Fate of Atmospheric HOX and Peroxides in the Particle Phase

by

Kaitlin Marie Badali

A thesis submitted in conformity with the requirements for the degree of Master of Science (M.Sc.)

Department of Chemistry University of Toronto

© Copyright by Kaitlin Marie Badali 2014

ii

Fate of Atmospheric HOX and Peroxides in the Particle Phase

Kaitlin Marie Badali

Master of Science (M.Sc.)

Department of Chemistry

University of Toronto

2014

Abstract

The odd hydrogen radicals hydroxy (OH) and hydroperoxy (HO2), collectively referred to as

HOX, along with peroxides, play a significant role in determining the oxidative capacity of

the atmosphere. These oxidants are key species in the removal of many harmful trace gases

in the atmosphere. When transition metal ions are present in an aqueous aerosol, they will

efficiently cycle with HOX radicals and affect the rate at which sink species are formed. The

plausibility of the coupling of copper and iron redox cycles is investigated using aqueous

solutions of organics and transition metals to represent aqueous aerosols. The formation of

secondary organic aerosols (SOA) has been shown to generate hydroperoxides. The yields

and stability of peroxides in SOA generated and aged under a variety of conditions is

investigated, with both atmospheric and indoor air quality implications.

iii

Acknowledgments

Throughout my time at the University of Toronto, I have had an overwhelming amount of

support from both the Environmental Chemistry department and my family. Firstly, I would

like to thank my supervisor, Jon Abbatt, for his advice and guidance throughout my graduate

degree. His knowledge of atmospheric chemistry and supportive supervision technique

allowed me to develop my critical thinking and research skills. I would also like to thank my

committee members, Andre Simpson for his input into this thesis, and Jamie Donaldson for

his insightful questions. Thank you to the Donaldson group for the lending of the

fluorometer, to the Mabury group for use of the Solar Simulator, and the Murphy group for

the modelling software. Thank you to Jinqiu Mao at Princeton University for his guidance

and suggestions through the first project.

I would like to thank the incredible collection of students and post-docs that comprise the

Abbatt research group – everyone brings such a unique contribution to the group dynamic

and helped me in different ways. Firstly, I would like to thank Maria for her friendship and

scientific support through both projects – I couldn’t be happier to share an office with you,

gracias! Thank you to Shouming for his SOA expertise and helping me through the chamber

experiments, to Dana for her extensive knowledge of the HRP-DCF assay, Ran for his

mechanistic help and insightful interpretations, and Rob for showing me the ropes the

summer I arrived. A special thanks to Jacquie and Michaela for your friendship and coffee

breaks! Thank you to each and every member of the group for helping me along in a special

way.

Finally, I would like to thank my family for the unwavering support through my entire

university career. My parents both emotionally and financially supported me along these past

7 years, I could not have done it without you and I am forever in debt (literally!). Graham,

you have been my rock through this all and I couldn’t thank you enough for everything

you’ve done. To my brothers, Alex and Chris, my grandma, Jen, Chester and everyone else –

your love and encouragement have helped me along the way!

iv

Table of Contents

Abstract ..................................................................................................................................... ii

Acknowledgments ................................................................................................................... iii

Table of Contents ..................................................................................................................... iv

List of Tables ........................................................................................................................... vi

List of Figures ......................................................................................................................... vii

Chapter 1 Introduction and Overview ...................................................................................... 1

1.1 Oxidative Capacity of the Atmosphere ........................................................................... 1

1.2 HOX Radicals in the Atmosphere .................................................................................... 2

1.3 Peroxides in the Atmosphere .......................................................................................... 5

1.3.1 Peroxide Measurement ............................................................................................. 7

1.4 Research Goals ................................................................................................................ 9

1.4.1 Effects of Transition Metal Redox Coupling on HOX Chemistry ............................ 9

1.4.2 Production and Stability of Peroxides in Secondary Organic Aerosol .................. 10

1.5 References ..................................................................................................................... 11

Chapter 2 Aqueous HOX Chemistry with Copper-Iron Redox Coupling ............................... 14

2.1 Introduction ................................................................................................................... 14

2.2 Methods ......................................................................................................................... 17

2.2.1 Photolysis Samples ................................................................................................. 17

2.2.2 Horseradish Peroxidase Assay ............................................................................... 19

2.2.3 Kinetic Modelling................................................................................................... 20

2.3 Results and Discussion ................................................................................................. 21

2.3.1 Acetaldehyde photolysis......................................................................................... 21

2.3.2 Aldehyde and Copper Photolysis ........................................................................... 23

2.3.3 Aldehyde and Iron Photolysis ................................................................................ 24

2.3.4 Aldehyde, Copper, and Iron ................................................................................... 27

2.3.5 Kinetic model results .............................................................................................. 30

2.3.6 Organic Iron Complexes ........................................................................................ 31

2.4 Conclusions and Atmospheric Implications ................................................................. 33

v

2.5 References ..................................................................................................................... 35

2.6 Appendix ....................................................................................................................... 37

Chapter 3 Yields and Stability of Peroxides from Secondary Organic Aerosol Formation ... 40

3.1 Introduction ................................................................................................................... 40

3.2 Methods ......................................................................................................................... 45

3.2.1 SOA Collection using Flow Tube .......................................................................... 45

3.2.2 SOA Collection using Environmental Chamber .................................................... 46

3.2.3 SOA Storage and Extraction .................................................................................. 48

3.2.4 Photolysis in Solar Simulator ................................................................................. 48

3.2.5 Indoor Photolysis using Fluorescent Lights ........................................................... 49

3.3 Results and Discussion ................................................................................................. 49

3.3.1 Peroxide Yields in SOA ......................................................................................... 49

3.3.2 Effect of SOA Mass Loading ................................................................................. 52

3.3.3 On-Filter Stability................................................................................................... 52

3.3.4 Effect of Parafilm ................................................................................................... 54

3.3.5 In-Solution Stability ............................................................................................... 55

3.3.6 SOA Photolysis in Solar Simulator ........................................................................ 56

3.3.7 SOA Photolysis under Fluorescence Lights ........................................................... 60

3.4 Conclusions and Environmental Implications .............................................................. 60

3.5 References ..................................................................................................................... 63

Chapter 4 Implications and Future Work ............................................................................... 65

4.1 Environmental Implications .......................................................................................... 65

4.2 Future Work .................................................................................................................. 66

vi

List of Tables

Table 2.1 – Reactions and rate constants included in kinetic model ...................................... 21

Table 2.2 – Hydrogen peroxide production from photolyzed aldehyde samples ................... 22

Table 2.3 – Peroxide production in a system containing 1 mM acetaldehyde with varying

copper to iron ratios ................................................................................................................ 27

Table 2.4 – Modelled reaction rates of iron(II) and HO2 as copper concentration increases . 28

Table 3.1 – Review of peroxide yields from fresh SOA samples reported in past studies .... 50

Table 3.2 – Peroxide yields of filters stored for 24 hours using parafilm .............................. 54

Table 3.3 – Peroxide yields from photolysis of SOA on-filter and in-solution for 7 hours

under fluorescent lights........................................................................................................... 60

vii

List of Figures

Figure 1.1 – Mechanism through which hydrogen peroxide produces fluorescent signal in

HRP-DCF assay, with DCF formation and HRP catalytic cycle.............................................. 8

Figure 1.2– HRP-DCF assay sensitivities to H2O2, ROOH (t-butylhydroperoxide, cumene

hydroperoxide), and ROOR (di-t-butyl peroxide) .................................................................... 9

Figure 2.1 – Coupling of copper and iron redox cycles with fates of HOX radicals .............. 15

Figure 2.2 – Experimental setup of Suntest CPS Solar Simulator fitted with xenon arc lamp,

with respective spectrum. ....................................................................................................... 18

Figure 2.3– Calibration curve of hydrogen peroxide obtained using the HRP-DCF assay .... 19

Figure 2.4 – Peroxide production from 3 hour photolysis of 3 mM acetaldehyde in solar

simulator ................................................................................................................................. 22

Figure 2.5 – Hydrogen peroxide concentrations produced from 1-hour photolysis of samples

containing 1 mM acetaldehyde, 1 mM CuSO4, and/or 1 μM FeSO4 ...................................... 23

Figure 2.6 – Modelled production of hydrogen peroxide in a system containing 1 mM

aldehyde with varying copper concentrations at 60 minutes of photolysis ............................ 24

Figure 2.7 - Concentration of hydrogen peroxide produced from solutions containing 1 mM

acetaldehyde with iron(II) varied from 0.01 to 1.0 µM. ......................................................... 25

Figure 2.8 – Modelled Fe(III) concentration in a system containing 1 mM aldehyde with

Cu/Fe ratios of 0 (1 μM Fe(II)), 0.1 (0.1 μM Cu(II) with 1 μM Fe(II)), 1 (1 μM Cu(II) with 1

μM Fe(II)), and 1000 (1 mM Cu(II) with 1 μM Fe(II)). Dashed lines represent simulations in

which the coupling reaction (R3) was removed from model. ................................................ 29

Figure 2.9 – Generation of hydrogen peroxide from photolysis of organic acids both with and

without iron(III). ..................................................................................................................... 31

Figure 2.10 – Proposed mechanism of oxidation of acetaldehyde to form oxalic acid, which

is shown to complex with iron(III). ........................................................................................ 32

Figure 2.11 – Kinetic Model output of key species from system containing 1 mM

acetaldehyde ........................................................................................................................... 37

Figure 2.12 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 mM copper(II) ............................................... 37

Figure 2.13 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 μM iron(II) .................................................... 38

viii

Figure 2.14 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 mM copper(II) and 1 μM iron (II)................. 39

Figure 3.1 – Deposition location of particles within the body as a function of particle size . 40

Figure 3.2 – Mechanism of limonene ozonolysis, showing select oxidation peroxide-related

products................................................................................................................................... 42

Figure 3.3 – Simplified schematic of flow tube set up for SOA generation and collection. .. 46 Figure 3.4 – Simplified schematic of environmental chamber set up for SOA generation and

collection................................................................................................................................. 47

Figure 3.5 – Stability of peroxides in SOA stored on-filter at room temperature and in freezer

at -20 °C. ................................................................................................................................. 53

Figure 3.6 – Stability of peroxides in aqueous extract of SOA sample at room temperature. 55

Figure 3.7 – Peroxide yields from SOA photolysis using a xenon arc lamp in the solar

simulator. ................................................................................................................................ 59

1

Chapter 1

Introduction and Overview

1.1 Oxidative Capacity of the Atmosphere

Atmospheric oxidants have been the focus of much research in recent decades due to

their direct influence on air quality and climate. These species play a primary role in

dictating the oxidative capacity of the atmosphere, which is a measure of its ability to cleanse

itself of reduced gases.[1, 2] The oxidation of species such as volatile organic compounds

(VOCs), methane (CH4), carbon monoxide (CO), halocarbons, and other greenhouse gases

will generate products of lower volatility.[2-4] As a result, these gases will condense into

particles or cloud droplets and can ultimately be removed from the atmosphere through wet

and dry deposition. While oxidation chemistry within the particle and aqueous phases has

been recognized, the most heavily studied systems, and thus most thoroughly understood, are

primarily centred on gas phase reactions. With recent environmental concerns growing

around issues such as climate change, acid rain, and greenhouse gas emissions, there is a

need to understand the oxidation chemistry of the atmosphere that will occur in all phases.

Three common oxidants within the atmosphere are the hydroxy (OH) and

hydroperoxy (HO2) radicals and peroxides.[5] The OH and HO2 radicals, collectively known

as HOX, will rapidly cycle with each other in the presence of different atmospheric species.

Peroxides, on the other hand, are considered to be a reversible sink of HOX, as they are

formed through a radical termination reaction of HO2. This sink is short lived, however, as

photolysis or other reactions of the peroxides will regenerate HOX through the production of

OH radicals.[6] Consequently, the abundance and fate of these three oxidants are closely

linked and dependent on each other.

Aqueous chemistry can occur in the atmosphere in either cloud droplets or aqueous

aerosols. The liquid water content of aqueous aerosols tends to be quite low, ranging from

7.6x10-7

– 4x10-5

cm3/m

3.[7] Consequently, species that are found within the aqueous phase

of the aerosol tend to be highly concentrated. Due to the comparatively high volume of water

in cloud droplets, the concentrations of species within the aqueous phase are much more

2

dilute. Typical cloud water content ranges from 0.05-3 cm3/m

3.[7] The surface to volume

ratio of aerosol particles will also be substantially different from cloud droplets, and will

therefore influence the rates and significance of reactions that will occur. As a result of the

high surface area, it is expected that surface reactions will be more significant in aerosol

chemistry than in cloud droplets.[8]

The peroxides and HOX radicals are expected to become incorporated into the particle

and aqueous phases primarily through uptake from the gas phase.[9] Many reactions that

occur in the aqueous phase are also known to produce HOX radicals and peroxides, and thus

can also be important aqueous phase sources of oxidants. Both peroxides and the HO2 radical

are highly water soluble, resulting in a rapid uptake into aqueous aerosols and cloud droplets

from the gas phase. Past studies have reported the presence of peroxides and OH radicals

within aqueous cloud droplets, demonstrating their importance as oxidants within the

aqueous phase.

A lack of information regarding indoor air chemistry has also been recognized in

recent years, resulting in an increase of studies looking at indoor air quality. Like outdoors,

some of the major oxidants indoors are believed to be peroxides, the hydroxy and

hydroperoxy radicals.[10] Several studies in the past decade have evaluated the production of

peroxide species within secondary organic aerosol that is known to form indoors. These

studies will be discussed in detail in Chapter 3. Overall, the abundance and reactivity of these

oxidants indoors is not yet well characterized.

1.2 HOX Radicals in the Atmosphere

The primary source of HOX in the atmosphere is the production of the hydroxyl

radical from the photolysis of ozone in the presence of water vapour. This reaction, as shown

below in R1 and R2, will occur at wavelengths of 315 nm and below.[4] The primary source

of HO2 radicals is the atmosphere is the photolysis of carbonyls, such as formaldehyde

shown in R3.[8] Additionally, HO2 radicals are generated through oxidation processes in the

atmosphere that are initiated by OH radicals.

3

O3 + hν O(1D)

O(1D) + H2O 2 OH

CH2O + hν CO + 2 HO2

Atmospheric concentrations of the HOX radicals vary widely and are dependent on

surrounding conditions. The concentrations of OH will be significantly influenced by both

the humidity and solar irradiation levels of its environment, as well as the concentrations of

various reactants.[2] Under typical clean-air conditions, the HO2 radical is approximately

100 times more abundant than OH.[4] A summary of recent HO2 measurements report an

average mixing ratio of 8.7 ± 3.7 pptv.[11] Despite its low concentration, the effectiveness of

OH as an oxidant is accounted for by it highly reactive nature. Under typical tropospheric

conditions, the lifetime of the OH radical is only 1 second.[12, 13] Conversely, the HO2

radical is less reactive within the gas-phase, typically having a lifetime of 100 seconds.[4]

Past studies have investigated oxidation by both OH and HO2 radicals within aerosol

particles and cloud droplets, stressing the importance of considering heterogeneous oxidation

by HOX when studying the chemistry of the atmosphere. These radicals can become

incorporated into the particle or aqueous phase through uptake processes, or formed by

reactions at the surface or within particles. Typically, uptake of HOX into particles will only

occur through interactions with HO2, as a result of its higher gas phase concentrations and

solubility in the aqueous phase. With a pKa value of 4.69, HO2 will tend to dissociate into O2-

and H+

upon uptake into the aqueous phase.[14] While the concentration of O2- may be

much lower than the protonated form in some particles, its reaction rates can be faster than

those of HO2.[11] Concentrations of OH in the atmosphere tend to be too low for its uptake

to be significant. The typical lifetime of the OH radical is also much shorter than its average

uptake rate.[8]

The uptake of HOX radicals into cloud droplets is typically limited by gas phase

diffusion to the droplet surface.[15] This limitation is not significant for aerosol particles,

however, due to their small size. One consequence of the small particle size of aerosols is

that dissolved species tend to re-evaporate from the particle at much faster rates than in a

cloud droplet.[8] It is common for atmospheric particles to possess an organic coating on

2 O2

R1

R2

R3

4

their surfaces. In these particles, the uptake across the coating will be inhibited, resulting in a

decrease in the levels of HOX radicals within these aerosols.[15]

The uptake rates of HO2 radicals into cloud droplets and aerosol particles under

various conditions have been the subject of much recent research. Due to the solubility and

dissociation of HO2 in the aqueous phase, it has been found that its uptake coefficients (γ)

can range from 0.05-0.1.[16] The rate of uptake has been found to be enhanced when the

particle is enriched in species such as transition metals, increasing to uptake coefficients as

high as >0.2 in the presence of copper.[15, 17] Much of the chemistry that will occur within

the particle following uptake is still not well understood, especially due to the presence of

varying species in different particle types.

