikes B. G., Lazrus A. L., Kok G. L., Kunen S. E., Gandrud B. W., Gitlin S. N. and Sperry P. D. (1982) Evidence for aqueous phase hydrogen peroxide synthesis in the troposphere. J. geophys. Res. 87, 3045- 3051. Hoffmann M. R. and Edwards J. 0. (1975) Kinetics of the oxidation of sulfite by hydrogen peroxide in acidic solution. J. phy. chem. 79, 2096- 2098. - 100 - Ho P. C. (1986) Photooxidation of 2,4-dinitrotoluene in aqueous solution in the presence of hydrogen peroxide. Environ. Sei. Teehnol. 20, 260- 267. Hoigne J. and Bader H. (1975) Ozonization of water: Role of hydroxyl radicals as oxidizing intennediates. Scienee.190, 782-784. Hood G. C. and Reilly (1960) Ionization of strong electrolytes. VIII. Temperature coefficient of dissociation of strong acids by proton magnetic resonance. J. ehern. Phys. 32, 127-130. Johnson C. A., Sigg L. and Zobrist J. (1987) Case studies on the chemical composition of fogwater: the influence of local gaseous emissions. Atmos. Environ. 21, 2365-2374. Jones P. Tobe M. L. and Wyhne-Jones W. F. K. (1959) Hydrogen peroxide + water mixture Part 5 Stabilization against the iron perchlorate catalyzed decomposition of hydrogen peroxide. Trans Faraday Soe. 55, 91-97. Joos F. and Baltensperger U. (1991) A field study on chemistry, S(IV) oxidation rates and vertical transport during fog conditions. Atmos. Environ. 25A 217-230. Kawamura K. and Kaplan 1. R. (1987) Motor exhaust emissions as a primary source for dicarboxylic acids in Los Angeles ambient air. Environ. Sei. Teehnol. 21, 105-110. Kawamura K., Steinberg S. and Kaplan 1. R. (1985) Capillary GC determination of short-chain dicarboxylic acids in rain, fog, and mist. Intern. J. Environ. Anal. Chem., 19, 175-188. Kelly T.J., Daum P.H. and Schartz S.E. (1985) Measurements of Peroxides in Cloudwater and Rain. J. geophys. Res. 90, 7861-7871. Kleinman L. 1. (1984) Oxidant requirements for the acidification of precipitation. Atmos. Environ. 18, 1453-1458. - 101 - Kleinman L. I. (1986) Photochemical formation of peroxides in the boundary layer. J. geophys. Res. 91, 10,889-10,904. Kormann C„ Bahnemann D. W. and Hoffmann M.R. (1988) Photocatalytic production of H202 and organic peroxides in aqueous suspensions of TiOz, ZnO, and desert sand. Envir. Sei. Techno/. 22, 798-806. Kusakabe K„ Aso S„ Hayashi J„ Isomura K. and Morooka S. (1990) Decomposition of humic acid and reduction of trihalomethane formation potential in water by ozone with U.V. irradiation. Wat. Res. 24, 781-785. Lacroix S. (1949) Action de quelques substances organiques sur le vegetaux. Ann. Chim. 4, 5-27. Leitis E. (1979) An investigation into the chemistry of the UV-ozone purification process, Annual Report to the National Science Foundation, NTIS Document Number PB 296485. Lelieveld J. and Crutzen P. J. (1990) Influences of cloud photochemical processes on tropospheric ozone. Nature 343, 227-233. Mallant R. K. A. M„ Slanina J„ Masuch G. and Kettrup A. (1986) Experiments on H202-containing fog exposures of young trees. in Aerosols (ed. Lee S. D. et al.) 901-912 Lewis, Chelsea Michigan. Martell A. M. and Smith R. M. (1977) Critical stability constants, Vol. 3 Plenum, New York. Martin L. R„ Easton M. P., Foster J. W. and Hill M. W. (1989) Oxidation of hydroxymethanesulfonic acid by Fenton's reagent. Atmospheric Environment 23, 563-568. Masuch G„ Kettrup A„ Mallant R. K. A. M. and Slanina J. (1986) Effects of H202-containing acidic fog on young trees. Intern. J. Environ. Anal. Chem. 27, 183-213. - 102 - McElroy W. J. (1986) Source of hydrogen peroxide in cloudwater. Atmos. Environ. 20, 427-438. Meagher J. F., Olszyna K. J., Weatherford F. P. and Mohnen V. A. (1990) Tue availability of H202 and 03 for aqueous phase oxidation of S02. the question of linearity. Atmos. Environ. 24A, 1825-1829. Möller D. (1989) Tue possible role of H202 in new-type forest decline. Atmos. Environ. 