In addition to uptake from the gas phase, HOX may form in cloud droplets through

both reactive and photolytic processes. Some of the most significant sources of HOX within

cloud droplets are the Fenton reaction (R4), the photolysis of NO3- (R5), NO2

- (R6),

hydrogen peroxide (R7), Fe(OH)2+

(R8), all generating OH radicals.[18-20]

Fe2+

+ H2O2 OH + OH- + Fe

3+

NO3- + hν OH + NO2

NO2- + hν OH + NO + OH

-

H2O2 + hν 2 OH

Fe(OH)2+

+ hν OH + Fe2+

One of the important HOX sinks in aerosol particles is the reduction of a transition

metal (M) ion by the HO2/O2- radical.[8] This reaction can either result in the reversible loss

of HOX through the production of hydrogen peroxide (R9), or a permanent loss of HOX

through the formation of oxygen (R10). These reactions are predicted to be relevant to

aerosols that contain mineral dust or crustal material, as there will be sufficient levels of

dissolved transition metal to allow the reactions to proceed.[8] Another possible sink of HOX

R4

R5

R6

R7

R8

H2O

H2O

5

in aerosol particles is the reaction of HO2 with dissolved halide species, however this sink is

typically only significant in the marine boundary layer where there are high levels of

halides.[21]

Mx + O2

- M

(x+1) + H2O2 + 2OH

-

Mx + O2

- M

(x+1) + O2

1.3 Peroxides in the Atmosphere

Of the peroxide family, hydrogen peroxide is most prevalent in the atmosphere,

typically accounting for 70-90% of total peroxide content.[22] The remaining portion of the

peroxides is thought to be composed of small organic peroxides, with methyl hydroperoxide

and peroxyacetic acid being most frequently measured.[23] Typical mixing ratios of

hydrogen peroxide are 0.5-5 ppb in the gas phase, while concentrations are approximately

250 µM in the aqueous phase of cloud droplets.[23, 24] Peak peroxide levels have been

measured in the summer during high pollution events, when solar irradiation and peroxide-

forming species are at their highest.[23]

One of the primary loss processes of HOX radicals in the atmosphere is the radical

termination reaction of HO2 to generate peroxides.[23] This reaction can form either

hydrogen peroxide through the self reaction of HO2 (R11), or organic hydroperoxides

through the reaction of HO2 with RO2 (R12).[25] These pathways to produce peroxides in

the atmosphere can occur both in the gas phase and within the aqueous phase. In the aqueous

phase, as the liquid water content of the aerosols or cloud droplets increases, the rate at

which peroxide will be formed through this self reaction will decrease.[23] This reduction in

the production rate is a result of the dilution of HO2 as the volume of water increases.

HO2 + HO2 H2O2 + O2

HO2 + RO2 ROOH + O2

R11

R12

H2O

R9

R10

6

The ozonolysis of volatile organic compounds has also been shown to be a potential

source of peroxides, through first forming an energetic radical known as a Criegee

intermediate. This intermediate can go on to react with water, generating hydroperoxides.[24,

26, 27]

Similar to the processes that occur with HO2 radicals, peroxides can be taken up into

the aqueous and particle phases from the gas phase.[27] Their favourable partitioning to the

aqueous phase is a result of their high solubility in water and their abundance in the gas

phase.[24, 28] Once in the aqueous phase, the peroxides may be held tightly in solution by

binding with ions present in solution. This results in the formation of a peroxohydrate adduct,

similar to a solvation shell that would form with water.[29]

In past studies, the photolysis of cloud and fog droplets has been shown to generate

peroxides.[30] Although this production was observed in samples collected from various

locations, the species responsible for this photolytic production have not yet been

identified.[18] The production of peroxides was also observed in aqueous solutions of

ambient automobile exhaust particles.[31] This generation is expected to be due to the redox

cycle of HO2 radicals with transition metals, quinone, and organic species that may be found

in various atmospheric particles.[24, 31]

In the aqueous phase, peroxides are known to be the most significant pathway to

oxidize sulphur. This oxidation, shown in R13 and R14, results in the production of sulphuric

acid.[32] Although peroxide is expected to participate in other oxidation reactions in the

aqueous or particle phases, these reactions have not been well characterized and their

significance is not known. As mentioned in the previous section, the peroxides may

potentially volatilize into the gas phase from small particles with low liquid water content.

SO2 + H2O HSO3- + H

+

HSO3- + H2O2 HSO4

- + H2O

Peroxides are of particular importance in the atmosphere because of the significant

role they play in oxidizing sulphur in the aqueous phase to form acid rain.[33] Peroxides also

have the potential to interfere with processes within the biosphere by damaging delicate

R13

R14

7

leaves and other components of various plant species. Once the peroxides have deposited

onto the plant, they have been shown to rapidly oxidize chlorophyll to a colourless

compound.[23] This hinders the plant’s photosynthesis ability, causing their eventual demise.

The association of peroxides with particles is also a concern for human health. Past

studies have shown that particles will carry oxidants much deeper into the respiratory system

than their gas-phase equivalents.[29, 34] Once deep within the lungs, the peroxides will

create an imbalance in oxidants and antioxidants within the cells, causing oxidative

stress.[35] The precise effects that peroxides will have on human health are not well

characterized, however, with various studies publishing conflicting results about the impacts

that particle-bound oxidants will have within the body.

1.3.1 Peroxide Measurement

Due to the challenges associated with measuring radical species, past studies have

used the measurements of peroxides as an indication of OH and HO2 radical levels.[32]

Currently, no real-time online measurement technique exists for the quantification of

peroxides. Within the community, the most commonly employed methods for detection

involve offline electrochemical reduction, chemiluminescence, and fluorescence

techniques.[27, 30] In both projects of this thesis, a fluorescent assay technique based on a

method by Keston and Brandt[36] was used to evaluate peroxide content within aqueous

samples. The experimental methods of this technique will be described in detail in Chapter 2.

a)

8

Figure 1.1 – Mechanism through which hydrogen peroxide produces fluorescent signal in

HRP-DCF assay, with a) DCF formation (adapted from Black & Brandt)[37] and b) HRP

catalytic cycle (adapted from Berglund et al.)[38]

In the assay, horseradish peroxidase (HRP) is used as a catalyst in the reaction of

peroxide with a fluorescent agent, dichlorofluorescein (DCF). The pathway through which

peroxide will generate the fluorescent molecule is found in Figure 1.1. A stock solution of

the fluorescent molecule is in the diacetate form (DCFH-DA), and must be converted to the

“activated” hydrolyzed form (DCFH) through a reaction with a base. The conversion of the

hydrolyzed molecule into the fluorescent form (DCF) in part a) of Figure 1 requires the

removal of two hydrogen atoms.[37] The catalytic cycle of horseradish peroxidase, seen in

part b) of Figure 1.1, consumes two hydrogen atoms to return the catalyst to its ground state

following its reaction with hydrogen peroxide.[38] This cycle will remove the protons from

DCFH, producing one fluorescing DCF for each hydrogen peroxide molecule consumed.

In a characterization study of the assay conducted by our group, the sensitivities of

the HRP-DCF assay to hydrogen peroxide and organic peroxides (ROOH, ROOR) were

tested. The results, as shown in Figure 1.2, demonstrate the assay’s overwhelming sensitivity

to hydrogen peroxide in aqueous samples.[39] Due to the extremely low responses by

organic peroxides, it is assumed that their contributions to the fluorescence signal will be

b)

9

negligible. Thus, it is inferred that fluorescence measured in all experimental samples will

exclusively come from hydrogen peroxide.

Figure 1.2– HRP-DCF assay sensitivities to H2O2, ROOH (t-butylhydroperoxide, cumene

hydroperoxide), and ROOR (di-t-butyl peroxide).[39]

1.4 Research Goals

In this thesis, two projects evaluate distinct aqueous- and particle-phase chemistry that

has the potential to affect levels of oxidants within the atmosphere. It is hoped that the results

of these projects will help to improve the understanding of atmospheric oxidation chemistry.

The research conducted in the second project not only occurs in the troposphere, but also has

implications for indoor air quality. In further developing our knowledge of oxidants within

aerosol particles, we will be able to better evaluate potential health effects caused by

inhalation of these particles.

1.4.1 Effects of Transition Metal Redox Coupling on HOX Chemistry

In the first project, the influence of transition metal-catalyzed redox cycles on

aqueous HOX chemistry was evaluated. It has long been accepted that transition metals will

10

participate in a redox cycles with HOX radicals in an aqueous environment.[40] In a recent

modelling paper by Mao et al., however, it was suggested that the cycles of copper and iron

would be coupled together when simultaneously present in an aerosol.[11] This chemistry

could have implications on the levels of HOX in the atmosphere, as it could form an

irreversible sink through the production of water.

In order to evaluate this coupling chemistry, the production of peroxides was

monitored in systems containing HOX radicals with only copper, only iron, and both copper

and iron together. Aqueous acetaldehyde was used as a photolytic source of HOX in all

samples, and hydrogen peroxide was measured using the HRP-DCF fluorescence assay. A

kinetic model was created in order to predict the fates and reaction pathways of different

species within each system studied experimentally.

1.4.2 Production and Stability of Peroxides in Secondary Organic Aerosol

In the second project, the production and stability of peroxides within aerosol

particles was investigated. Several past studies have reported the generation of

hydroperoxides during the oxidation of volatile organic compounds to form secondary

organic aerosols.[41] These peroxides will typically be formed as a temporary sink of gas-

phase HOX radicals. It is currently expected that most peroxides are found within particles as

the result of an uptake process from the gas phase, so this chemistry could represent an

important source of peroxide within the particle phase. The particles have been shown to

form both in the atmosphere and indoors, resulting in the potential to have effects on both

climate and health.

The yields of peroxides were measured in secondary organic aerosol formed through

the ozonolysis of α-pinene and limonene using the HRP-DCF fluorescence assay. To gain a

more complete understanding of these peroxides, their stabilities were measured under dry

and aqueous conditions at room temperature and within a freezer. The photolytic stability of

the peroxides was determined through a series of experiments using a solar simulator to

represent atmospheric conditions and fluorescent light source to replicate an indoor

environment.

11

1.5 References

[1] R.G. Prinn, The cleansing capacity of the atmosphere, in Annual Review of

Environment and Resources, 2003, pp. 29-57.

[2] R. Commane, C.F.A. Floquet, T. Ingham, D. Stone, M.J. Evans, and D.E. Heard,

Observations of OH and HO2 radicals over West Africa, in Atmospheric Chemistry

and Physics, 2010, pp. 8783-8801.

[3] J. Lelieveld, and P.J. Crutzen, The role of clouds in tropospheric photochemistry, in

Journal of Atmospheric Chemistry, 1991, pp. 229-267.

[4] D.R. Crosley, The measurement of OH and HO2 in the atmosphere, in Journal of the

Atmospheric Sciences, 1995, pp. 3299-3314.

[5] A.M. Thompson, The oxidizing capacity of the Earth's atmosphere - probably past

and future changes, in Science, 1992, pp. 1157-1165.

[6] L. Jaegle, D.J. Jacob, P.O. Wennberg, C.M. Spivakovsky, T.F. Hanisco, E.J.

Lanzendorf, E.J. Hintsa, D.W. Fahey, E.R. Keim, M.H. Proffitt, E.L. Atlas, F. Flocke,

S. Schauffler, C.T. McElroy, C. Midwinter, L. Pfister, and J.C. Wilson, Observed OH

and HO2 in the upper troposphere suggest a major source from convective injection

of peroxides, in Geophysical Research Letters, 1997, pp. 3181-3184.

[7] H. Herrmann, Kinetics of aqueous phase reactions relevant for atmospheric

chemistry, in Chemical Reviews, 2003, pp. 4691-4716.

[8] D.J. Jacob, Heterogeneous chemistry and tropospheric ozone, in Atmospheric

Environment, 2000, pp. 2131-2159.

[9] Z.M. Chen, H.L. Wang, L.H. Zhu, C.X. Wang, C.Y. Jie, and W. Hua, Aqueous-phase

ozonolysis of methacrolein and methyl vinyl ketone: a potentially important source of

atmospheric aqueous oxidants, in Atmospheric Chemistry and Physics, 2008, pp.

2255-2265.

[10] S. Gligorovski, and C.J. Weschler, The Oxidative Capacity of Indoor Atmospheres, in

Environmental Science & Technology, 2013, pp. 13905-13906.

[11] J. Mao, S. Fan, D.J. Jacob, and K.R. Travis, Radical loss in the atmosphere from Cu-

Fe redox coupling in aerosols, in Atmospheric Chemistry and Physics, 2013, pp. 509-

519.

[12] J. Lelieveld, W. Peters, F.J. Dentener, and M.C. Krol, Stability of tropospheric

hydroxyl chemistry, in Journal of Geophysical Research-Atmospheres, 2002, pp. 11.

[13] R. Sommariva, A.L. Haggerstone, L.J. Carpenter, N. Carslaw, D.J. Creasey, D.E.

Heard, J.D. Lee, A.C. Lewis, M.J. Pilling, and J. Zador, OH and HO2 chemistry in

clean marine air during SOAPEX-2, in Atmospheric Chemistry and Physics, 2004,

pp. 839-856.

[14] S.E. Schwartz, Gas-phase and aqueous-phase chemistry of HO2 in liquid water

clouds, in Journal of Geophysical Research-Atmospheres, 1984, pp. 1589-1598.

[15] M. Mozurkewich, P.H. McMurry, A. Gupta, and J.G. Calvert, Mass accommodation

coefficient for HO2 radicals on aqueous particles, in Journal of Geophysical

Research-Atmospheres, 1987, pp. 4163-4170.

[16] J.P.D. Abbatt, A.K.Y. Lee, and J.A. Thornton, Quantifying trace gas uptake to

tropospheric aerosol: recent advances and remaining challenges, in Chemical

Society Reviews, 2012, pp. 6555-6581.

12

[17] F. Taketani, Y. Kanaya, and H. Akimoto, Kinetics of heterogeneous reactions of HO2

radical at ambient concentration levels with (NH4)(2)SO4 and NaCl aerosol

particles, in Journal of Physical Chemistry A, 2008, pp. 2370-2377.

[18] B.C. Faust, C. Anastasio, J.M. Allen, and T. Arakaki, Aqueous-phase photochemical

formation of peroxides in authentic cloud and fog waters, in Science, 1993, pp. 73-75.

[19] H. Herrmann, B. Ervens, H.W. Jacobi, R. Wolke, P. Nowacki, and R. Zellner,

CAPRAM2.3: A chemical aqueous phase radical mechanism for tropospheric

chemistry, in Journal of Atmospheric Chemistry, 2000, pp. 231-284.