23, 1625-1627. Munger J. W., Jacob D. J., Waldman J. M. and Hoffmann M. R. (1983) Fogwater chemistry in an urban atmosphere. J. geophys. Res. 88, C9, 5109-5121. Neftel A., Jacob P. and Klockow D. (1984) Measurements of hydrogen peroxide in polar ice samples. Nature 311, 43-45. Norton R. B., Roberts J. M. and Huebert B. J. (1983) Tropospheric oxalate. Geophys. Res. Lett. 10, 517-520. Olszyna K. J., Meagher J. F. and Bailey E. M. (1988) Gas-phase, cloud and rain-water measurements of hydrogen peroxide at a high-elevation site. Atmos. Environ. 22, 1699-1706. Parker C. A. and Hatchard C. G. (1959) Photodecomposition of complex oxalates, some preliminary experiments by flash photolysis. J. phys. Chem. 63, 22-26. Penkett S. A., Jones B. M. R., Brice K. A. and Eggleton A. E. J. (1979) Tue importance of atmospheric 03 and H202 in oxidizing S02 in cloud and rainwater. Atmos. Environ. 13, 123-137. Petasne R. G. and Zika R. G. (1987) Fate of Superoxide in coastal sea water. Nature 325, 516-518. Peirson D. H., Cawse P. A. and Cambray R. S. (1974) Chemical uniformity of airbome particulate material, and a maritime effect. Nature, 251, 675-679. - 103 - Rabani J. and Nielsen S. 0. (1969) Adsorption spectrum and decay kinetics radiolysis. J. phy. Chem. of o2- and H02 in equeous solutions by pulse 73, 3736-3744. Rowan N. S„ Hoffman M. Z. and Milburn R. M. (1974) lntermediates in the photochemistry of tris(oxalato)cobaltate(ill) ion in aqueous solution. free and coordinated radicals. J. Amer. Chem. Soc. 96, 6060-6067; Rush J. D. and Bielski H. J. (1985) Pulse radiolytic studies of the reactions of H 02/02- with Fe(Il)/Fe(III) ions. The reactivity of H02/02- wi th ferric ions and its implication on the occurrence of the Haber-Weiss reaction. J. phys. Chem. 89, 5062-5066. Schaap W. B„ Laitinen H. A. Bailar J. C. Jr. (1954) Polarography of iron oxalates, malonates and succinates. J. Am. Chem. Soc. 76, 5862-5872. Schneider W. (1988) Iron hydrolysis and the biochemistry of iron -the interplay of hydroxide and biogenic ligands. Chimia 42, 9-20. Schwartz S. E. (1984) Gas- and aqueous-phase chemistry of H02 in liquid water clouds. J. geophys. Res. 89, 11,589-11,598. Siffert C. and Sulzberger B. (1991) Light-induced dissolution of hematite in the presence of oxalate: a case study. Langmuir. 7, 1627-1634. Sigg A. and Neftel A. (1991) Evidence for a 50% increase in H202 over the past 200 years from a Greenland ice core. Nature 351, 557-559. Sigg L. and Behra P. (1991) EAWAG, unpublished results. Simonaitis R„ Olszyna K. J. and Meagher J. F. (1991) Production of hydrogen peroxide and organic peroxides in the gas phase reactions of ozone with natural alkenes. Geophys. Res. Lett. 18, 9-12. Steinberg S„ Kawamura K. and Kaplan 1. R. (1985) The determination of a keto acids and oxalic acid in rain, fog and mist by HPLC. Intern. J. Environ. Anal. Chem. 19, 251-260. - 104 - Stumm W. and Morgan J. J. (1981) Aquatic Chemistry Wiley-Interscience New York. Sturzengger V. (1989) Wasserstoffperoxid in oberflaechengewaessern: photochemische produktion und abbau. Ph.D. Thesis No. 9004 ETHZ, Zurich. Sulzberger B., Soter D., Siffert C., Banwart S., and Stumm W. (1989) Dissolution of Fe(III) (hydr)oxides in natural waters; laboratory assessment on the kinetics controlled by surface coordination. Mar. Chem. 28, 127-144. Sutton J. (1952) Formation of a ferric-perchlorate ion-pair complex. Nature 169, 71-72. Tamura H., Goto K., Yotsnyanagi T. and Nagayama M. (1974) Spectro- photometric determination of iron(II) with 1,10-phenanthroline in the presence of large amounts of iron (III). Talanta 21, 314-318. Tsai W., Cohen Y., Sakugawa H. and Kaplan 1. R. (1991) Hydrogen peroxide levels in Los Angeles: a screening-level evalution. Atmos. Environ. 25B, 67-78. Tuncel S. G., Olmez 1., Parrington J. R., Gordon G. E. and Stevens R. K. (1985) Composition of fine particle regional sulfate component in Shenandoah Valley. Envir. Sei. Techno[. 19, 529-537. Vincze L. and Papp S. (1987) Individualquantum yields of Fe3+0Xn2-Hm+ complexes in aqueous acidic solutions (OX2- = C2042-, n = 1 - 3, m = 0, 1). J. Photochem. 36, 289-296. von Piechowski M. (1991) Der einfluss von kupferionen auf die redox- chemie des atmosphaerischen wassers: kinetische untersuchungen. PhD Thesis ETHZ No. 9512. Waldman J. M., Lioy P. J., Zelenka M., Jing L., Lin Y. N. He Q. C., Qian Z. M., Chapman R. and Wilson W. E. (1991) Wintertime - 105 - measurements of aerosol acidity and trace elements in Wuhan, a city in central China. Atmos. Environ. 