[20] B. Ervens, S. Gligorovski, and H. Herrmann, Temperature-dependent rate constants

for hydroxyl radical reactions with organic compounds in aqueous solutions, in

Physical Chemistry Chemical Physics, 2003, pp. 1811-1824.

[21] B.M. Matthew, I. George, and C. Anastasio, Hydroperoxyl radical (HO2 center dot)

oxidizes dibromide radical anion (Br-center dot(2)-) to bromine (Br-2) in aqueous

solution: Implications for the formation of Br-2 in the marine boundary layer, in

Geophysical Research Letters, 2003, pp. 5.

[22] Y. Ren, A.J. Ding, T. Wang, X.H. Shen, J. Guo, J.M. Zhang, Y. Wang, P.J. Xu, X.F.

Wang, J. Gao, and J.L. Collett, Measurement of gas-phase total peroxides at the

summit of Mount Tai in China, in Atmospheric Environment, 2009, pp. 1702-1711.

[23] D.W. Gunz, and M.R. Hoffmann, Atmospheric chemistry of peroxides - a review, in

Atmospheric Environment Part a-General Topics, 1990, pp. 1601-1633.

[24] C. Arellanes, S.E. Paulson, P.M. Fine, and C. Sioutas, Exceeding of Henry's law by

hydrogen peroxide associated with urban aerosols, in Environmental Science &

Technology, 2006, pp. 4859-4866.

[25] W. Hua, Z.M. Chen, C.Y. Jie, Y. Kondo, A. Hofzumahaus, N. Takegawa, C.C.

Chang, K.D. Lu, Y. Miyazaki, K. Kita, H.L. Wang, Y.H. Zhang, and M. Hu,

Atmospheric hydrogen peroxide and organic hydroperoxides during PRIDE-PRD'06,

China: their concentration, formation mechanism and contribution to secondary

aerosols, in Atmospheric Chemistry and Physics, 2008, pp. 6755-6773.

[26] A.S. Hasson, A.W. Ho, K.T. Kuwata, and S.E. Paulson, Production of stabilized

Criegee intermediates and peroxides in the gas phase ozonolysis of alkenes 2.

Asymmetric and biogenic alkenes, in Journal of Geophysical Research-Atmospheres,

2001, pp. 34143-34153.

[27] P. Mertes, L. Pfaffenberger, J. Dommen, M. Kalberer, and U. Baltensperger,

Development of a sensitive long path absorption photometer to quantify peroxides in

aerosol particles (Peroxide-LOPAP), in Atmospheric Measurement Techniques,

2012, pp. 2339-2348.

[28] A.S. Hasson, and S.E. Paulson, An investigation of the relationship between gas-

phase and aerosol-borne hydroperoxides in urban air, in Journal of Aerosol Science,

2003, pp. 459-468.

[29] S.K. Friedlander, and E.K. Yeh, The submicron atmospheric aerosol as a carrier of

reactive chemical species: Case of peroxides, in Applied Occupational and

Environmental Hygiene, 1998, pp. 416 - 420.

[30] A. Mutzel, M. Rodigast, Y. Iinuma, O. Boege, and H. Herrmann, An improved

method for the quantification of SOA bound peroxides, in Atmospheric Environment,

2013, pp. 365-369.

13

[31] Y. Wang, C. Arellanes, and S.E. Paulson, Hydrogen Peroxide Associated with

Ambient Fine-Mode, Diesel, and Biodiesel Aerosol Particles in Southern California,

in Aerosol Science and Technology, 2012, pp. 394-402.

[32] M.H. Lee, B.G. Heikes, and D.W. O'Sullivan, Hydrogen peroxide and organic

hydroperoxide in the troposphere: A review, in Atmospheric Environment, 2000, pp.

3475-3494.

[33] H. Liang, Z.M. Chen, D. Huang, Y. Zhao, and Z.Y. Li, Impacts of aerosols on the

chemistry of atmospheric trace gases: a case study of peroxides and HO2 radicals, in

Atmospheric Chemistry and Physics, 2013, pp. 11259-11276.

[34] X. Chen, and P.K. Hopke, Secondary organic aerosol from alpha-pinene ozonolysis

in dynamic chamber system, in Indoor Air, 2009, pp. 335-345.

[35] F. Tao, B. Gonzalez-Flecha, and L. Kobzik, Reactive oxygen species in pulmonary

inflammation by ambient particulates, in Free Radical Biology and Medicine, 2003,

pp. 327-340.

[36] A.S. Keston, and R. Brandt, Fluorometric analysis of ultramicro quantities of

hydrogen peroxide, in Analytical Biochemistry, 1965, pp. 1-&.

[37] M.J. Black, and R.B. Brandt, Spectrofluorometric anaylsis of hydrogen-peroxide, in

Analytical Biochemistry, 1974, pp. 246-254.

[38] G.I. Berglund, G.H. Carlsson, A.T. Smith, H. Szoke, A. Henriksen, and J. Hajdu, The

catalytic pathway of horseradish peroxidase at high resolution, in Nature, 2002, pp.

463-468.

[39] D. Aljawhary, Personal Communication, 2014.

[40] L. Deguillaume, M. Leriche, K. Desboeufs, G. Mailhot, C. George, and N.

Chaumerliac, Transition metals in atmospheric liquid phases: Sources, reactivity, and

sensitive parameters, in Chemical Reviews, 2005, pp. 3388-3431.

[41] K.S. Docherty, W. Wu, Y.B. Lim, and P.J. Ziemann, Contributions of organic

peroxides to secondary aerosol formed from reactions of monoterpenes with O-3, in

Environmental Science & Technology, 2005, pp. 4049-4059.

14

Chapter 2

Aqueous HOX Chemistry with Copper-Iron Redox Coupling

2.1 Introduction

It is common for atmospheric aerosol particles and cloud droplets to be enriched with

transition metals through both anthropogenic and natural processes.[1-4] Iron is the most

abundant transition metal within the Earth, and consequently many crustal aerosols and dust

particles contain iron, as well as manganese and copper.[4-6] The metal ions in these

particles tend to be highly soluble at the acidic pH levels that are commonly found in cloud

water and aqueous aerosols, and thus many transition metals will be found in the aqueous

phase.[2, 7] The aerosol concentrations of transition metals are heavily dependent on the

source and aqueous environment of the particles, and as a result widespread values have been

reported. Literature values of total transition metal concentration are 0.02-670 µM, 0.0006-

23 µM, and 0.00055-80 µM for iron, copper, and manganese respectively.[2] There is some

discrepancy with these concentration values when applied to cloud chemistry however, as the

speciation of the metals is not known.[4] Reported concentrations of dissolved iron, copper

and manganese within cloud and fog droplets range from 0.02-134 µM, 0.0007-7 µM, and

0.0009-4 μM respectively.[2, 8, 9] It is believed that much of the iron(II) is found in the free

Fe2+

form; on the other hand, it is expected that almost all of the dissolved iron(III) will be

complexed with organic compounds, SO42-

, and OH-.[10] It is expected that much of the

copper present in the cloud water will be found in organic complexes, regardless of its

oxidation state.[4] The reactivity of copper complexes with HOX radicals and other

atmospheric species is significantly lower than that of the free copper ions.[10, 11]

It has long been known that transition metals can catalytically cycle with HOX

radicals; in particular, copper, iron and manganese most efficiently participate in the redox

cycling.[2, 12] The cycle is initiated by the reduction of the metal by reaction with HO2,

shown in R1 and pathway a) of Figure 2.1. The reduced metal can then either regenerate HO2

through a reaction with oxygen (R2a), or consume an additional HO2 radical to form

hydrogen peroxide (R2b, pathway b) in Figure 2.1). Both pathways will regenerate

15

Figure 2.1 – Coupling of copper and iron redox cycles with fates of HOX radicals (adapted

from Mao et al. [12])

the oxidized metal, allowing the cycle to continue indefinitely. Although the reaction rate

constants of the copper cycle are much larger than the other transition metals, the iron cycle

may be just as significant due to atmospheric concentrations of iron typically being an order

of magnitude greater than copper.[12] At commonly reported atmospheric concentrations of

transition metals, it is expected that the catalytic transition metal-catalyzed sink of HOX will

be more significant than the self-reaction of HO2 to form H2O2.[4]

Mn+ HO Mn-1+ O + H

+

Mn-1+ O H+

Mn+ HO

Mn-1+ HO H+

Mn+ H O

It is possible that both copper and iron will be present within an aqueous aerosol

simultaneously, with iron being the more abundant species. A recent modeling paper

published by Mao et al. proposed the coupling of copper and iron redox cycles within

R1

R2a

R2b

a)

b)

c)

d)

e)

f)

16

atmospheric aerosols.[12] Under these conditions, it is suggested that iron(III) will be

reduced to iron(II) through reaction with copper(I) (R3, pathway c) in Figure 2.1), in place of

a reaction with HO2 (R1).

Cu(I) + Fe(III) Cu(II) + Fe(II)

The fate of the iron(II) can have significant implications on the chemistry of aqueous HOX,

ultimately affecting the oxidative capacity of the atmosphere. Products such as hydrogen

peroxide are considered to be reversible sinks of HOX, as photolysis or reaction of peroxide

will regenerate OH radicals. Alternatively, the production of water is a permanent sink of

HOX, as it will be unable to regenerate the radicals.[12] A reaction of iron(II) with HO2 (R4a,

pathway d in Figure 2.1) will result in an overall null cycle of HOX through production of

H2O2. If the iron(II) reacts with H2O2 (R4b, pathway e in Figure 2.1), the production of OH

will regenerate some of the consumed HOX. The net result of this pathway, however, is a net

HOX sink through the formation of water. The reaction of iron(II) with OH (R4c, pathway f

in Figure 2.1) will be a net sink of HOX as only water will be produced.

Fe(II) + HO2: HO + HO Cu/Fe H O + O

Fe(II) + H2O2: HO + H O Cu/Fe OH + O + H O

Fe(II) + OH: HO + OH Cu/Fe O + H O

The purpose of this project is to investigate the proposed coupling of copper and iron

in an aqueous environment in a laboratory setting in order to evaluate the likelihood of this

chemistry proceeding in the environment. Acetaldehyde was used as a photolytic source of

aqueous-phase HOX radicals in all experiments. Upon photolysis, acetaldehyde will

dissociate into either HCO or H radicals, which will react with oxygen to produce HO2 (R5-

R6).[13] The effects that the presence of copper and iron will have on the cycling of HOX

radicals was evaluated by measuring the concentration of hydrogen peroxide in samples

R4b

R4c

R3

R4a

17

while varying the concentrations of the transition metals. In samples that contain only the

transition metal with the HOX radicals, it is expected that the same amount of peroxide will

be measured as in a system that contains only HOX radicals. In samples that contain both

transition metals with the HOX radicals, it is anticipated that a decrease in peroxide will be

observed if the coupling reaction is proceeding. The coupling will result in the loss of HOX

radicals through the production of water, thus limiting the amount of peroxide that will form.

CH3CHO + hν CH3 + HCO

HCO + O2 HO2 + CO

CH3CHO + hν CH3CO + H

H + O2 HO2

2.2 Methods

2.2.1 Photolysis Samples

Transition metal stock solutions of copper (II) (CuSO4∙5H2O, 25 mM), iron (II)

(FeSO4∙7H2O, 0.25 mM), and iron (III) (FeCl3, 0.25 mM) were prepared daily using

deionized water (18 mΩ, Millipore). A stock solution of acetaldehyde (CH3CHO, ≥99.5%,

Sigma-Aldrich) was prepared and immediately diluted for the photolysis study. Photolysis

samples were prepared using a variety of concentrations of acetaldehyde (0.01-3.0 mM),

copper(II) (0.01-3.0 mM), iron(II) (1 μM), and iron(III) (1 μM). Solution pH was adjusted to

approximately 3.0 using sulphuric acid (∼2 M, H2SO4), measured before and after

photolysis. For all samples, the pH did not change over the course of the experiment.

Two equivalent sample sets were prepared through transferring 5 mL of each solution

to custom-made quartz holders; one sample set was photolyzed while the other was kept in

the dark at room temperature as a dark control. A solar simulator (Figure 2.2, Suntest CPS,

Atlas) fitted with a xenon arc lamp was employed as the light source. In a previous

characterization study of this solar simulator, actinometry was used to report that the

R5b

R6a

R6b

R5a

18

intensity of the system was 8.5 einstein/min.[14] Inside the simulator, samples were placed

on a cooled plate held at 4 °C and covered with a quartz case. A cooling stream of air was

flowed through the simulator while the lamp was in operation. During a typical photolysis

experiment, aqueous sample temperatures did not exceed 33 °C. Unless otherwise stated,

samples were photolyzed for 1 hour, followed by immediate analysis of hydrogen peroxide

concentration.

Figure 2.2 – a) Experimental setup of Suntest CPS Solar Simulator fitted with xenon arc

lamp, with respective b) spectrum (not calibrated).

40x1015

30

20

10

0

Flu

x (

pho

tons

s-1 c

m-2

)

800700600500400300

Wavelength (nm)

a)

b)

19

2.2.2 Horseradish Peroxidase Assay

A stock solution of ’,7’-dichlorofluorescein diacetate (1 mM DCFHDA,

C24H14Cl2O7, Sigma-Aldrich) was prepared in methanol and stored at -20°C. The DCFHDA

was converted to the hydrolyzed ’,7’-dichlorofluorescein (DCFH) form by adding 1.0 mL

DCFHDA to 4.0 mL sodium hydroxide (0.01 M, NaOH) and allowing the hydrolysis to

proceed for 30 minutes at room temperature. The fluorescing solution, referred to as HRP-

DCF, was prepared by mixing 4.0 mL of DCFH with 7.0 mg of peroxidase from horseradish

(HRP, Type I, Sigma-Aldrich), brought to a final volume of 100 mL with pH 7.2 phosphate

buffer (7.35 mM KH2PO4, 17.6 mM Na2HPO4). The HRP-DCF solution was kept in amber

bottles on ice until needed.

Hydrogen peroxide standards were freshly prepared for the calibration of each assay.

A stock solution (1 mM H2O2, 30% wt ACS reagent, Sigma Aldrich) was prepared in

deionized water, used for standard dilutions ranging from μM to 0 μM H2O2. Standards

were stored in amber bottles on ice until needed. The calibration curve was fit to a second-

order polynomial trend, where most assay samples fell within the linear regime of the curve.

A sample calibration curve can be found in Figure 2.3. The limit of detection of the HRP-

DCF assay is 0.1 μM H2O2.

Figure 2.3– Calibration curve of hydrogen peroxide obtained using the HRP-DCF assay.

Concentrations of H2O2 standards are μM to 0 μM.

y = -6E+08x2 + 34455x - 0.016

R² = 0.9995

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0.45

0.00E+00 5.00E-06 1.00E-05 1.50E-05 2.00E-05 2.50E-05

Bla

nk

Co

rrec

ted

Flu

ore

scen

ce (

50

0 n

m)

at

52

3.9

nm

H2O2 Concentration (M)

20

Assay samples were prepared by adding 50 μL of the sample with 2.25 mL HRP-

DCF in Teflon-capped amber vials. Samples were briefly mixed before reacting at room

temperature in the dark for 30 minutes. The reaction was quenched by putting vials on ice for

30 seconds, before measuring the sample fluorescence. A spectrometer/fluorometer

(SpectroVis Plus, Vernier) was operated in fluorescence mode with a 500 nm excitation light

source, measuring sample spectra at 523.9 nm. Data were collected using Logger Pro

software (Version 3.8.2, Vernier).

2.2.3 Kinetic Modelling

A kinetic model of the copper/iron chemistry was created using the React modelling

program (version 1.2, Alchemy Software). A summary of reactions used in the model can be

found in Table 2.1. The model was set to run for 60 minutes using 1 second time steps. The

photolysis of acetaldehyde was simplified in the model to a reaction converting the aldehyde

directly to HO2 (reaction 2), with the rate constant selected to match the model to

experimental results. The model’s initial conditions were as follows: 1.0 mM H+ to match

experimental pH of 3, 0.235 mM O2 to match approximated dissolved oxygen at photolysis

sample temperature of 33 °C, and 55 M H2O. Initial acetaldehyde, copper(II), and iron(II)

concentrations were varied in order to study different experimental systems. The rate

constants are expressed in units of M-1

s-1

, where 8.8(-7) represents 8.8 x10-7

M-1

s-1

.

As will be discussed further in the following section, it was found that 1.6 μM H2O2

was generated upon the photolysis of an aqueous solution of 1 mM acetaldehyde using our

experimental set up. By running the model to represent a system containing 1 mM

acetaldehyde and no transition metals, the rate constant for the photolysis of acetaldehyde

(reaction 2) was adjusted so that the output produced 1.6 μM H2O2 after a one hour

simulation. The rate constant representing the conversion of acetaldehyde to HO2 was found

to be 8.8 x10-7

M-1

s-1

. This suggests that the photolytic lifetime of acetaldehyde is

approximately 13 days under the experimental conditions used in this study.

21

Table 2.1 – Reactions and rate constants included in kinetic model

Reaction Rate

Constant

(M-1

s-1

)

Reference

1 H2O H+ + OH

- 1.0(-14)

a

2 Aldehyde HO2 8.8(-7)b

Experimental

3 HO2 O2- + H

+ 2.1(-5)

a Bielski, 1978 [15]

4 HO2 + HO2 H2O2 + O2 8.6(5) Bielski, 1978 [15]

5 OH + HO2 H2O + O2 7.0(9) Sehested et al., 1968 [16]

6 OH + O2-

OH

- + O2

1.0(10) Sehested et al., 1968 [16]

7 OH + H2O2 H2O + HO2 2.7(7) Christensen et al., 1982 [17]

8 HO2 + O2- H2O2 + O2 + OH

- 1.0(8) Bielski, 1978 [15]

9 HO2 + Cu2+ Cu

+ + O2 + H

+ 1.0(8) Rabani et al., 1973 [18]

10 O2- + Cu

2+ Cu

+ + O2 8.0(9) Rabani et al., 1973 [18]

11 OH + Cu+

Cu

2+ + OH

- 3.0(9) Goldstein et al., 1992 [19]

12 O2 + Cu+ Cu

2+ + O2

- 4.6(5) Bjergbakke et al., 1976 [20]

13 H2O2 + Cu+ Cu

2+ + OH + OH

- 7.0(3) Berdnikov, 1973 [21]

14 HO2 + Cu+ + H

+ Cu

2+ + H2O2 3.5(9) Berdnikov, 1973 [21]

15 O2- + Cu

+ + 2H

+ Cu

2+ + H2O2 9.4(9) vonPiechowski et al., 1993 [11]

16 H2O2 + Fe2+ Fe

3+ + OH + OH

- 70 Christensen et al., 1993 [17]

17 H2O2 + Fe3+

Fe2+

+ HO2 + H+

2.0(-3) Walling and Goosen, 1973 [22]

18 O2- + Fe

2+ + 2H

+ Fe

3+ + H2O2 1.0(7) Rush and Bielski, 1985 [23]

19 O2- + Fe

3+

Fe

2+ + O2 1.5(8) Rush and Bielski, 1985 [23]

20 HO2 + Fe2+

+ H+ Fe

3+ + H2O2 1.2(6) Jayson et al., 1973 [24]

21 OH + Fe2+ Fe(OH)

2+ 4.6(8) Christensen&Sehested,1980[25]

R3 Fe3+

+ Cu+ Cu

2+ + Fe

2+ 1.3(7) Bjergbakke et al., 1976 [20]

a Denotes equilibrium constant

b Denotes first order rate constant in units of s

-1

2.3 Results and Discussion

2.3.1 Acetaldehyde photolysis

The photolysis of acetaldehyde samples produced detectable levels of hydrogen

peroxide, suggesting that it is a suitable source of aqueous-phase HOX radicals. Samples

containing less than 1 mM acetaldehyde produce peroxide levels near the method’s limit of

detection (0.1 μM) based on a 1-hour photolysis time, and were therefore not used in this

study. A summary of the concentration of hydrogen peroxide produced from a 1-hour

photolysis of various acetaldehyde concentrations can be found in Table 2.2.

22

Table 2.2 – Hydrogen peroxide production from photolyzed aldehyde samples

Concentration of

Acetaldehyde

Amount of H2O2

Produced (μM)

Standard

Deviation (μM)

0.10 mM 0.06 0.02

1.0 mM 1.6 0.52

3.0 mM 6.1 1.4

The photolytic production of hydrogen peroxide from acetaldehyde was linear as a

function of photolysis time, as seen in Figure 2.4. Based on these results, an experimental

photolysis time of 1 hour was selected for the remainder of the study; this is the shortest

experimental time that produces a signal that lies confidently above the method’s detection

limit and within the linear portion of the calibration curve. The growth in error bars as time

progressed was a result of placement of samples within the solar simulator. Samples that

were situated closest to the inlet of the cooling air contained slightly lower levels of peroxide

than those further away, however this effect only becomes significant at longer photolysis

times.

Figure 2.4 – Peroxide production from 3 hour photolysis of 3 mM acetaldehyde in solar

simulator

y = 0.0738x + 1.3677

R² = 0.9585

0

2

4

6

8

10

12

14

0 20 40 60 80 100 120 140

H2O

2 P

rod

uce

d (

μM

)

Photolysis Time (min)

23

2.3.2 Aldehyde and Copper Photolysis

The hydrogen peroxide concentrations of photolysis of samples containing both 1

mM acetaldehyde and 1mM copper were not found to be significantly different from samples

containing only 1 mM acetaldehyde, as seen in Figure 2.5. In other experiments, the

dependence of peroxide production on copper concentration was tested by preparing samples

containing 3 mM acetaldehyde with CuSO4 concentrations ranging from 0.1-3 mM. It was

found that the peroxide production of these samples was also the same as samples with only

acetaldehyde, and thus is independent of the copper concentration. For the remaining

experiments of this study, 1 mM of CuSO4 was used in all copper-containing samples.

Figure 2.5 – Hydrogen peroxide concentrations produced from 1-hour photolysis of samples

containing 1 mM acetaldehyde, 1 mM CuSO4, and/or 1 μM FeSO4

Although copper acts as a catalyst in the cycling of HOX and the production of

hydrogen peroxide, it does not produce significantly different amounts of peroxide than a

system with only HOX. It is expected that at very short times in the sample’s photolysis, the

reaction of copper(I) and HO2 is the dominant HOX reaction (Table 2.1, reaction 14),

resulting in the rapid production of H2O2 . However, as the amount of H2O2 begins to build up

0

2

4

6

8

10

12

14

16

18

H2O

2 C

on

cen

tra

tio

n (

μM

)

Aldehyde

Aldehyde+Copper

Aldehyde+Iron

Aldehyde+Copper+Iron

24

as the photolysis continues, the reaction of copper(I) with H2O2 (reaction 13) begins to

become more significant. This results in a loss of H2O2, thus slowing the overall production

rate of H2O2. This result is confirmed in the modeling of the system, where at very short

reaction times (on the order of milliseconds) samples containing copper produce

approximately 1.2 times more peroxide than those containing only acetaldehyde. Within the

full hour of photolysis, however, the model predicts that samples containing 0.1-3 mM

copper will produce 10-40 % less peroxide than samples containing only acetaldehyde, as

shown in Figure 2.6. This effect was not observed experimentally, suggesting that the model

over-predicts the loss of peroxide. Model predictions show that copper is in excess with

respect to HOX radicals by several orders of magnitude. Thus, over the range of

concentrations used experimentally, the copper always remains in excess, accounting for no

observed dependency of peroxide production on copper concentration.

Figure 2.6 – Modelled production of hydrogen peroxide in a system containing 1 mM

aldehyde with varying copper concentrations at 60 minutes of photolysis

2.3.3 Aldehyde and Iron Photolysis

The photolysis of samples containing aldehyde with 1 mM FeSO4 produced no

detectable levels of hydrogen peroxide. Under these conditions, the lifetime of hydrogen

25

peroxide in the presence of Fe(II) is 14 seconds (reaction 16). In order for the peroxide to

persist through the timescale of the experiment, concentrations of FeSO4 cannot exceed 1

μM. Under these conditions, the lifetime of H2O2 will be 240 minutes.

The photolysis of samples containing both acetaldehyde and 1 μM iron produced

hydrogen peroxide concentrations that were approximately one order of magnitude greater

than those containing no iron, seen in Figure 2.5. Unlike the copper system, the photolytic

generation of H2O2 does depend on the sample’s concentration of FeSO4. As the

concentration of Fe(II) increased from 0.01 μM to 1 μM, the concentration of peroxide

linearly increased, seen in Figure 2.7.

Figure 2.7 - Concentration of hydrogen peroxide produced from solutions containing 1 mM

acetaldehyde with iron(II) varied from 0.01 to 1.0 µM (note: no error bars are shown due to

single measurement for each experiment)

In order to better understand the chemistry of the iron system, samples were prepared

in which iron was added to the solutions both before and after photolysis. It was found that

samples photolyzed in the presence of iron produced 13 μM peroxide; this is an order of

magnitude more hydrogen peroxide produced than samples in which iron was added after

photolysis, which only produced 1.5 μM. The high levels of peroxide observed in samples

y = 7.9154x + 3.6917

R² = 0.9978

0

2

4

6

8

10

12

14

0.0 0.2 0.4 0.6 0.8 1.0 1.2

H2O

2 C

on

cen

tra

tio

n (

μM

)

Fe(II) Concentration (μM)

26

photolyzed in the presence of iron suggests that the formation of a complex that efficiently

releases HOX is occurring, as opposed to the iron interfering with the HRP-DCF solution in

the fluorescence analysis.

Based on iron’s high affinity to form complexes with organic ligands, it is likely that

there is a high degree of complexation within the photolysis samples. As the acetaldehyde

photodissociates, there will be a mixture of organic radicals and stable molecules available to

complex. Upon photolysis of the complex, an organic radical will be formed from the ligand.

This radical can go on to react with O2, ultimately forming HO2, as shown in R7-R9.[26]

This additional source of HO2 within the system may account for the increase in peroxide in

iron-containing samples. A more thorough discussion of iron-organic complexes can be

found in a later section of this chapter.

Fe(III)-L Fe(II) + L∙

L∙ + O2 O2-∙ + L

+

O2-∙ + H

+ HO2∙

The increase in peroxide may also be due to the production of OH radicals from an

iron-hydroxy complex. In an aqueous solution, iron(III) ions will be in equilibrium with their

complexed form, [Fe(III)(OH)(H2O)5]2+

. At the sample pH of 3 that is used in this study, it

has been shown that the majority of iron in the solution will be found in this complexed

form.[27, 28] Upon photolysis, OH radicals will be produced as shown below in R10.[27]

These radicals have the potential to further react with acetaldehyde to first generate a

carboxylic acid and HO2 radical, as shown in a later section in Figure 2.10, product a). The

extra production of HO2 would then lead to an increase in the amount of hydrogen peroxide

formed.

[Fe(III)(OH)(H2O)5]2+

+ H2O + hν [Fe(II)(H2O)6]2+

+ OH∙

Experiments in which aqueous solutions of iron at pH 3 were photolyzed with

benzoic acid, a common OH radical trap, have shown the production of OH radicals under

our experimental conditions.[29] As OH radicals are generated, they will rapidly oxidize

benzoic acid to form hydroxybenzoic acid. In this experiment, benzoic acid (≥99.5%, Sigma-

R7

R8

R9

R10

27

Aldrich) was added to a final concentration of 1 mM in solution with 1 µM iron(III).

Solutions were photolyzed for 1 hour using the solar simulator. The production of OH

radicals was monitored through measuring the formation of hydroxybenzoic acid using a

direct analysis in real time mass spectrometer (DART-MS). The observed formation of

hydroxybenzoic acid in the iron(III) samples confirms that there is production of OH radical.

Thus, there is potential that the OH radicals will react with excess acetaldehyde in the iron-

only samples, resulting in the production of hydrogen peroxide.

2.3.4 Aldehyde, Copper, and Iron

In samples containing acetaldehyde in the presence of both copper and iron, it is

found that the photolytic production of hydrogen peroxide is significantly less than samples

containing acetaldehyde with only iron. However, there is an increase of approximately two

times the concentration of H2O2 found in samples with no transition metals, seen in Figure

2.5. In these experiments with both metals, acetaldehyde and iron concentrations were held at

1 mM and 1 µM respectively, while copper concentrations ranged from 0.1 µM to 1.0 mM.

This allowed the study of the effect of changing the copper to iron ratio, Cu/Fe. As the

amount of copper increased within the system, it was found that the amount of hydrogen

peroxide produced decreased, as seen in Table 2.3.

Table 2.3 – Peroxide production in a system containing 1 mM acetaldehyde with varying

copper to iron ratios

Ratio of Cu/Fe Amount of H2O2

Produced (μM) Standard Deviation (μM)

0 14.8 2.0

0.1 13.1 2.3

1 13.4 2.6

1000 4.48 0.79

It is believed that the observed decrease in peroxide produced from samples

containing aldehyde and both transition metals compared to samples containing only iron is

due to a hindrance of the formation of the iron complexes. Both the organic and hydroxy

complexes contain iron in the +3 oxidation state, however the samples are prepared using

28

iron in the +2 oxidation state. This decrease in iron(III) could be due to either a slowed

production through the reaction of iron(II) and HO2, or an increase in iron(III) loss through

the coupling reaction of iron(III) and copper(I).

In the system containing only iron, iron(II) will react with HO2 (reaction 20),

generating iron(III). The iron(III) can then go on to form the complexes, photolytically

generating peroxide. In the presence of copper and iron however, the reaction of HO2 with

copper (reaction 9) is favoured over that with iron. Thus, as the amount of copper in the

system increases, the reaction of iron(II) and HO2 will become less dominant, slowing the

production of iron(III) and ultimately reducing the degree of complexation. This effect is

seen in the model simulations, where the reaction rate between iron(II) and HO2 decreases as

the amount of copper increases, as seen in Table 2.4.

Table 2.4 – Modelled reaction rates of iron(II) and HO2 as copper concentration increases

Copper/Iron ratio Rate of Reaction 20 (M/s)

0 Cu/Fe 3.5 x10-10

0.1 Cu/Fe 2.7 x10-10

1 Cu/Fe 1.5 x10-10

1000 Cu/Fe 5.3 x10-12

To evaluate the effect of the coupling reaction on the concentration of iron(III), the

model simulation was run both with and without reaction R3. As seen in Figure 2.8, the

concentration of iron(III) decreases as a function of increasing copper concentration when

the coupling reaction is included. However, the modelled Fe(III) concentrations from the

simulation without the coupling reaction predicts a more similar result to what was observed

experimentally. In this case, the model predicts that a system with only iron will have

approximately the same amount of iron(III) as a system with 0.1 and 1 Cu/Fe. This would

mean that a similar amount of iron-complexes will form in these three systems, producing

approximately the same amount of hydrogen peroxide. Consequently, it is not anticipated

that the coupling reaction of iron(III) with copper(I) is a significant reaction under these

experimental conditions. It is clear, however, that there is some form of coupling occurring

between the copper and iron redox cycles, as the modelled concentrations in the 0.1 and 1

29

Cu/Fe systems without coupling do not follow the same trend as the iron-only system

through the whole simulation.