25B 113-120. Walling C. (1975) Fenton's reagent revisited. Ace. Chem. Res. 8, 125-131. Weinstein-Lloyd J. and Schwartz S. E. (1991) Low-intensity radiolysis study of free-radical reactions in cloudwater: H202 production and destruction. Environ. Sei. Technol. 25, 791-800. Weschler C. J., Mandich M. L. and Graedel T. E. (1986) Speciation, photosensitivity, and reactions of transition metals ions in atmospheric droplets. J. geophys. Res. 91, 5189-5204. Weiss J. (1947) On the nature of the "active" oxalic acid. Faraday Soc. 2, 188-196. Westall J. C. (1979) "MICROQL", a chemical equilibrium program in Basic. Intemal report, EAW AG, Dübendorf. Williams P. T., Radojevic M. and Clarke A. G. (1988) Dissolution of trace metals from particles of industrial origin and its influence on the composition of rainwater. Atmos. Environ. 22, 1433-1442. Xue H., Goncalves M. L. Reutlinger M. Sigg L. and Stumm W. (1991) Copper(l) in fogwater: determination and interactions with sulfite. Environ. Sei. Technol. 25, 1716-1722. Yamamoto S. and Back R. A. (1985) Tue gas.phase photochemistry of oxalic acid. J. phys. Chem. 89, 622-625. Yamamoto Y., Niki E., Shiokawa H. and Kamiya Y. (1979) Ozonation of organic compounds. 2. Ozonation of phenol in water. J. org. Chem. 44, 2137-2142. Yoshizumi K., Aoki K., Nouchi 1., Okita T., Kobayashi T., Kamakura S. and Tajima M. (1984) Measurements of the concentration in rainwater and of the Henry's law constant of hydrogen peroxide. Atmos. Environ. 18, 395-401. - 106 - Zafiriou 0. C. (1977) Marine organic photochemistry previewed. Mar. Chem. 5, 497-522. Zepp R., Braun A.M., Hoigne J. and Leenheer J.A.(1987) Photoproduction of hydrated electrons from natural organic solutes in aquatic environments. Environ. Sei. technol. 21, 485-490. Zepp R. (1989) U.S. Environmental Protection Agency, College Station Road, Athens, GA, personal communication. Zepp R., Faust B. and Hoigne J. (1992) Hydroxyl radical formation in aqueous reactions (pH 3-8) of iron (II) with hydrogen peroxide: the photo-Fenton reaction. Envir. Sei. Techno[. 26, 313-319. Zhuang G., Duce R.A., and Kester D.R. (1990) The dissolution of atmospheric iron in surface seawater of the open ocean. J. geophys. Res. 95, C 16207-16216. Zhuang G., Zhen Y., Duce R. A. and Brown P. R. (1992) Link between iron and sulphur cycles suggested by detection of Fe(II) in remote marine aerosols. Nature 355, 537-539. Zika R., Saltzman E., Chameides W. L. and Davis D. D. (1982) H202 levels in rainwater collected in south Florida and the Bahama islands. J. geophys. Res. 87, 5015-5017. Zika R. G. (1981) Marine organic photochemistry in Marine Organic Chemistry edited by Duursma E. K. and Dawson R. Elsevier Scientific Publishing company Amsterdam. Zika R. G. and Saltzman E.S. (1982) Interaction of ozone and hydrogen peroxide in water: implications for analysis of H202 in air. Geophys. Res. Lett. 9, 231-234. Zuo Yuegang and Jürg Hoigne (1992) Formation of H202 and depletion of oxalic acid by photolysis of Fe(ID)-oxalato complexes in atmospheric waters. (1992) Environ. Sei. Techno[. (in press) - 107 - Curriculum vitae YuegangZuo Hubei, China 17Sept1958 Born in Echeng, in Echeng 1964 - 1970 Primary school 1970 -1976 High school in Echeng 1976 Diploma Echeng 1976 - 1978 Huarong High School, Chemistry Lecturer 1978 - 1982 Wuhan University in Analytical Chemistry 1982 Bachelor' s Degree Academia Sinica and Institute of 1982 - 1984 Graduate School, Research Center for Environmental Chemistry (late called Sinica, Beijing Eco-environmental Sciences) , Academia in Environmental Chemistry 1984 Master' s Degree for Eco-environmental Sciences, 1984 - 1988 Research Center Academia Sinica, Beijing, China Research Assistant (1984 - 1986) Research Associate (1986 - 1988) Water Institute for Water Resources and 1988 - 1992 Swiss Federal / Swiss Federal Pollution Control (BAW AG), Dübendorf Institute of Technology (ETH), Zurich Ph.D. Thesis