Figure 2.8 – Modelled Fe(III) concentration in a system containing 1 mM aldehyde with

Cu/Fe ratios of 0 (1 μM Fe(II)), 0.1 (0.1 μM Cu(II) with 1 μM Fe(II)), 1 (1 μM Cu(II) with 1

μM Fe(II)), and 1000 (1 mM Cu(II) with 1 μM Fe(II)). Dashed lines represent simulations in

which the coupling reaction (R3) was removed from model.

Despite the fact that the presence of copper will reduce the extent of complexation of

iron with organic ligands (particularly in the system containing 1 mM copper(II) with 1 μM

iron(II), 1000 Cu/Fe), it is still anticipated that complexes will form in small amounts.

Although the production rate of iron(III) will be significantly slower in a system containing

copper, this reaction will still proceed, allowing some iron(III) to form, albeit to a much

lesser extent than the iron-only system. These complexes are still expected to photolytically

generate H2O2, accounting for the higher levels observed in this system as compared to the

aldehyde-only and copper containing systems.

30

2.3.5 Kinetic model results

The kinetic model was primarily used to determine the fate of each species in each

experimental system. Modelled concentrations of key species for each system studied can be

found in Figures 2.11-2.14 in the Appendix. The most heavily studied systems studied

through the model contained 1 mM acetaldehyde with 1 mM copper(II), and/or 1 μM iron(II)

at a solution pH of 3.

In a system containing 1 mM acetaldehyde and 1mM copper(II), it was found that the

dominant pathway of peroxide production was the reaction of copper(I) with the HO2 radical

(reaction 14), proceeding at a rate of 3.5 x10-10

M/s. The reaction of copper(I) with the

dissociated O2- radical (reaction 15) was the second largest production of hydrogen peroxide,

with a reaction rate of 2.0 x10-11

M/s. Finally, the self reaction of HO2 (reaction 4) and O2-

(reaction 8) were the least significant sources of hydrogen peroxide, proceeding at rates of

3.5 x10-17

and 8.4 x10-17

M/s respectively.

In a system containing 1 mM acetaldehyde with 1 μM iron(II), the model predicted

that the dominant pathway of hydrogen peroxide production is through the reaction of

iron(II) with HO2 (reaction 20), with a rate of 3.5 x10-10

M/s. The second pathway of

peroxide formation is from the reaction of iron(II) with O2- (reaction 18), proceeding at a rate

of 6.0 x10-11

M/s. Like in the copper-only system, the self reaction of HO2 (reaction 4) and

O2- (reaction 8) were the least significant sources of peroxide, with rates of 2.2 x10

-13 and 5.3

x10-13

M/s respectively. The model did not include any iron complex chemistry, and thus a

significant amount of the peroxide that was observed experimentally is unaccounted for

using this model.

In a system containing 1 mM acetaldehyde with both 1 mM copper(II) and 1 μM

iron(II), it was found that the dominant peroxide production is from the reaction of copper(I)

with HO2 (reaction 14), with a rate of 3.3 x10-10

M/s. The peroxide is then generated by the

reaction of copper(I) with O2- (reaction 15, 1.8 x10

-11 M/s), iron(II) with HO2 (reaction 20,

5.2 x10-12

M/s) iron(II) with O2- (reaction 18, 9.0 x10

-13 M/s), HO2

with O2

- (reaction 8, 8.0

x10-17

M/s), and the HO2 self reaction (reaction 4, 3.4 x10-17

M/s), given in order of

decreasing significance. In this system, the reaction of copper(I) with HO2 was more

31

significant than the reaction of copper(I) with iron(III) (reaction R3, with a rate of 5.6 x10-11

M/s), suggesting that the coupling chemistry under these conditions would not be as

significant as previously predicted.

2.3.6 Organic Iron Complexes

The unanticipated increase in peroxide generated in systems containing aldehyde and

iron warranted further investigation into iron-organic complexes. A series of solutions

containing 1 mM of an organic acid were prepared with and without 1µM iron(III), and

photolyzed in the solar simulator for 1 hour. The concentration of hydrogen peroxide

generated in each sample can be found in Figure 2.9.

Figure 2.9 – Generation of hydrogen peroxide from photolysis of organic acids both with

and without iron(III).

The photolysis of aqueous pyruvic acid both with and without iron produced

approximately 30 µM of hydrogen peroxide. Solutions of oxalic acid in the presence of iron

produced the greatest amount of peroxide, approximately 50 µM, while aqueous oxalic acid

without iron did not produce a significant amount of peroxide. This result is in agreement

0

10

20

30

40

50

60

70

H2O

2 C

on

cen

tra

tio

n (

µM

)

Pyruvic

Pyruvic + Iron(III)

Oxalic

Oxalic + Fe(III)

Glyoxylic

Glyoxylic + Fe(III)

Tricarballylic

Tricarballylic + Fe(III)

Citric

Citric + Fe(III)

Pinonic

Pinonic + Fe(III)

32

with previous studies that found oxalate complexes favourably with iron.[30, 31] There was

a significant increase in the amount of peroxide produced from the photolysis of glyoxylic

acid when iron was included, however the amount of peroxide was insignificant when

compared to that from oxalic acid with iron. The other organic acids tested in this study did

not produce significant amounts of peroxide, both with and without iron.

Acetaldehyde can also be oxidized in solution to form organic compounds that are

known to complex with iron, including oxalic acid, glyoxylic acid, and pyruvic acid.[26, 32]

A proposed oxidation mechanism can be found in Figure 2.10, with oxalic and glyoxylic

acids labelled with b and c respectively. Upon the photolysis of acetaldehyde, it is also likely

that several organic fragments are generated in solution. This complex mixture of organics

have the potential to form compounds that will favourably complex with iron(III) in solution.

Figure 2.10 – Proposed mechanism of oxidation of acetaldehyde to form a) acetic acid,

which is further oxidized to b) glyoxylic acid, and ultimately forming c) oxalic acid, which is

shown to complex with iron(III) to generate peroxide.

b) c)

a)

33

2.4 Conclusions and Atmospheric Implications

The redox cycling of HOX radicals in the presence of transition metals was evaluated

by measuring hydrogen peroxide production in photolyzed aqueous solutions. In particular,

the changes in peroxide production in a system containing both copper and iron

simultaneously was used to evaluate if a proposed coupling of the two metal redox cycles

was occurring in solution.

In aqueous solutions of a HOX source with transition metals, the amounts of hydrogen

peroxide generated after one hour in the presence and absence of only copper ions are not

significantly different. This result was anticipated, showing that the observed peroxide

production is a result of HOX radicals participating in the transition metal catalyzed redox

cycle.

In the presence of only iron, an order of magnitude more peroxide is generated than

in its absence, suggesting that the iron ions are complexed in solution. The complexes could

be the result of iron interacting with organics in the solution, such as oxalate, forming an

iron-organic complex, or due to the speciation of iron under acidic aqueous conditions,

forming an iron-hydroxy complex. Both types of complex will efficiently lead to peroxide

production after photolysis.

When copper and iron are present simultaneously in solution, there is a significant

increase in peroxide when compared to a system without transition metals. The amount of

peroxide, however, is less than that in a system containing only iron. It is believed that the

presence of copper will slow the conversion of iron(II) to iron(III), thus reducing the amount

of iron(III) available to form complexes.

The efficacy of the copper-iron coupling suggested by Mao et al. in their recent

modelling paper could not be demonstrated in a simple system containing one organic

species. In an environmental system, it is likely that there will be a complex mixture of many

different, highly concentrated organic species. Thus, it is possible that in the atmosphere the

degree of complexation with iron may be comparable or higher than that observed in this

study. As a result, it would be necessary to conduct similar experiments in more

34

atmospherically relevant samples in order to evaluate if this coupling could be significant in

the atmosphere.

35

2.5 References

[1] K.V. Desboeufs, A. Sofikitis, R. Losno, J.L. Colin, and P. Ausset, Dissolution and

solubility of trace metals from natural and anthropogenic aerosol particulate matter, in

Chemosphere, 2005, pp. 195-203.

[2] L. Deguillaume, M. Leriche, K. Desboeufs, G. Mailhot, C. George, and N. Chaumerliac,

Transition metals in atmospheric liquid phases: Sources, reactivity, and sensitive

parameters, in Chemical Reviews, 2005, pp. 3388-3431.

[3] R.L. Siefert, S.O. Pehkonen, Y. Erel, and M.R. Hoffmann, Iron photochemistry of

aqueous suspensions of ambient aerosol with added organic-acids, in Geochimica Et

Cosmochimica Acta, 1994, pp. 3271-3279.

[4] D.J. Jacob, Heterogeneous chemistry and tropospheric ozone, in Atmospheric

Environment, 2000, pp. 2131-2159.

[5] C. Weller, S. Horn, and H. Herrmann, Photolysis of Fe(III) carboxylato complexes:

Fe(II) quantum yields and reaction mechanisms, in Journal of Photochemistry and

Photobiology a-Chemistry, 2013, pp. 24-36.

[6] B.C. Faust, and J. Hoigne, Photolysis of Fe(III)-hydroxy complexes as sources of OH

radicals in clouds, fog, and rain, in Atmospheric Environment Part a-General Topics,

1990, pp. 79-89.

[7] C.J. Weschler, M.L. Mandich, and T.E. Graedel, Speciation, photosensitivity, and

reactions of transition-metal ions in atmospheric droplets, in Journal of Geophysical

Research-Atmospheres, 1986, pp. 5189-5204.

[8] C. Anastasio, B.C. Faust, and J.M. Allen, Aqueous-phase photochemical formation of

hydrogen-peroxide in authentic cloud waters, in Journal of Geophysical Research-

Atmospheres, 1994, pp. 8231-8248.

[9] D.L. Sedlak, J. Hoigne, M.M. David, R.N. Colvile, E. Seyffer, K. Acker, W. Wiepercht,

J.A. Lind, and S. Fuzzi, The cloudwater chemistry of iron and copper at Great Dun Fell,

UK, in Atmospheric Environment, 1997, pp. 2515-2526.

[10] L. Deguillaume, M. Leriche, A. Monod, and N. Chaumerliac, The role of transition metal

ions on HOx radicals in clouds: a numerical evaluation of its impact on multiphase

chemistry, in Atmospheric Chemistry and Physics, 2004, pp. 95-110.

[11] M. Vonpiechowski, T. Nauser, J. Hoigne, and R.E. Buhler, O2- decay catalyzed by Cu2+

and Cu+ ions in aqueous-solutions - A pulse-radiolysis study for atmospheric chemistry,

in Berichte Der Bunsen-Gesellschaft-Physical Chemistry Chemical Physics, 1993, pp.

762-771.

[12] J. Mao, S. Fan, D.J. Jacob, and K.R. Travis, Radical loss in the atmosphere from Cu-Fe

redox coupling in aerosols, in Atmospheric Chemistry and Physics, 2013, pp. 509-519.

[13] P. Warneck, and G.K. Moortgat, Quantum yields and photodissociation coefficients of

acetaldehyde in the troposphere, in Atmospheric Environment, 2012, pp. 153-163.

[14] D.A. Jackson, and S.A. Mabury, Environmental properties of pentafluorosulfanyl

compounds: Physical properties and photodegradation, in Environmental Toxicology

and Chemistry, 2009, pp. 1866-1873.

[15] B.H.J. Bielski, Reevaluation of the spectral and kinetic properties of HO2 and O2- free

radicals, in Photochemistry and Photobiology, Blackwell Publishing Ltd, 1978, pp. 645-

649.

[16] K. Sehested, O.L. Rasmussen, and H. Fricke, Rate constants of OH with HO2,O2-, and

H2O2+ from hydrogen peroxide formation in pulse-irradiated oxygenated water, in The

Journal of Physical Chemistry, American Chemical Society, 1968, pp. 626-631.

36

[17] H. Christensen, K. Sehested, and H. Corfitzen, Reactions of hydroxyl radicals with

hydrogen peroxide at ambient and elevated temperatures, in The Journal of Physical

Chemistry, American Chemical Society, 1982, pp. 1588-1590.

[18] J. Rabani, D. Klug-Roth, and J. Lilie, Pulse radiolytic investigations of the catalyzed

disproportionation of peroxy radicals. Aqueous cupric ions, in The Journal of Physical

Chemistry, American Chemical Society, 1973, pp. 1169-1175.

[19] S. Goldstein, G. Czapski, H. Cohen, and D. Meyerstein, Deamination of β-alanine

induced by hydroxyl radicals and monovalent copper ions. A pulse radiolysis study, in

Inorganica Chimica Acta, 1992, pp. 87-93.

[20] E. Bjergbakke, K. Sehested, and O.L. Rasmussen, The Reaction Mechanism and Rate

Constants in the Radiolysis of Fe2+ and Cu2+ Solutions, in Radiation Research,

Radiation Research Society, 1976, pp. 433-442.

[21] V. Berdnikov, Catalytic activity of the hydrated copper ion in the decomposition of

hydrogen peroxide, in Russian Journal of Physical Chemistry, 1973, pp. 1060-1162.

[22] C. Walling, and A. Goosen, Mechanism of the ferric ion catalyzed decomposition of

hydrogen peroxide. Effect of organic substrates, in Journal of the American Chemical

Society, American Chemical Society, 1973, pp. 2987-2991.

[23] J.D. Rush, and B.H.J. Bielski, Pulse radiolytic studies of the reaction of

perhydroxyl/superoxide O2- with iron(II)/iron(III) ions. The reactivity of HO2/O2- with

ferric ions and its implication on the occurrence of the Haber-Weiss reaction, in The

Journal of Physical Chemistry, American Chemical Society, 1985, pp. 5062-5066.

[24] G.G. Jayson, B.J. Parsons, and A.J. Swallow, Oxidation of ferrous ions by perhydroxyl

radicals, in Journal of the Chemical Society, Faraday Transactions 1: Physical

Chemistry in Condensed Phases, The Royal Society of Chemistry, 1973, pp. 236-242.

[25] H. Christensen, and K. Sehested, Pulse radiolysis at high temperatures and high

pressures, in Radiation Physics and Chemistry (1977), 1980, pp. 183-186.

[26] Y.G. Zuo, and J. Hoigne, Evidence for photochemical formation of H2O2 and oxidation

of SO2 in authentic fog water, in Science, 1993, pp. 71-73.

[27] G.S. Zhuang, Z. Yi, R.A. Duce, and P.R. Brown, Link between iron and sulfur cycles

suggested by detection of Fe(II) in remote marine aerosols, in Nature, 1992, pp. 537-539.

[28] W. Feng, and D. Nansheng, Photochemistry of hydrolytic iron (III) species and

photoinduced degradation of organic compounds. A minireview, in Chemosphere, 2000,

pp. 1137-1147.

[29] X.L. Zhou, and K. Mopper, Determination of photochemically produced hydroxyl

radicals in seawater and fresh-water, in Marine Chemistry, 1990, pp. 71-88.

[30] Y.G. Zuo, and J. Hoigne, Formation of hydrogen-peroxide and depletion of oxalic-acid

in atmospheric water by photolysis of iron(III) oxalato complexes, in Environmental

Science & Technology, 1992, pp. 1014-1022.

[31] C. Weller, S. Horn, and H. Herrmann, Effects of Fe(III)-concentration, speciation,

excitation-wavelength and light intensity on the quantum yield of iron(III)-oxalato

complex photolysis, in Journal of Photochemistry and Photobiology a-Chemistry, 2013,

pp. 41-49.

[32] Y.G. Zuo, and J. Hoigne, Photochemical decomposition of oxalic, glyoxalic and pyruvic-

acid catalyzed by iron in atmospheric waters, in Atmospheric Environment, 1994, pp.

1231-1239.

37

2.6 Appendix

Figure 2.11 – Kinetic Model output of key species from system containing 1 mM

acetaldehyde

Figure 2.12 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 mM copper(II)

a)

b)

38

Figure 2.13 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 μM iron(II)

a)

b)

39

Figure 2.14 – Kinetic model output of key a) HOX and b) transition metal species from

system containing 1 mM acetaldehyde with 1 mM copper(II) and 1 μM iron (II)

a)

b)

40

Chapter 3

Yields and Stability of Peroxides from Secondary Organic Aerosol

Formation

3.1 Introduction

Aerosol particles have been the focus of much research in recent years due to their

impacts on both climate and human health. Secondary organic aerosol (SOA) is made of

particles that are formed in the atmosphere by the processing of primary gas-phase organic

emissions to form compounds of lower volatility. Depending on the environmental

conditions, source of the precursor species, and degree of atmospheric processing, the

particles formed may contain not only organic molecules such as organic acids, peroxides,

and other functionalized species, but may also become internally mixed with inorganics,

including metals.

Of particular importance, peroxides can be found in the gaseous and particle

phases.[1] Peroxides that are in the particle phase have the potential to be carried much

deeper into the respiratory system upon inhalation than those in the gaseous phase. As a

result of their high solubility, gas-phase peroxides tend to be lost to fluids in the upper

regions of the respiratory tract.[2, 3] It has long been suggested that particles will have

increasingly more negative effects on health the deeper they penetrate into the respiratory

system, as seen in Figure 3.1.[3, 4] Studies have shown that ultrafine particles are more

detrimental to health as a result of their deposition locations among other factors.[5, 6]

Figure 3.1 – Deposition location of particles within the body as a function of particle size.[7]

41

As the particles bring peroxides deeper within the body, the natural balance of

oxidants and antioxidants is disrupted, causing an excess of oxidizing species.[8] This

imbalance results in oxidative stress within the body, leading to health complications such as

exacerbation of asthma, pneumonia, DNA damage causing cancer, and extreme cases

resulting in mortality.[6, 9]

Many cleaning products and air fresheners that are used indoors contain organic

precursors to secondary organic aerosols known as terpenes. One of the most commonly

detected volatile organic compound (VOC) indoors is limonene.[10] Limonene is primarily

found as a green solvent in household cleaners, a component of wax finishing products, and

the active ingredient of air fresheners, but can also be found in the peels of many citrus

fruits.[11, 12] Another prevalent VOC is α-pinene, which is produced through both biogenic

and anthropogenic means. Like limonene, α-pinene is frequently used as a component in

household cleaners and air fresheners, however it may also be emitted by wood products

used indoors such as furniture.[13-15] Past studies have measured typical indoor terpene

mixing ratios of up to 500 ppb, with the detected levels being heavily dependent on terpene-

containing product use and room ventilation.[14] In a recent study by Singer et al., it was

found that limonene levels reached as high as 1400 ppb following the use of an orange-oil

based cleaning product, resulting in a mass loading of SOA of up the 300 μg/m3 in the

presence of ozone.[16] In another study by Salthammer, terpene levels reached 900 ppb

following the use of an oil treatment on wood indoors.[14]

Due to the limited levels of sunlight indoors, the formation of secondary organic

aerosols is expected to be primarily through the ozonolysis of alkenes.[13, 17] The main

source of indoor ozone is through the transport of outdoor air through cracks and imperfectly

sealed buildings, and as a result typical indoor concentrations will be heavily influenced by

the environment outside.[13, 18] It has been reported that indoor concentrations of ozone are

typically 30-70 % of those measured outdoors.[13, 18] Additional sources of ozone indoors

include office and household devices, such as photocopiers and laser printers.[12] A recent

consumer trend involves the use of ozone generators indoors in order to “purify” the air.

While the release of ozone from these generators is reported by manufacturers to be within a

42

safe human exposure limit (less than 100 ppb), ozone concentrations of 2500 ppb have been

measured in a room equipped with a generator.[13, 18]

Figure 3.2 – Mechanism of limonene ozonolysis, showing select oxidation peroxide-related

products (adapted from Walser et al, 2008).[19]

A general reaction mechanism for the ozonolysis of terpenes, specifically the

oxidation of limonene, can be found in Figure 3.2.[19] The reaction between ozone and

terpenes is initiated by the addition of the ozone molecule across the double bond of the

organic compound, forming a primary ozonide. The ozonide molecule will decompose to

form a carbonyl and an energetic radical called a Criegee intermediate. There are several

fates of this intermediate, including the formation of peroxides (as shown in Figure 3.2),

Primary

Ozonide

Criegee

Intermediate

Peroxy

radical

Organic

hydroperoxide

43

organic acids, aldehydes, and ketones through the stabilization of the radical.[18] The

oxidized organic products of this reaction will have lower vapour pressures than the initial

terpene, causing the products to condense and form a SOA particle.

The association of hydroperoxides with secondary organic aerosol has been

investigated in several past studies.[12, 13, 15, 20-25] In a study conducted by Li et al., the

concentrations of total organic and hydrogen peroxides formed during the ozonolysis of

limonene were measured in both the gas and particle phases simultaneously. It was estimated

that 0.91 ±0.42 ppb of peroxide was generated from the reaction at limonene and ozone

concentrations relevant to indoor conditions.[13] A study by Docherty et al. measured the

yields of peroxides in SOA formed through the ozonolysis of α- and β-pinene, Δ-3carene, and

sabinene, reporting yields of 47-85 % of the SOA mass.[20] A series of studies conducted by

Chen and Hopke measured peroxides associated with particles formed using α-pinene,

limonene, and linalool VOC precursors.[12, 15, 24] It was found that the peroxide yields

were 1.8-26 x10-10

mole peroxide/μg SOA. The stability of the peroxides on the particles was

evaluated for 24 hours at room temperature for linalool and limonene SOA and in a

refrigerator for limonene SOA, reporting losses of 15-69% of peroxides at room temperature

and 10% in the refrigerator. A study by Wang et al. investigated peroxide associate with

SOA formed through the oxidation of α- and β-pinene and toluene precursors, reporting

yields of 2.7 x10-11

, 6.2 x10-11

mole peroxide/μg SOA for α- and β-pinene.[21] No peroxides

were detected for the fresh measurement of toluene SOA. The stability of the peroxides both

on-particle and in an aqueous solution were also measured, finding an exponential decay of

peroxides on-particle, and an increase in peroxide within the first 20 hours after extraction,

followed by a period of stability in aqueous solution. Mertes et al. generated α-pinene SOA,

measuring peroxide yields of 12-34 % of the SOA mass.[23] A study by Bateman et al.

measured the yields of peroxides from limonene SOA, reporting a value of 2% in terms of

the moles of SOA collected.[22] The stability of the peroxides under photolytic conditions

were also evaluated, finding that there was no significant change in peroxide levels following

14 hours of photolysis. A study by Mutzel et al. attempted to improve the quantification of

peroxides in SOA by measuring peroxides in α-pinene SOA, however no yields were

reported in terms of the amount of SOA analyzed.[25] In all of the studies mentioned,

different techniques were used to measure and quantify the peroxide levels within the SOA.

44

As will be discussed in a later section, the techniques have different sensitivities to hydro and

organic peroxides, accounting for some of the widespread variation in the reported peroxide

yields.

The purpose of this study is to investigate the yields and stability of peroxides in

secondary organic aerosols in order to develop a more thorough understanding of their

potential contribution to oxidation chemistry in the atmosphere and indoor environments.

The analysis of the peroxides in this study is done using the horseradish peroxidase-

dichlorofluorescien (HRP-DCF) assay as described in the previous chapter, with sensitivity

to primarily hydrogen peroxide. Aerosols are generated through the ozonolysis of α-pinene

and limonene using both a flow tube apparatus and environmental chamber in order to

evaluate the effect of aerosol mass loading on peroxide yields. The stability of peroxides will

be measured both on-filter and in aqueous solution. This will allow us to gain a sense of the

peroxide lifetimes in particles that remain dry upon emission into the atmosphere or indoors,

and in particles that take up water to form either aqueous aerosol or cloud droplets. The

photolytic stability of the peroxides in aqueous solution is evaluated using a solar simulator

to represent the outdoor environment, while a fluorescent light source was used to evaluate

the stability of the peroxides indoors both on a dry particle and in an aqueous solution.

Finally, the stability of the peroxides on-filter was measured in a -20 °C freezer in an attempt

to validate the common protocol within the community of freezing filter samples prior to

analysis. The results of the thermal stability experiments will be used to validate the peroxide

yields obtained in previous studies. To our knowledge, no previous study has observed the

photolytic stability of peroxides in any aerosol under indoor conditions, or in α-pinene SOA

outdoors.

The results of this project will lead to a more thorough understanding of the peroxides

associated with aerosol. Knowledge of the peroxide lifetimes within the particle will make it

possible to evaluate the length of time of which a particle will remain a health risk due to

oxidative stress, as opposed to solely being a respiratory risk. Understanding the nature

peroxides within particles will also improve the understanding of oxidation chemistry that

will occur in particulate and aqueous phases.

45

3.2 Methods

3.2.1 SOA Collection using Flow Tube

Secondary organic aerosol samples were collected using a flow tube setup with a

headspace bubbler providing the organic precursor, as seen in Figure 3.3. The flow tube was

used to collect a high mass of particles that was required to perform the HRP-DCF assay in a

short period of time, ensuring the freshness of the SOA. The flow tube has dimensions of 60

cm in length, 3.74 cm in inner radius, with a volume of 2.64 L. During SOA collection, the

flow tube was fully covered with foil to prevent exposure to light. All flows are reported as

volumetric measurements and were controlled by mass flow controllers (MFC) which were

calibrated using a Gilibrator air flow calibrator (Gilibrator-2, Sensidyne Instrumentation).

Ozone was generated by flowing air over a mercury lamp, with a flow rate of 150 ccm (MFC

1). Nitrogen was passed through a glass bubbler containing the VOC precursor ((-)-α-pinene,

≥99%, Sigma-Aldrich ; (R)-(+)-Limonene, 97%, Sigma-Aldrich) at a rate of 10 ccm (MFC

2). The bubbler was chilled at -7 °C for α-pinene experiments, and 2 °C for limonene

experiments. A carrier flow of clean air was added to the VOC flow after the bubbler, at 25

ccm (MFC 3). An excess dilution flow of 1500 ccm of clean air was incorporated into the

flows through a needle valve, with an open vent in the line before the air mixed with the

ozone flow. This excess flow is used to supply extra air to the flow tube in the case of

fluctuations in the other flows, ensuring that the total flow through the flow tube remains 1

lpm. The ozone and VOC flows were introduced to the top of the flow tube using a custom-

made stainless steel T-fitting that prevented mixing until the flows reached the body of the

flow tube. The flow rate through the flow tube was held at 1000 ccm using a pump (MFC 4).

Aerosol samples were collected on supported polytetrafluoroethylene (PTFE) filters (Zefluor,

Pall Life Sciences, 47 mm diameter, .0 μm pore size). Filter samples for stability and

fluorescent light photolysis experiments were collected for 2 hours, with approximately 500

μg of SOA generated. Filter samples for photolysis experiments using the solar simulator

were collected overnight (approximately 16 hours), yielding approximately 4.5 mg of SOA.

46

Figure 3.3 – Simplified schematic of flow tube set up for SOA generation and collection.

The amount of ozone entering the flow tube was periodically measured using an

ozone analyzer (model 49C, Thermo Environmental Instruments Inc.). Approximately 1700

ppb of ozone was introduced to the flow tube, while an excess of 870 ppb of ozone was

detected at the exit of the flow tube. Based on the consumption of ozone, it is estimated that

approximately 830 ppb of α-pinene and limonene are consumed within the flow tube. The

mixing ratios of the VOCs introduced into the flow tube were not measured. The mass

loadings of SOA within the flow tube were 3500 and 5000 μg/m3 for α-pinene and limonene

experiments, respectively.

3.2.2 SOA Collection using Environmental Chamber

An environmental chamber was used to collect secondary organic aerosol samples to

evaluate peroxide yields at lower mass loads (see Figure 3.4). The chamber is a 1 m3 teflon

bag supported by a Teflon-coated frame. The bag is externally surrounded by stainless steel

panels, preventing exposure to outside light. As in the flow tube set up, all chamber flows are

reported as volumetric measurements and were controlled by mass flow controllers. Ozone

47

was generated by passing air through a mercury lamp at a flow of 6 lpm (MFC 5). A dilution

flow of air of 7 lpm (MFC 4) is mixed with the ozone flow prior to being introduced into the

chamber through a stainless steel port. Limonene ((R)-(+)-Limonene, 97%, Sigma-Aldrich)

was introduced to the chamber using a 10 ccm flow of nitrogen through a headspace bubbler

chilled at 5 °C (MFC 1). α-pinene was introduced to the chamber through a 12 ccm flow

from a custom cylinder with a certified concentration (3 0 ppm ± 0 % α-pinene in nitrogen,

Air Liquide). The VOC flows were carried with a dilution flow of air of 500 ccm (MFC 2),

meeting an additional dilution flow of 2.2 lpm (MFC 3) before entering the chamber through

a stainless steel port. The flow rate through the chamber was held at 15 lpm using a pump

(MFC 6). Aerosol samples were collected on supported PTFE filters (Zefluor, Pall Life

Sciences, 47 mm diameter, .0 μm pore size) for hours, collecting an average of 185 μg

and 900 μg of α-pinene and limonene SOA respectively.

Figure 3.4 – Simplified schematic of environmental chamber set up for SOA generation and

collection.

The ozone levels within the chamber were periodically measured using an ozone

analyzer during collection. Approximately 320 ppb of ozone was introduced to the chamber,

while excesses of 85 ppb and 140 ppb of ozone were measured at the exit of the chamber

during α-pinene and limonene experiments respectively. A proton-transfer mass spectrometer

48

(PTR-MS, Ionicon Analytik GmbH) was used to measure the mixing ratio of limonene

within the chamber; it was found that 250 ppb of limonene was introduced into the chamber

using the bubbler. The flow of α-pinene used from the certified cylinder resulted in a mixing

ratio of 250 ppb within the chamber. The mass loadings of SOA within the chamber were

100 μg/m3 for α-pinene experiments, and 500 μg/m

3 for limonene experiments, considerably

lower than in the flow tube.

3.2.3 SOA Storage and Extraction

Following collection, SOA samples for on-filter stability testing were weighed and

immediately placed in a sealed plastic filter holder (Analyslide Petri Dish, Pall Life

Sciences). Samples were either placed in a dark cupboard at room temperature, or placed in a

sealed plastic bag and into a dark freezer at -20 °C. Filter samples were left for 1 to 18 days

before being reweighed and extracted for peroxide analysis. Frozen samples were brought to

room temperature before being removed from the filter holder to prevent the condensation of

contaminants onto the filter surface. SOA samples for in-solution stability or photolysis

experiments were weighed and extracted immediately following collection.

Filter samples were placed in a foil-covered Teflon bottle with deionized water (15

mL for stability samples; the volume for photolysis was varied to achieve SOA concentration

of 1 mM in solution, assuming the molecular weight of SOA is 200 g/mole[22]) and placed

on a shake table for 15 minutes at 420 rpm. Following extraction, an aliquot of each sample

was immediately prepared for analysis using the HRP-DCF assay, as described in Chapter 2.

Extracted solutions were covered with parafilm and foil, and stored in the dark at room

temperature to evaluate in-solution peroxide stability.

3.2.4 Photolysis in Solar Simulator

Following extraction, 1 mM SOA samples (prepared under the assumption that the

molecular weight of SOA is 200 g/mole) were transferred to custom-made quartz sample

holders. A volume of aqueous iron(III) chloride was added to some SOA samples for a final

49

concentration of 1 μM. Hydrogen peroxide samples were also prepared with and without iron

as photolysis standards. An equivalent set of all samples was prepared and stored in the dark

as a control. A xenon arc lamp in the Suntest CPS solar simulator was used as the light

source, as described previously in Chapter 2. Samples were photolyzed for 3 hours, with

aliquots of each sample being taken every hour for peroxide analysis.

3.2.5 Indoor Photolysis using Fluorescent Lights

Photolysis experiments using fluorescent lights in the laboratory were conducted with

the SOA sample both on filter and in solution. For on filter photolysis experiments, filter

samples were immediately placed in a pyrex covered petri dish after collection. In solution

samples were extracted in 15 mL of deionized water before being transferred to the petri

dish. Samples were positioned on the bench top under the laboratory’s fluorescent lights for

7 hours at room temperature. Equivalent samples were prepared and placed in a dark

cupboard for 7 hours to serve as a dark blank. Following photolysis, on filter samples were

extracted in water for 15 minutes. Solutions were then analyzed for peroxide content using

the HRP-DCF assay.

3.3 Results and Discussion

3.3.1 Peroxide Yields in SOA

Peroxide yields from “fresh” SOA samples are considered to be the measurements

made immediately following the 15 minute aqueous extraction of an SOA sample that was

promptly extracted after collection. A summary of fresh peroxide yields can be found in

Table 3.1, showing values reported by both this study and previously published studies. The

yields have been reported in three ways: 1) % mole, moles of peroxides/moles of SOA

x100%, where the molecular weight of SOA is assumed to be 200 g/mole; 2) % mass, mass

of peroxides/mass of SOA collected x100%, where molecular weight of peroxides is

assumed to be 34 g/mole; and 3) normalized yield, moles of peroxides/mass of SOA

collected. Yields will be reported as % mole for the remainder of these results.

50

Table 3.1 – Review of peroxide yields from fresh SOA samples reported in past studies

α-Pinene Detection method Yield (% mole) Yield (% mass) Normalized (mole/ug)

This study HRP-DCF 4.2 (± 0.6) 0.7 (± 0.1) 2.1 x10-10

(± 3.2 x10-11

)

Docherty et al.,

2005[20] Iodide -- 47 (± 12) --

Chen and Hopke,

2009[15] HRP-DCF -- -- 1.8 x10

-10 (± 7.9 x10

-11)

Wang et al.,

2011[21] HRP-PHOPAA

a -- -- 2.7 x10

-11 (± 1.1 x10

-11)

Mertes et al.,

2012 Iodide -- 34 (± 4) --

a-PHOPAA is para-hydroxyphenyl-acetic acid

The yields of peroxides from α-pinene and limonene SOA only comprise

approximately 4% of the total SOA molecular mass. These values are comparable to the

results reported in the series of studies conducted by Chen and Hopke, which also used the

HRP-DCF assay to measure peroxides from SOA formed through α-pinene and limonene

ozonolysis.[12, 15] The studies by both Docherty et al. and Mertes et al. used the iodide

technique to quantify their peroxides.[20, 23] This technique is sensitive to both hydro and

organic peroxides, and therefore the yields reported by these studies are considered to be the

total peroxide content of the SOA. As the HRP-DCF assay is largely sensitive to only

hydrogen peroxide, the difference in sensitivities between the techniques is suspected to

account for the lower yields reported in this study. The study by Bateman et al. also used the

iodide method to quantify peroxides, however hydrogen peroxide was used to calibrate their

fluorescence response.[22] The other studies using the iodide method calibrated the assay

using benzoyl peroxide, an organic peroxide that is expected to have a greater sensitivity to

the assay than hydrogen peroxide. This is likely the cause of their lower reported yields when

Limonene Detection method Yield (% mole) Yield (% mole) Normalized (mol/ug)

This study HRP-DCF 3.8 (± 0.7) 0.6 (±0.1) 1.9 x10-10

(± 3.5 x10-11

)

Chen and Hopke,

2010[12] HRP-DCF -- -- 1.6 x10

-10 (± 1.2 x10

-11)

Bateman et al.,

2011 Iodide 2 -- ∼1 x10

-10

51

compared to other studies using the iodide technique. Finally, the study by Wang et al.

reported yields that were approximately an order of magnitude lower than the α-pinene yields

in this study. This study also used horseradish peroxidase as their catalyst, however they used

a HPLC technique to speciate the peroxides, and para-hydroxyphenyl-acetic acid was used in

the place of dichlorofluoroscein as their fluorescing agent. Unfortunately, the origin of the

discrepancies between the values is not known.

As seen in Chapter 1, the sensitivity of the HRP-DCF assay heavily favours hydrogen

peroxide, although small sensitivity to simple organic peroxides exists. For this reason, it is

believed that the peroxides detected arising from the SOA are primarily hydrogen peroxide

with the possibility of having small amounts of other organic hydroperoxide species present.

That being said, a control experiment was conducted in which gas phase hydrogen peroxide

was passed through a Teflon filter. Upon extraction in water, it was found that no measurable

amount of hydrogen peroxide had stuck to the filter. This result suggests that the peroxides

measured in this work were not in the form of hydrogen peroxide when present in the aerosol

particles and on the filter, but rather converted to hydrogen peroxide when the filter materials

were put into solution for the assay. Indeed, hydrogen peroxide is so volatile that one would

not expect it to significantly partition to particles.

While we cannot confidently identify the form of the peroxides on the filters, it is

likely that they are in some form of small organic hyhdroperoxides. In particular, a general

class of compounds, the α-hydroxyhydroperoxides, exist in equilibrium with hydrogen

peroxide and organics containing carbonyl functional groups, especially aldehydes.[26] Their

abundance is much higher in the elevated precursor concentrations present in aerosol than in

dilute water solutions, i.e. it is expected that they decompose to form hydrogen peroxide

when dissolved in water. The chemistry of this class of organic peroxides and other classes is

not well enough known to be confident that this is the class of molecules that ultimately

gives rise to the signals observed.

52

3.3.2 Effect of SOA Mass Loading

The yields of peroxides in SOA generated using the chamber set up can be found in

Figure 3.5a. The mass loadings of SOA in the chamber were approximately 35 times lower

than those in the flow tube. The fresh yields of peroxides in chamber SOA were found to be

.9 % mole for α-pinene at 5.2 % mole for limonene. These yields are similar to those from

SOA collected using the flow tube (4.2 % mole and 3.8 % mole respectively). The yields

from chamber SOA that was stored at room temperature for 24 hours were found to be 1.1 %

mole and .4 % mole for α-pinene and limonene experiments; this corresponds to a loss of

peroxides of approximately 50%. Similar trends were observed in SOA collected using the

flow tube, as will be discussed in the following section. The results demonstrate that mass

loading does not have a large effect on peroxide yields in SOA, and as a result the remainder

of the experiments will use SOA generated through the flow tube set up.

3.3.3 On-Filter Stability

The yields of peroxide from α-pinene and limonene SOA can be found in Figure 3.5

as a function of filter age before extraction. At room temperature, a loss of peroxide was

observed as the filter age increased for both α-pinene and limonene SOA. Within 24 hours at

room temperature, a decay of approximately 50% of the total peroxide yield was observed.

An insignificant amount of peroxide remained in both SOA types after 7 days.

Unlike the stability at room temperature, it was found that the peroxide yields were

stable when the filters were stored in the freezer at -20 °C. No significant change in peroxide

yields was observed in α-pinene and limonene SOA filters that were stored in the freezer for

up to 7 days.

53

Figure 3.5 – Stability of peroxides in SOA stored on-filter at (a) room temperature and (b) in

freezer at -20 °C (note: points without error bars represent experiments where a single

measurement was performed)

The decay of peroxides from the SOA samples at room temperature on filter suggests

that the species are lost either through volatilization or through chemical reactions. If the

peroxides begin as small organic peroxides, they could either volatilize directly from the

0.0

1.0

2.0

3.0

4.0

5.0

6.0

0 50 100 150 200

Yie

ld (

% m

ole

)

Filter age (hrs)

α-pinene - flow tube

limonene - flow tube

α-pinene - chamber

limonene - chamber

0.0

1.0

2.0

3.0

4.0

5.0

6.0

0 50 100 150 200

Yie

ld (

% m

ole

)

Filter age (hrs)

α-pinene

limonene

(a)

(b)

54

particle, or as hydrogen peroxide following a conversion reaction within the particle phase.

Although organic peroxides are volatile, hydrogen peroxide will volatilize at a faster rate

given its low molecular weight. As a specific example, if the peroxides are in the form of α-

hydroxyhydroperoxides present on the filter as discussed above, then these species may be

slowly decomposing to form hydrogen peroxide, which rapidly volatilizes.

At low temperature, the volatilization and reaction rates of the peroxides are expected

to be much slower than at room temperature. Consequently, the loss of the peroxides through

these processes will be less significant, accounting for the stable yields observed. The

demonstrated stability of the peroxides at -20 °C for the entire period of observation suggests

that it is acceptable to store collected samples prior to analysis for hydrogen peroxide. Thus,

there is no recommended change to the current protocol commonly used in the community.

3.3.4 Effect of Parafilm

The effect of using parafilm to seal samples in the filter holders was investigated

using both room temperature and freezer limonene SOA samples. Following collection, the

filters were placed in the filter holders and immediately parafilmed where applicable. Filter

samples were stored for 24 hours before extraction and peroxide analysis. The results are

summarized in Table 3.2. It was found that storing filter samples with parafilm had no

significant effect on the yield of peroxides.

Table 3.2 – Peroxide yields of filters stored for 24 hours using parafilm

Yield (% mole)

With Parafilm Without Parafilm

Room Temperature 1.81 1.66

Freezer 4.13 5.31

55

3.3.5 In-Solution Stability

The stability of peroxides in aqueous solutions of α-pinene and limonene SOA had

very similar trends following extraction. The yields of peroxides in aqueous solutions of

SOA samples extracted 24 hours after collection can be found in Figure 3.6 as a function of

solution age. Within the first 48 hours after extraction, the amount of peroxide in solution

increased by approximately 50% of its original yield. After 48 hours, the peroxide slowly

decayed in the solution, reaching a yield of approximately 0% within 7 days.

A control experiment was performed in which the Teflon filter was removed from the

aqueous solution following the initial 15 minute extraction period. In this sample, the same

increase in peroxide yields was observed within the first 48 hours following extraction. This

result confirms that the measured increase in yields is due to a chemical process within the

extracted solution, as opposed to a more efficient extraction of the filter at longer extraction

times.

Figure 3.6 – Stability of peroxides in aqueous extract of 24 hr SOA sample at room

temperature. (note: no error bars shown due to single measurements performed for

experiment)

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

0 50 100 150 200

Yie

ld (

% m

ole

)

Time Since Extraction (hr)

α-pinene

limonene

56

If the peroxides in the particle begin as organic hydroperoxides, it is possible that the

increase in yields is due to a gradual decomposition reaction to form hydrogen peroxide

within the first 48 hours. Although the total peroxide content in the extract may not be

changing, a larger fraction would be present in the H2O2 form. Since the HRP-DCF assay is

sensitive to only hydrogen peroxide, this could account for the increase in yields.

At times longer than 48 hours following extraction, there is a slow decay of the

peroxides observed in the solution. It is possible that there is a continuous chemical loss of

the peroxides by other species present in the SOA extract. At short times, this loss is masked

by the conversion of organic hydroperoxides to hydrogen peroxide, resulting in an observed

increase. As the conversion reaction slows, however, the loss process dominates, resulting in

an overall decrease in peroxide yields.

In aqueous solution, the exponential loss that was observed on-filter at room

temperature is no longer significant. In water, hydrogen peroxide and organic

hydroperoxides have a very high solubility, and as a result they will not be volatile. A control

experiment was performed in which a standard of hydrogen peroxide was kept in the dark at

room temperature under the same conditions as the SOA extracts. It was found that the

concentration of the hydrogen peroxide was stable for the initial 3 days of storage, before

gradually decaying. This suggests that at least a portion of the observed loss of peroxides in

aqueous SOA samples is due to the volatilization from the solution.

3.3.6 SOA Photolysis in Solar Simulator

Aqueous extracts of α-pinene and limonene SOA were photolyzed using the solar

simulator for 3 hours. The results of these experiments are found in Figure 3.7, with the

closed points representing photolyzed samples and the open points representing the

corresponding dark control sample. No significant change in peroxide yields was observed in

α-pinene solutions that were photolyzed or kept in the dark. In solutions containing α-pinene

and iron, a decrease in peroxide levels were observed in photolyzed solutions, with a slightly

less significant decrease in the dark solution. A linear increase in peroxide yields was

observed in photolyzed limonene solutions, while the dark solution did not have a significant

57

change in peroxide yields. In solutions containing limonene SOA and iron, an increase in

peroxide yields was observed after 1 hour of photolysis, however there was no significant

change in yields during the 2 subsequent hours of photolysis. The dark limonene solution

containing iron did not have a significant change in peroxide yields.

1.50

2.00

2.50

3.00

3.50

4.00

4.50

0 50 100 150 200

Yie

ld (

% m

ole

)

Photolysis time (min)

α-pinene - light

α-pinene - dark

(a)

58

1.50

2.00

2.50

3.00

3.50

4.00

0 50 100 150 200

Yie

ld (

% m

ole

)

Photolysis time (min)

α-pinene + Fe - light

α-pinene + Fe - dark

1.50

2.00

2.50

3.00

3.50

4.00

4.50

0 50 100 150 200

Yie

ld (

% m

ole

)

Photolysis Time (min)

limonene - light

limonene - dark

(b)

(c)

59

Figure 3.7 – Peroxide yields from SOA photolysis using a xenon arc lamp in the solar

simulator.

In photolyzed solutions containing only α-pinene, no change in peroxide was

observed, suggesting that there is no photolytic process involving the peroxides. This result

is consistent with a control experiment in which an aqueous hydrogen peroxide standard was

photolyzed and no change was observed. In the presence of iron, however, the decrease in

peroxide yields suggests that there is a photolytic loss process involving iron. Again, this

result is consistent with the photolysis of a hydrogen peroxide standard containing iron. The

observed decrease of hydrogen peroxide in the presence of iron is likely due to photo-Fenton

chemistry, in which iron will consume peroxide to generate HOX radicals. The increase in

peroxide yields observed in photolyzed limonene extracts indicates that there is a photolytic

production of peroxides. In the presence of iron, there is an increase in peroxide within the

first hour of photolysis, followed by stable yields for the remaining measurements. It is

suspected that there is a competition within the solution between the photolytic production of

peroxides from the limonene SOA and the photolytic loss of peroxides from the iron. This

general topic of photochemical production of peroxide from aqueous limonene SOA

photolysis is potentially interesting, but we would require better speciated measurements of

the peroxides to make definitive claims about this mechanism.

1.50

2.00

2.50

3.00

3.50

4.00

4.50

0 50 100 150 200

Yie

ld (

% m

ole

)

Photolysis Time (min)

limonene + Fe - light

limonene + Fe - dark

(d)

60

3.3.7 SOA Photolysis under Fluorescence Lights

The results of SOA photolysis using the fluorescent lights in the laboratory are

summarized in Table 3.3. There is no significant change in the yields of peroxides between

light and dark samples from limonene SOA both on filter and in solution. This result

suggests that, in an indoor environment, there will be no photolytic production or loss of

peroxides in the SOA material.

Table 3.3 – Peroxide yields from photolysis of SOA on-filter and in-solution for 7 hours

under fluorescent lights

Light Dark

Average Yield

(% mole)

Standard

Deviation

Average Yield

(% mole)

Standard

Deviation

α-pinene

filter 2.11 0.08 2.07 0.13

solution 5.24 0.81 5.37 0.90

Limonene

filter 1.98 0.20 2.24 0.12

solution 4.47 0.34 4.47 0.53

3.4 Conclusions and Environmental Implications

The yields of peroxides in secondary organic aerosols generated through α-pinene

and limonene ozonolysis were measured using the HRP-DCF assay. The stability of the

peroxides was measured both on particle and in aqueous solution under a variety of

atmospherically and indoor related conditions. It was found that the yields of peroxides in

both α-pinene and limonene SOA were not significantly different, comprising 4% of the

SOA molecular mass. Comparing yields of SOA generated in the flow tube set up to those

collected using the environmental chamber demonstrated that the mass loading of particles

did not have an effect on the peroxide yields.

On-particle at room temperature, the peroxides were exponentially lost, with a 50%

decrease in yield within 24 hours. In the freezer, however, there was no loss in peroxides

observed within 1 week of storage. These results suggest that, at room temperature, there is a

rapid volatilization or reaction within the particle that is causing the loss of peroxides. Since

it has been shown that the peroxides will not exist as hydrogen peroxide on the particle, one

61

possibility is that they are in the form of α-hydroxyperoxides. These species are known to be

in equilibrium with hydrogen peroxide and organics with carbonyl functional groups.

In the freezer, however, the loss processes observed at room temperature become

insignificant due to the slowing of rates caused by the reduced temperatures. Consistent with

this observation, in aqueous solution, there was an observed increase in peroxide yields

within the first 48 hours following extraction at room temperature. This increase is likely due

to the chemical decomposition reaction occurring within the extracted solution, of the

conversion of organic hydroperoxides to hydrogen peroxide.

The photolysis of the extracted SOA solutions using the solar simulator produced

inconclusive results that warrant further investigation. No change in peroxide yields were

observed during the photolysis of α-pinene SOA, however a linear increase was observed in

the photolysis of limonene SOA. In the presence of iron, there was an observed decrease in

peroxide in the α-pinene extract. With iron, the yields of peroxide in limonene SOA

increased within the first hour of photolysis, stabilizing for the remaining 2 hours of

photolysis. Using the fluorescent light source to reproduce indoor photolysis, no changes in

peroxide yields were observed in-solution or on-particle.

Based on the results of this study, it is expected that the peroxides begin in the

aerosol particle as small organic hydroperoxides. With time, these peroxides are converted to

hydrogen peroxide both on-particle and in-solution. Although the peroxides only comprise a

small portion of the SOA, they have potential to influence chemistry indoors and to affect

human health. Outdoors, it is anticipated that these peroxides will be available to participate

in oxidation chemistry within the aerosol particles in the atmosphere, for example through

Fenton like chemistry or photolysis to form OH.[1] Oxidation chemistry due to the peroxides

would have a greater opportunity to proceed within aqueous particles, as the lifetime of

peroxides was demonstrated to be longer in solution. If the peroxides are lost from the

particles through volatilization, this process could be a source of oxidants to the gas phase.

Indoors, SOA can be either inhaled directly or it can deposit on indoor surfaces where

it can become part of the semi-volatile organic matter that is known to be ubiquitously

present.[27] For the airborne particles, peroxides associated with the SOA, as one component

62

of the general class of molecules referred to as Reactive Oxygen Species (ROS), have the

potential to cause oxidative stress upon inhalation. Based on our work, it is expected that the

peroxides exist as hydroperoxides in the particles – and not as hydrogen peroxide – thus

allowing them to be carried deep into the respiratory system; being so volatile and soluble, it

is unlikely that hydrogen peroxide itself can pass long distances into the respiratory

pathways. Once the particles deposit and dissolve in the lung fluid, the hydroperoxides will

likely decompose to form hydrogen peroxide, a potent ROS constituent. Given that air

exchange rates indoors are typically on the hour of orders, the loss of peroxides at room

temperature that we observed occurring on the day-to-multiday timescale is likely to be

unimportant.

A second potential impact of these peroxides is if they deposit onto surfaces indoors.

In this case, they can be potentially ingested after touching theses surfaces or, probably more

importantly, they may participate in condensed phase chemistry that could impact the

composition of the films. In particular, one question that we are now evaluating is whether

these peroxides can participate in the formation of more active oxidants, such as the

condensed phase hydroxyl radical, under indoor conditions. To be important, this chemistry

would have to occur in the films on the first day or two of SOA deposition given the

timescale we observed for on-filter decay of these species.

63

3.5 References

[1] D.W. Gunz, and M.R. Hoffmann, Atmospheric chemistry of peroxides - a review, in

Atmospheric Environment Part a-General Topics, 1990, pp. 1601-1633.

[2] S.K. Friedlander, and E.K. Yeh, The submicron atmospheric aerosol as a carrier of

reactive chemical species: Case of peroxides, in Applied Occupational and

Environmental Hygiene, 1998, pp. 416 - 420.

[3] L.A. Morio, K.A. Hooper, J. Brittingham, T.H. Li, R.E. Gordon, B.J. Turpin, and

D.L. Laskin, Tissue injury following inhalation of fine particulate matter and

hydrogen peroxide is associated with altered production of inflammatory mediators

and antioxidants by alveolar macrophages, in Toxicology and Applied

Pharmacology, 2001, pp. 188-199.

[4] M. Lippmann, and R.E. Albert, Effect of particle size on regional deposition of

inhaled aerosols in human respiratory tract, in American Industrial Hygiene

Association Journal, 1969, pp. 257-&.

[5] G. Oberdorster, Pulmonary effects of inhaled ultrafine particles, in International

Archives of Occupational and Environmental Health, 2001, pp. 1-8.

[6] F. Tao, B. Gonzalez-Flecha, and L. Kobzik, Reactive oxygen species in pulmonary

inflammation by ambient particulates, in Free Radical Biology and Medicine, 2003,

pp. 327-340.

[7] D. Moore, G.D. Robson, and A.P.J. Trinci, 21st Century Guidebook to Fungi,

Cambridge University Press, Cambridge, United Kingdom, 2011.

[8] S.A. Gurgueira, J. Lawrence, B. Coull, G.G.K. Murthy, and B. Gonzalez-Flecha,

Rapid increases in the steady-state concentration of reactive oxygen species in the

lungs and heart after particulate air pollution inhalation, in Environmental Health

Perspectives, 2002, pp. 749-755.

[9] L. Risom, P. Moller, and S. Loft, Oxidative stress-induced DNA damage by

particulate air pollution, in Mutation Research-Fundamental and Molecular

Mechanisms of Mutagenesis, 2005, pp. 119-137.

[10] C.J. Weschler, and H.C. Shields, Indoor ozone/terpene reactions as a source of

indoor particles, in Atmospheric Environment, 1999, pp. 2301-2312.

[11] J.Y. Zhang, K.E.H. Hartz, S.N. Pandis, and N.M. Donahue, Secondary organic

aerosol formation from limonene ozonolysis: Homogeneous and heterogeneous

influences as a function of NOx, in Journal of Physical Chemistry A, 2006, pp.

11053-11063.

[12] X. Chen, and P.K. Hopke, A chamber study of secondary organic aerosol formation

by limonene ozonolysis, in Indoor Air, 2010, pp. 320-328.

[13] T.H. Li, B.J. Turpin, H.C. Shields, and C.J. Weschler, Indoor hydrogen peroxide

derived from ozone/d-limonene reactions, in Environmental Science & Technology,

2002, pp. 3295-3302.

[14] J. Toftum, S. Feund, T. Salthammer, and C.J. Weschler, Secondary organic aerosols

from ozone-initiated reactions with emissions from wood-based materials and a

"green" paint, in Atmospheric Environment, 2008, pp. 7632-7640.

[15] X. Chen, and P.K. Hopke, Secondary organic aerosol from alpha-pinene ozonolysis

in dynamic chamber system, in Indoor Air, 2009, pp. 335-345.

64

[16] B.C. Singer, B.K. Coleman, H. Destaillats, A.T. Hodgson, M.M. Lunden, C.J.

Weschler, and W.W. Nazaroff, Indoor secondary pollutants from cleaning product

and air freshener use in the presence of ozone, in Atmospheric Environment, 2006,

pp. 6696-6710.

[17] Z.H. Fan, P. Lioy, C. Weschler, N. Fiedler, H. Kipen, and J.F. Zhang, Ozone-initiated

reactions with mixtures of volatile organic compounds under simulated indoor

conditions, in Environmental Science & Technology, 2003, pp. 1811-1821.

[18] C.J. Weschler, Ozone in indoor environments: Concentration and chemistry, in

Indoor Air-International Journal of Indoor Air Quality and Climate, 2000, pp. 269-

288.

[19] M.L. Walser, Y. Desyaterik, J. Laskin, A. Laskin, and S.A. Nizkorodov, High-

resolution mass spectrometric analysis of secondary organic aerosol produced by

ozonation of limonene, in Physical Chemistry Chemical Physics, 2008, pp. 1009-

1022.

[20] K.S. Docherty, W. Wu, Y.B. Lim, and P.J. Ziemann, Contributions of organic

peroxides to secondary aerosol formed from reactions of monoterpenes with O-3, in

Environmental Science & Technology, 2005, pp. 4049-4059.

[21] Y. Wang, H. Kim, and S.E. Paulson, Hydrogen peroxide generation from alpha- and

beta-pinene and toluene secondary organic aerosols, in Atmospheric Environment,

2011, pp. 3149-3156.

[22] A.P. Bateman, S.A. Nizkorodov, J. Laskin, and A. Laskin, Photolytic processing of

secondary organic aerosols dissolved in cloud droplets, in Physical Chemistry

Chemical Physics, 2011, pp. 12199-12212.

[23] P. Mertes, L. Pfaffenberger, J. Dommen, M. Kalberer, and U. Baltensperger,

Development of a sensitive long path absorption photometer to quantify peroxides in

aerosol particles (Peroxide-LOPAP), in Atmospheric Measurement Techniques,

2012, pp. 2339-2348.

[24] X. Chen, and P.K. Hopke, A chamber study of secondary organic aerosol formation

by linalool ozonolysis, in Atmospheric Environment, 2009, pp. 3935-3940.

[25] A. Mutzel, M. Rodigast, Y. Iinuma, O. Boege, and H. Herrmann, An improved

method for the quantification of SOA bound peroxides, in Atmospheric Environment,

2013, pp. 365-369.

[26] R. Zhao, A.K.Y. Lee, R. Soong, A.J. Simpson, and J.P.D. Abbatt, Formation of

aqueous-phase alpha-hydroxyhydroperoxides (alpha-HHP): potential atmospheric

impacts, in Atmospheric Chemistry and Physics, 2013, pp. 5857-5872.

[27] Q.T. Liu, R. Chen, B.E. McCarry, M.L. Diamond, and B. Bahavar, Characterization

of polar organic compounds in the organic film on indoor and outdoor glass

windows, in Environmental Science & Technology, 2003, pp. 2340-2349.

65

Chapter 4

Implications and Future Work

4.1 Environmental Implications

In the first project, the potential coupling of metal-catalyzed redox cycles was

investigated by evaluating peroxide production in systems in the presence and absence of

transition metals. With a coupling of the redox cycles, the production of peroxide was

expected to decrease, demonstrating an irreversible loss of HOX. Under the experimental

conditions used, the results were not able to demonstrate this coupling, as iron’s affinity to

form complexes dominated over its participation in the coupling reaction. To better evaluate

the significance of this coupling chemistry in the atmosphere, a more atmospherically

relevant sample should be used.

The results of this project demonstrated that a substantial amount of peroxide could

be produced from a system containing one simple organic with iron. This suggests that there

will be the potential to have a larger extent of iron complexation within atmospheric

aerosols, which are known to have a much more diverse mixture of many concentrated

organic species. The photolysis of these iron complexes could be a significant source of

peroxides within the particle phase, however this will be heavily dependent on the organics

that are present, as not all complexes have been shown to generate peroxide. This could

increase the amount of oxidants present within the particle, ultimately increasing its

oxidative capacity.

In the second project, the production of peroxides within secondary organic aerosols

generated from α-pinene and limonene precursors was studied. The stability of the peroxides

was investigated under a range of indoor and atmospherically relevant conditions. It was

found that the yields of peroxides in SOA generated using the different precursors were not

significantly different. While the yields of hydrogen peroxide increased in an aqueous

solution within the initial days following formation (likely due to the conversion of organic

hydroperoxides into hydrogen peroxide), there was a rapid decay of peroxides from the dry

particles.

66

In the atmosphere, the peroxides that are associated with the aerosol would be

available to participate in oxidation chemistry within the particle phase. In an aqueous

aerosol, the peroxides are expected to remain in the aqueous phase for several days due to

their demonstrated longer lifetimes. On a dry particle however, the peroxides are expected to

be rapidly lost. The volatilization of the peroxides from the particle could be a source of HOX

to the gas phase.

Indoors, the peroxides in the SOA particles can both influence oxidation chemistry

and have health implications. Upon inhalation, the peroxides associated with the particles are

expected to be carried deep within the lungs. The transport of these oxidants will cause an

imbalance between oxidants and antioxidants within the cells, causing oxidative stress. The

particles also have the potential to be deposited onto organic films commonly found indoors.

Here, the peroxides could promote oxidation within the films, changing their composition.

4.2 Future Work

In order to deepen the understanding of the results of the first project, it is

recommended that the iron complexes are investigated. To evaluate the significance of the

iron hydroxyl complex, it will be valuable to have a quantitative measurement of the OH

radicals. This could be accomplished by performing a benzoic acid OH trap experiment as

described previously, and quantifying the samples using a more sensitive mass spectrometer

than previously used. Using this concentration of OH, the amount of peroxide that would be

generated by the oxidation of acetaldehyde could be used to determine the contribution of

this reaction to the overall increase in peroxide that was observed in the acetaldehyde and

iron system. To investigate the iron organic complex, an analytical technique such as ion

chromatography with an electrical conductivity detector could be used to confirm the

production of oxalic acid through the oxidation of acetaldehyde in the system. By performing

a quantitative analysis, it will be possible to determine the contribution of the organic

complex to the increase in peroxide from the acetaldehyde and iron system.

As a consequence of measuring only hydrogen peroxide, much of the chemistry that

occurred in the transition metal samples is not well understood. In order to have a better

67

understanding of the processes in the sample, other species in the samples such as the

transition metal ions, OH and HO2 radicals could be measured using techniques such as ion

chromatography and electron spin resonance. This would confirm the concentrations and

significance of the pathways predicted using the models, and help to explain how the redox

cycles may be influenced by the metals.

The experiments conducted in the first project could be performed using an

alternative source of aqueous HOX radicals. Propanaldehyde is an example of a simple

organic that is expected to generate HO2 radicals upon photolysis, while not likely to form

oxalic acid upon oxidation. If the same trends in peroxide production are confirmed, this

would strengthen the prediction that the coupling chemistry is not likely to be significant in

an ambient particle. In order to better evaluate the efficacy of the coupling chemistry,

however, the inclusion of other species commonly found in aerosol particles and cloud

droplets at relevant concentrations is important. By increasing the diversity of the species in

the samples, it will be possible to better predict the interferences some species may have on

the coupling reaction.

A more thorough speciation of the peroxides measured in the second project would

improve the understanding of the species that are generated through SOA formation. This

could be done by using the iodide technique to measure the total peroxide concentration

(including hydrogen peroxide, organic hydroperoxides, and organic peroxides). Measuring

the initial yields and changes of the individual peroxides under the various conditions would

deepen the understanding of the conversion and loss processes affecting the observed yields

in this study.

To evaluate the broader impact of the second project, it would be interesting to see if

the peroxides associated with the SOA particles will promote oxidation chemistry in the

particle phase or indoor films. The peroxides could generate HOX radicals through reaction

or photolysis, which could then oxidize other material within the particle or film. Having a

more thorough understanding of the oxidation that may occur following particle formation

could improve the knowledge of the aging processes that will change the composition of

aerosols and indoor films.